Chemistry: Atoms First 2e · Chemical Bonding and Molecular Geometry

Formal Charges and Resonance

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In 30 seconds
Why this matters
The college version
Eli explains
Worked example
Key takeaway
Check yourself
Study tools
Sources & references

In 30 seconds

When several Lewis structures can be drawn for a molecule, chemists need a way to choose among them — and to describe molecules no single structure captures. is a bookkeeping tool that assigns each atom a charge by comparing its valence electrons with those it "owns" in the drawing. describes molecules using two or more Lewis structures differing only in electron placement; the real molecule is a blend, or , of those forms. Together they pick the most plausible structure and explain why ozone, nitrate, and benzene behave as they do.

Why this matters

Formal charge is the tie-breaker for choosing the best Lewis structure: among valid drawings, the one with formal charges closest to zero — negative charges on the most electronegative atoms — is usually best. Resonance explains facts a single structure cannot: the carbonate ion's three equivalent C–O bonds, ozone's identical O–O bonds, and benzene's unusual stability. In biology, resonance describes proteins' planar amide bonds, delocalized charge in phosphates, and the stability of many drugs.

The college version

Core Concepts

Formal charge: the bookkeeping equation

Formal charge compares an atom's valence electrons with those assigned to it in a structure:

$FC = V - L - 12 B$

where FC is formal charge, V is the free atom's valence electrons, L is lone-pair electrons, and B is bonding electrons (each bond contributes 2; shared pairs count fully for both atoms). The sum of all formal charges always equals the overall charge — a built-in check.

Choosing the best Lewis structure

Among structures that satisfy the octet rule, prefer the one that:

Has formal charges closest to zero (fewest nonzero charges).
Puts negative formal charge on the most electronegative atom.
Avoids like charges on adjacent atoms.

Formal charge is bookkeeping, not real charge; an atom at +1 is not necessarily a cation. Real charge distribution also depends on electronegativity.

Resonance: multiple structures, one molecule

Resonance applies when two or more valid Lewis structures of the same molecule differ only in electron placement while atoms stay fixed. These resonance forms are connected by a double-headed arrow. The real molecule is a resonance hybrid — not a flipping mixture, but a single structure averaging all contributors. Averaging gives intermediate bond lengths and bond orders and lowers energy: resonance stabilizes molecules.

Recognizing valid resonance forms

A valid must: keep all atoms in place; preserve the total electron count; never break a sigma bond (only pi electrons and lone pairs move); and keep the total charge. Moving atoms creates an isomer or a different molecule. Practice with ozone, carbonate, and nitrate until curved-arrow notation feels automatic.

How It Works / Step-by-Step Process

Calculating formal charges

Draw a complete Lewis structure with all bonds and lone pairs.
For each atom, record V (free-atom valence electrons from the group number).
Count L: lone-pair electrons on that atom.
Count B: bonding electrons involving that atom (each bond contributes 2).
Apply FC = V - L - 12B for every atom.
Check: sum of formal charges = overall charge.

Drawing resonance forms

Identify pi bonds and lone pairs that can move (delocalize).
Move electrons with curved arrows: a lone pair becomes a pi bond, or a pi bond becomes a lone pair.
Keep every atom fixed; only electrons move.
Verify each form keeps the same atoms, electron count, and total charge.
Connect all forms with double-headed arrows; the hybrid averages them.

Common Confusions

Do Not Confuse	With	The Difference
Formal charge	Real (partial) charge	Formal charge is bookkeeping based on electron counting; real charge depends on electronegativity and is usually fractional
Resonance forms	Different molecules or isomers	Resonance forms have identical atom positions and differ only in electron placement
Resonance hybrid	A rapidly interconverting mixture	The hybrid is one molecule whose electron density averages the forms; there is no flipping
Double-headed arrow (resonance)	Equilibrium arrows (reaction)	Resonance arrows connect structures of the same molecule; equilibrium arrows connect different species
"The structure is X"	"The structure is an average of X and Y"	For resonance molecules, the single-structure description is incomplete; the hybrid is the reality
Bond order 1.5	Half a bond	One shared pair spread over two equivalent positions — a full but slightly weakened bond in each

Eli explains

The same idea, in plain words

Explain it like I’m 10

Formal charge is like keeping score in a card game: count the electrons an atom brought, count what it keeps, and the difference is its score. Resonance is like describing a horse that is neither all black nor all white — you draw two pictures, one mostly black and one mostly white, and the real animal is the blend. That is why resonance molecules have bonds that look halfway between single and double.

Worked example

Example 1: Formal charges in the nitrate ion (NO₃⁻)

Draw the common form with one N=O double bond and two N–O single bonds. For nitrogen, V = 5, L = 0, B = 8 (four bonds total, counting the double bond as two):

$FC N = 5 - 0 - 12 (8) = 5 - 4 = +1$

For the doubly bonded oxygen, V = 6, L = 4, B = 4:

$FC O = 6 - 4 - 12 (4) = 6 - 4 - 2 = 0$

For each singly bonded oxygen (carrying the negative charge), V = 6, L = 6, B = 2:

$FC O = 6 - 6 - 12 (2) = 6 - 6 - 1 = -1$

Sum: +1 + 0 + (-1) + (-1) = -1, matching the ion's charge. The three resonance forms are equivalent, so the real ion has three identical N–O bonds, bond order 43 (four shared pairs over three positions).

Example 2: Resonance and bond order in ozone (O₃)

Ozone has 18 valence electrons total. The two resonance forms are O=O–O and O–O=O; the central oxygen is +1 and the singly bonded terminal oxygen is −1 in each form. Because both forms are equivalent, the two O–O bonds are identical in the real molecule. The bond order is:

$BO = number of bonds in all forms number of forms = 1 + 2 2 = 1.5$

That is why both O–O bonds share one length, intermediate between an O–O single bond (about 148 pm in hydrogen peroxide) and an O=O double bond (about 121 pm in O₂). A single-structure drawing would wrongly predict two different bond lengths.

Example 3: Choosing the best structure for the cyanate ion (OCN⁻)

Count electrons: 6 (O) + 4 (C) + 5 (N) + 1 (charge) = 16. Two candidates satisfy the octet rule: O–C≡N and O=C=N. In the first, oxygen: 6 - 4 - 2 = 0, carbon: 4 - 0 - 4 = 0, nitrogen: 5 - 2 - 3 = 0. In the second, oxygen: 6 - 2 - 3 = +1, carbon: 0, nitrogen: 5 - 4 - 2 = -1. Structure 1 has all formal charges zero, so it is the better description — no formal charges while the ion as a whole still carries −1.

Key takeaways

Formal charge equation: FC = V - L - 12B; sum of FC values = overall charge.
Best structure: fewest nonzero formal charges; negative FC on the most electronegative atom.
Formal charge is bookkeeping, not real charge — electronegativity controls the real distribution.
Resonance forms differ only in electron placement (pi bonds/lone pairs); atoms never move.
The resonance hybrid averages all forms: bond orders and lengths fall between the extremes.
Resonance stabilizes molecules; more resonance forms generally mean greater stability.
Use a double-headed arrow between resonance forms — never an equilibrium arrow.

Check yourself

5 review questions from the chapter. Try each one, then open the answer.

Calculate the formal charge on the central nitrogen in ammonium, NH₄⁺.
Show answer
FC = 5 - 0 - 12(8) = +1. Sum of the four H (0 each) and N (+1) equals the +1 charge.
Why must the sum of formal charges equal the overall charge of an ion?
Show answer
Because formal charge is a redistribution of the atom's own valence electrons; the total number of electrons (and hence total charge) cannot change.
What two features must all resonance forms of a molecule share?
Show answer
Identical atom positions and the same total number of electrons (and same total charge); only pi electrons and lone pairs move.
Ozone's O–O is 1.5. What bond length would you predict compared with an O–O single bond?
Show answer
Shorter than a single bond and longer than a double bond — intermediate, as observed (~128 pm in ozone).
When choosing between two valid Lewis structures, what rule decides the winner?
Show answer
Prefer the structure with formal charges closest to zero, with any negative charge on the most electronegative atom.

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Study tools & related lessonsKey vocabulary · Related

Key vocabulary

formal charge: Bookkeeping charge on an atom: valence electrons minus electrons assigned in the structure
resonance: Describing a molecule with two or more structures differing only in electron placement
resonance hybrid: The real molecule, an average of all resonance forms
resonance form: One individual Lewis structure contributing to the hybrid
delocalization: Electrons spread over several atoms rather than one bond
bond order: Number of bonds between two atoms, averaged over resonance forms

Sources & references

openstax.org — Chemistry Atoms First 2e

This lesson was adapted from the open educational references above; their licenses and attributions are preserved. See Copyright & Licensing.

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