Resonance and Formal Charge

Learning Objective 2.6.A Represent a molecule with a Lewis diagram that accounts for resonance between equivalent structures or that uses formal charge to select between nonequivalent structures.

Quick Notes

Resonance occurs when more than one valid Lewis structure exists for a molecule.
Resonance structures differ only in the position of electrons, not the position of atoms.
The actual structure is a resonance hybrid — an average of all valid structures.
Formal charge helps choose the best Lewis structure when multiple possibilities exist.

Formal charge = (valence electrons) – (nonbonding electrons) – (½ × bonding electrons)

Best structures:
- Minimize formal charges
- Place negative charges on more electronegative atoms
- Obey the octet rule, unless expanded octets apply
Lewis structures are models — they don’t explain everything, especially with odd-electron species or delocalized bonding.

Full Notes

In some molecules, a single Lewis structure does not accurately represent the distribution of electrons. Instead, the true structure is best represented by multiple contributing resonance forms. Additionally, when comparing possible structures, we use formal charge to decide which is most appropriate.

What Is Resonance?

Resonance occurs when two or more valid Lewis structures can be drawn for a molecule by moving electrons only (not atoms). These are called resonance structures.

The actual molecule is a resonance hybrid, which averages the bond characteristics from all contributing structures. Resonance is a way of representing delocalized electrons (e.g. in a carboxylate group or benzene ring).

Rules for Resonance:

Only electrons (usually lone pairs or π bonds) can move
All structures must be valid Lewis structures with correct total valence electrons
Atom positions must remain the same

Examples of Resonance Structures

a. Ozone (O₃)

Two valid structures:

AP Chemistry ozone resonance: two valid Lewis structures and resonance hybrid depiction with delocalized bonding

One double bond between the left O and middle O
One double bond between the middle O and right O
Actual structure: a resonance hybrid where both O–O bonds are intermediate in length and strength

b. Carbonate ion (CO₃²⁻)

AP Chemistry carbonate ion resonance: three equivalent Lewis structures and the resonance hybrid

Three resonance structures, each with one C=O and two C–O⁻ bonds
Actual structure: resonance hybrid where all C–O bonds are intermediate in length and strength and are equivalent

c. Benzene (C₆H₆)

AP Chemistry benzene resonance: two Kekulé structures and resonance hybrid ring with delocalized electrons

Six carbon atoms in a ring, with alternating single and double bonds
Resonance hybrid with delocalized electrons forming a stable ring

What Is Formal Charge?

Formal charge is a tool for evaluating how well electrons are distributed in a Lewis structure.

Formula:

AP Chemistry formal charge equation: valence electrons minus nonbonding electrons minus one-half bonding electrons

formal charge = (valence electrons) – (nonbonding electrons) – (½ × bonding electrons)

Where:

Valence electrons: from group number in periodic table (e.g. 5 for N, 6 for O)
Non-bonding electrons: lone pair electrons on the atom
Bonding electrons: total electrons shared in bonds around the atom

Guidelines for Using Formal Charge

A preferred Lewis structure will:

Have the smallest possible formal charges
Avoid like charges on adjacent atoms
Place negative formal charges on more electronegative atoms
Have a total charge equal to the overall charge on the molecule or ion

Example:The Nitrate Ion (NO₃⁻)

AP Chemistry nitrate ion Lewis structure with formal charge assignments on nitrogen and oxygen atoms

In the nitrate ion:

N has a formal charge of +1
The double bonded O has a formal charge of 0
Each single bonded O has a formal charge of -1
This means overall the charge of the nitrate ion is -1 as (+1) + (-1) + (-1) = -1

Worked Example

Determine which is the better Lewis structure for NO₂⁻ (nitrite ion):

Structure A:

Both oxygen atoms single-bonded to nitrogen
Formal charges: N = +1, O = −1 (on both)

Structure B (resonance form):

AP Chemistry nitrite ion Structure B with one N=O double bond and one single bond; minimized formal charges

One oxygen double-bonded, the other single-bonded
Formal charges: N = 0, one O = 0, one O = −1

Answer: Structure B is better because it minimizes formal charges and spreads the negative charge over the more electronegative oxygen atoms. However, remember the true, actual structure is a resonance hybrid of two such forms.

Limitations of Lewis Structures

Lewis structures are useful but have limitations:

Cannot represent delocalized electrons well (e.g. in aromatic compounds)
Struggle with odd-electron species (radicals)
Do not account for molecular shape — this is handled by VSEPR
Do not directly show bond strength or length when resonance is involved

Summary

Resonance structures are used when a single Lewis diagram cannot represent the full distribution of electrons in a molecule. The actual structure is a resonance hybrid, which averages the possibilities. Formal charge helps choose between competing Lewis structures by identifying the most stable and realistic arrangement of electrons. While useful, Lewis structures are simplified models and do not explain every property of molecules.

Key points to remember:
- Resonance = multiple valid electron arrangements
- The true structure is an average of all resonance forms
- Formal charge = valence – nonbonding – ½ bonding
- Best structures minimize formal charges and follow the octet rule
- Lewis structures have limitations but are a valuable model