Chemistry 2e · Chemical Bonding and Molecular Geometry
Formal Charges and Resonance
On this page 9 sections
In 30 seconds
Sometimes more than one Lewis structure can be drawn for the same molecule, and chemists need a way to choose among them. formal charge Bookkeeping charge on an atom, shared electrons split evenly Full entry → is the bookkeeping tool: it assigns each atom a charge as if every bond were split evenly. When several valid structures differ only in where the electrons sit — not where the atoms are — the molecule shows resonance Multiple valid Lewis structures for one atomic skeleton Full entry →: the real molecule is a blend, or hybrid, of all contributors, with electrons delocalized over several atoms. This topic explains how to calculate formal charges, how to use them to pick the most plausible structure, and why resonance explains the equal bond lengths and stability of carbonate, nitrate, and ozone.
Why this matters
Formal charge and resonance turn "a valid Lewis structure" into "the best description of a real molecule." Carbonate, CO₃²⁻, is central to blood pH buffering; nitrate, NO₃⁻, matters in agriculture and water quality; ozone, O₃, shields the atmosphere from UV radiation. All three are resonance-stabilized — which is why they exist as stable species. On exams, formal-charge rules quickly eliminate poor candidates, and resonance explains facts (equal bond lengths, fractional bond orders) that no single Lewis structure can.
The college version
Core Concepts
What formal charge is — and what it is not
Formal charge (FC) is an electron-bookkeeping device: count an atom's valence electrons V, subtract its lone-pair electrons L and half its shared (bonding) electrons B/2:
FC = V - L - B2
Two checks make it trustworthy: the formal charges of all atoms must sum to the overall charge, and FC is bookkeeping, not a prediction of where real charge density sits. In CO, the formal charge lands on carbon even though oxygen is more electronegative.
Using formal charge to choose the best structure
When two or more Lewis structures are plausible, prefer the one that:
- Has formal charges closest to zero on every atom,
- Places negative formal charges on the most electronegative atoms,
- Avoids large like-signed formal charges on adjacent atoms.
These are guidelines — the lowest-formal-charge structure usually matches reality, but electronegativity and octet satisfaction still matter.
Resonance: several structures, one real molecule
Resonance occurs when a molecule or ion can be drawn with the same atomic skeleton in more than one valid Lewis structure. Ozone, O₃, has two equivalent structures — double bond left or right — that differ only in electron placement; atoms never move. The structures are connected by a double-headed arrow, ↔ , and no single drawing is "the real molecule." The real molecule is a resonance hybrid The real molecule, an average of all contributing structures Full entry → — an average of the contributors, with electrons delocalized across all three atoms. Resonance lowers the molecule's energy, which is why hybrids are more stable than any single contributor predicts.
Consequences: equal bonds and fractional bond order
A resonance hybrid averages the bond orders of its contributors. Ozone's two structures each show one single and one double O–O bond, so the hybrid has bond order (1 + 2)/2 = 1.5 for both. Experiment agrees: ozone's O–O bonds are identical in length, between a single bond (about 148 pm) and a double bond (about 121 pm). The same reasoning gives carbonate and nitrate bond orders of (1 + 2 + 2)/3 = 4/3 per bond. Identical bond lengths within an ion are a fingerprint of resonance.
How It Works / Step-by-Step Process
- Draw a valid Lewis structure; identify each atom's V, L, and B.
- Apply FC = V - L - B/2 to every atom.
- Sum the formal charges; the total must equal the molecule's or ion's charge.
- If other structures are possible, repeat and apply the preference rules.
- If structures differ only in electron placement, label them resonance contributors and describe the hybrid.
Common Confusions
| Do not confuse | With | Difference |
|---|---|---|
| Formal charge | Actual (real) charge | FC assumes equal sharing; real charge depends on electronegativity (see CO) |
| Resonance structures being real molecules that interconvert | A single hybrid molecule | The hybrid is real; the drawings are limiting descriptions, with no actual flipping |
| The resonance arrow ↔ | The equilibrium arrow ⇌ | ↔ = electron delocalization within one molecule; ⇌ = reaction at equilibrium |
| "The most stable resonance structure is the real molecule" | "The hybrid is the real molecule" | The best drawing is just a drawing; the molecule averages all contributors |
| A bond shown as double in one structure being truly double | A fractional bond order | In O₃ every O–O bond is 1.5 — between single and double |
| Formal charge predicting reactivity directly | A structure-selection tool | FC guides which structure to draw; reactivity depends on real electron density |
Eli explains
The same idea, in plain words
Explain it like I’m 10
Sometimes the same molecule can be drawn two or three correct ways, like one dance pose photographed from different angles. The real molecule is not flipping between pictures — it is a blend of all of them, with electrons spread over the whole molecule. The blend is more stable than any single picture.
Worked example
Example 1: Formal charges in carbon monoxide, CO
Carbon monoxide has 10 valence electrons (4 + 6 = 10); its best Lewis structure is a triple bond with a lone pair on each atom: :C ≡ O:.
For carbon, V = 4, L = 2, B = 6:
FCC = 4 - 2 - 62 = -1
For oxygen, V = 6, L = 2, B = 6:
FCO = 6 - 2 - 62 = +1
The charges sum to (-1) + (+1) = 0, matching the neutral molecule. Note the surprising result: the formal charge is negative on carbon, the less electronegative atom. Formal charge assumes equal sharing, so it is bookkeeping — the real molecule has only a small dipole. This is the classic warning against mistaking formal charge for real charge.
Example 2: Ozone, O₃, and its two resonance structures
Ozone has 18 valence electrons (3 × 6 = 18). Two equivalent structures are possible: A, O=O-O, and B, O-O=O.
For the central oxygen, V = 6, L = 2, B = 6:
FCcentral = 6 - 2 - 62 = +1
The double-bonded terminal oxygen has V = 6, L = 4, B = 4:
FC = 6 - 4 - 42 = 0
The single-bonded terminal oxygen has V = 6, L = 6, B = 2:
FC = 6 - 6 - 22 = -1
The charges +1, 0, -1 sum to zero; Structure B mirrors A. The real molecule is the hybrid: both O–O bonds have order 1.5 and equal length (about 128 pm). The +1 always sits on the central oxygen; the -1 rides whichever terminal oxygen carries the single bond.
Example 3: Choosing between structures with formal charge
Sulfur dioxide, SO₂, has 18 valence electrons (6 + 2 × 6 = 18); its two resonance structures place the double bond on either oxygen. In either structure, sulfur has V = 6, L = 2, B = 6:
FCS = 6 - 2 - 62 = +1
The double-bonded oxygen has FC = 0, the single-bonded oxygen FC = -1. The +1 on sulfur and -1 on oxygen are acceptable because oxygen, being more electronegative, carries the negative charge better. This matches the observed equal S–O bond lengths (bond order 1.5).
Key takeaways
- Formal charge: FC = V - L - B/2, with B = shared (bonding) electrons.
- The sum of all formal charges must equal the overall charge — use it as a check.
- Prefer structures with formal charges closest to zero; put negative formal charges on the most electronegative atoms.
- Resonance structures differ only in electron placement; atoms never move; connect with ↔ , not ⇌ .
- The real molecule is a resonance hybrid — a weighted average of contributors — and is more stable than any single structure.
- Identical bond lengths and fractional bond orders (O₃: 1.5; CO₃²⁻: 4/3) are experimental evidence of resonance.
Check yourself
5 review questions from the chapter. Try each one, then open the answer.
Calculate the formal charge on nitrogen in NH₄⁺.
Show answer
Nitrogen has V = 5, L = 0, B = 8 (four N–H bonds): FC = 5 - 0 - 8/2 = +1. Hydrogens are 0 each, so the total +1 matches the ion's charge.
Why are all three N–O bonds in NO₃⁻ the same length?
Show answer
Resonance delocalizes electrons over all three N–O bonds, averaging the bond order to 4/3 for each — so none is shorter or longer.
What is the difference between ↔ and ⇌ ?
Show answer
↔ links resonance structures of the same molecule; ⇌ separates reactants and products of a reversible reaction.
In cyanate, NCO⁻, which atom should carry a negative formal charge — and why?
Show answer
Oxygen — the preference rules place negative formal charge on the most electronegative atom.
What must the formal charges of every atom in an ion sum to?
Show answer
The overall charge of the molecule or ion — zero for a neutral molecule, -2 for CO₃²⁻. Use this as a built-in check.
Study tools & related lessonsKey vocabulary · Related
Key vocabulary
- formal charge
- Bookkeeping charge on an atom, shared electrons split evenly
- valence electron
- Electron in the outermost shell, counted as V
- lone pair
- Pair of electrons owned entirely by one atom, counted as L
- bonding electrons
- Shared electrons, counted as B, split B/2 per atom
- resonance
- Multiple valid Lewis structures for one atomic skeleton
- resonance hybrid
- The real molecule, an average of all contributing structures
- delocalization
- Spreading electrons over several atoms rather than fixing them in one bond
- fractional bond order
- Averaged bond order from resonance, e.g., 1.5 in O₃
Sources & references
This lesson was adapted from the open educational references above; their licenses and attributions are preserved. See Copyright & Licensing.
Educational content only. It is not medical, legal or professional advice. Found an error? Tell us.