Resonance and formal charge are essential tools in understanding how electrons are distributed in molecules, helping predict structure stability and molecular behavior.
Resonance
What Is Resonance?
In molecular chemistry, resonance describes a situation where a single Lewis structure cannot accurately depict the actual electron arrangement in a molecule or ion. Instead, the true structure is represented by a resonance hybrid, which is the weighted average of all the valid resonance structures.
Resonance structures differ only in the placement of electrons, not in the positions of atoms.
The actual molecule does not oscillate between these structures. Instead, it exists as a stable intermediate form—a blend or hybrid of all the structures.
Electrons involved in resonance are referred to as delocalized electrons, meaning they are not associated with a single atom or bond but spread across multiple atoms.
Resonance allows for a better explanation of molecules with equal bond lengths despite appearing to have single and double bonds in different resonance forms.
Key Characteristics of Resonance Structures
The skeleton of atoms (connectivity and placement) remains unchanged.
Only the positions of electrons (lone pairs and multiple bonds) vary.
Resonance involves π electrons or nonbonding electrons, typically in conjugated systems where double or lone-pair electrons are adjacent to a single bond.
The actual structure is more stable than any single resonance form, a phenomenon known as resonance stabilization.
How to Recognize Resonance
A molecule may have resonance if:
It has multiple atoms capable of forming π bonds.
It includes lone pairs adjacent to π bonds or empty p orbitals.
When drawing the lewis structure of polyatomic ions, make sure you include the brackets and charge.
There are multiple atoms with differing electronegativity that could share electrons.
Common examples include:
Ozone (O3)
Nitrate (NO3−)
Nitrite (NO2−)
Sulfate (SO4²−)
Carbonate (CO3²−)
Benzene (C6H6)
Drawing Resonance Structures
Example: Nitrate Ion (NO3−)
Calculate total valence electrons:
Nitrogen (N) contributes 5 electrons.
Oxygen (O) contributes 6 electrons per atom, and there are 3 O atoms: 6 × 3 = 18.
The ion has a −1 charge, so add 1 electron.
Total valence electrons = 5 + 18 + 1 = 24 electrons.
Choose the central atom:
Nitrogen is the central atom (only one atom of N, and it’s less electronegative than oxygen).
Form bonds:
Connect each oxygen to the nitrogen with a single bond.
Complete the octets:
Add lone pairs to each oxygen to complete their octets.
Count the electrons used. If there are too many, form double bonds by converting a lone pair into a bonding pair.
Add a double bond to one oxygen to reduce the total number of electrons to 24.
Recognize resonance:
The double bond could be placed with any of the three oxygen atoms.
Thus, three equivalent resonance structures exist, each with a double bond to a different oxygen.
These structures are connected with double-headed arrows (↔).
Bond Order in Resonance
Bond order reflects the average number of chemical bonds between a pair of atoms.
For nitrate:
4 total N–O bonds (3 single and 1 double over 3 atoms).
Bond order = total bonds / positions = 4 / 3 = 1.33.
Each N–O bond in nitrate is stronger than a single bond but weaker than a double bond, with equal length and energy in reality.
Resonance Does Not Imply Flipping
A common misconception is that the molecule rapidly shifts between forms. Instead, all resonance forms coexist simultaneously as contributors to the real structure.
The resonance hybrid is the most accurate representation, often drawn as a single structure with dashed lines representing partial bonds.
AP Exam Requirement
On the AP Chemistry or Biology exam:
You may be asked to draw all resonance structures.
Use brackets and include the overall charge.
Connect the structures using double-headed arrows (↔).
Do not use equilibrium arrows (⇌) which imply a chemical reaction or shift, not resonance.
Formal Charge
What Is Formal Charge?
Formal charge is a way to estimate the distribution of electric charge within a molecule. It assumes equal sharing of bonding electrons, even between atoms of different electronegativities.
It helps determine:
The most likely placement of electrons.
The most stable resonance form.
Whether a Lewis structure is valid.
A structure where atoms have formal charges closest to zero is generally preferred.
Formal Charge Formula
Use the following formula:
Formal Charge = (Valence electrons) − (Nonbonding electrons) − 1/2(Bonding electrons)
Or, for simplicity:
Formal Charge = Valence electrons − Dots − Dashes
Where:
Dots = lone pair electrons
Dashes = bonds (each dash = 1 bond)
Example: SCN⁻ (Thiocyanate Ion)
Let’s examine sulfur:
Sulfur has 6 valence electrons.
In one Lewis structure, sulfur has:
2 lone electrons (dots)
3 bonding electrons (dashes to carbon)
Formal charge = 6 − 2 − 3 = +1
Repeat this for each atom to calculate their respective formal charges.
When to Use Formal Charge
When multiple Lewis structures are possible.
When assessing resonance forms.
When predicting stability and reactivity of a molecule.
When breaking the octet rule, especially for atoms beyond period 2.
Formal charge helps rank resonance structures in terms of their contribution to the overall hybrid.
Example: Phosphate Ion (PO4³⁻)
Total valence electrons:
Phosphorus = 5
Oxygen = 6 × 4 = 24
Charge = −3 → add 3 electrons
Total = 5 + 24 + 3 = 32 electrons
Initial structure:
P in center with single bonds to 4 O atoms.
Complete octets = use 32 electrons.
Formal charge calculation:
P = 5 − 0 − 4 = +1
Each O = 6 − 6 − 1 = −1
Total charge = +1 + (4 × −1) = −3, matches the ion charge
Improving the structure:
Add a double bond between P and one O to reduce P's formal charge.
New charges:
P = 5 − 0 − 5 = 0
3 O = 6 − 6 − 1 = −1
1 O = 6 − 4 − 2 = 0
Now, the more electronegative atoms carry the negative charges, and the central atom is neutral—this structure is more stable.
Guiding Principles for Formal Charge
The sum of formal charges must equal the overall charge of the molecule or ion.
Minimize non-zero formal charges.
Negative charges should be on more electronegative atoms.
The best resonance form has the most atoms with a formal charge of zero.
Resonance and Formal Charge Together
Resonance structures with formal charges help determine the dominant contributor to the hybrid.
Example: Fulminic Acid (HCNO)
Two structures are possible. By calculating formal charges:
Structure A:
Carbon = 0
Oxygen = −1
Structure B:
Carbon = −1
Oxygen = 0
Structure A is favored because the negative charge is on oxygen, a more electronegative atom. Therefore, A contributes more to the resonance hybrid.
AP Free-Response Practice Examples
HCO3⁻ (Bicarbonate Ion)
This is what College Board had on their site for full credit, but their arrow is the same as the double-headed arrow we discussed. Just please make sure you do not draw an equilibrium arrow by accident.
Question from AP Chemistry 2016 – FRQ #2e:
Carbon is the central atom.
Hydrogen is attached to one of the three oxygen atoms.
Total electrons = 24.
Three resonance structures are possible with different placement of the C=O double bond.
Bond order for the carbon–oxygen bonds not attached to hydrogen is 1.33.
Draw all structures with brackets, charges, and double-headed arrows.
Fulminic Acid (HCNO)
AP Chemistry 2017 – FRQ #2a:
Justify which structure is better based on formal charge.
Electronegativity plays a key role in determining the favored structure.
Negative charges should be placed on more electronegative atoms.
S2Cl2 Molecule
AP Chemistry 2017 – FRQ #1c:
Requires accurate Lewis structure with lone pairs.
Emphasizes correct use of electron pairs and bonding arrangements.
Key Terms
Resonance: The phenomenon where multiple valid Lewis structures represent the same molecule.
Formal Charge: Hypothetical charge on an atom in a molecule based on equal sharing of bonding electrons.
Bond Order: The average number of bonds between atoms in a molecule.
Valence Electrons: Outermost electrons involved in chemical bonding.
Lewis Structure: Diagram showing the arrangement of atoms, bonds, and lone pairs in a molecule.
Resonance Hybrid: The true structure formed by blending all valid resonance structures.
Delocalized Electrons: Electrons that are shared across multiple atoms in a molecule.
Polyatomic Ion: A charged group of covalently bonded atoms.
Electronegativity: An atom’s ability to attract bonding electrons.
By applying resonance and formal charge principles, students can more accurately predict and draw molecular structures, evaluate stability, and better understand chemical behavior on the AP Biology exam.
FAQ
Resonance typically involves delocalized π (pi) electrons, which originate from double or triple bonds. However, certain molecules or ions without traditional double bonds can exhibit resonance if lone pairs are adjacent to atoms capable of forming multiple bonds or if there are empty orbitals that allow electron delocalization. While rare, some species like carbocations or radicals can participate in resonance due to their empty or half-filled p orbitals.
For resonance to occur:
There must be a system of adjacent p orbitals.
Electrons, including lone pairs, must be able to delocalize.
Double or triple bonds are not strictly necessary, but conjugation or orbital overlap is essential.
An example is the allyl cation (C3H5⁺), which has resonance without any traditional double bonds between all carbon atoms.
The resonance hybrid reflects the true electronic structure of a molecule and represents a delocalized electron cloud that spreads out charge and bonding over multiple atoms. This delocalization leads to greater stability than any single Lewis structure can show.
Electron delocalization reduces electrostatic repulsion between negative charges.
The hybrid has lower potential energy, making the molecule more stable.
Bond lengths and strengths in the hybrid are intermediate, which distributes bond strain more evenly.
The molecule benefits from resonance energy, which is the energy difference between the hybrid and the most stable contributing structure.
This stability is why molecules like benzene are more stable than expected based on a single Lewis structure.
Resonance affects molecular geometry because delocalized electrons influence electron cloud distribution, which in turn alters electron repulsion patterns and bond angles predicted by VSEPR theory.
Resonance structures lead to equalization of bond lengths, which affects bond angles.
In molecules like CO3²⁻, all C–O bond lengths are equal, even though some structures show a double bond.
Delocalization causes electron density to spread, often slightly expanding bond angles beyond those expected for localized lone pairs or bonds.
The molecular shape, however, is based on the resonance hybrid, not individual resonance forms.
Geometry is determined by regions of electron density, not the fluctuating positions of electrons in resonance structures.
Yes, resonance can occur in molecules where atoms—especially those in period 3 or beyond—have expanded octets. These atoms can hold more than eight electrons due to available d-orbitals, which allows for more bonding options and delocalization.
Phosphorus, sulfur, and chlorine often participate in resonance with expanded octets.
Example: In SO4²⁻, sulfur forms four S–O bonds with double-bond character due to delocalization.
Formal charge considerations are critical in these cases because electrons may be shifted in different ways across the structure.
Expanded octets allow multiple resonance contributors that wouldn't be possible under the octet rule alone.
The resonance in such molecules enhances stability by distributing electrons over a larger framework.
Resonance can influence a compound’s color by affecting the energy levels of electrons involved in transitions between molecular orbitals. When delocalized electrons are present, the energy required to excite them is altered, shifting the absorption wavelength into the visible light range.
Conjugated systems—molecules with alternating double and single bonds—often display resonance.
The more extensive the delocalization, the smaller the energy gap between molecular orbitals.
This smaller energy gap means light of longer wavelengths (visible light) can be absorbed.
The observed color is the complementary color of the light absorbed.
Example: β-carotene (in carrots) has a long chain of conjugated double bonds, resulting in orange color due to extensive resonance.
Practice Questions
The nitrate ion (NO3−) contains three equivalent nitrogen–oxygen bonds. Using your understanding of resonance, explain how the Lewis structure accounts for this observation. Include bond order in your explanation.
The nitrate ion has three resonance structures, each showing a double bond between nitrogen and a different oxygen atom. These structures do not exist independently but contribute to a single resonance hybrid where the electrons are delocalized across all three N–O bonds. As a result, each bond in the hybrid has the same length and strength, which are intermediate between a single and a double bond. The bond order is calculated as 4 bonding electrons over 3 positions, giving a bond order of 1.33. This fractional bond order explains the equivalent nature of all three N–O bonds in nitrate.
Draw two possible Lewis structures for the fulminic acid molecule (HCNO). Using formal charge calculations, justify which structure is the more stable resonance contributor.
One structure places the negative charge on oxygen and assigns a neutral formal charge to carbon. The other places the negative charge on carbon and gives oxygen a formal charge of zero. Using formal charge calculations, the first structure is preferred because oxygen is more electronegative and better stabilizes the negative charge. Formal charge is calculated by subtracting the number of bonds and lone electrons from the valence electrons of the atom. Since the most stable resonance structure places negative charge on the most electronegative atom, the structure with the negative charge on oxygen contributes more to the resonance hybrid.
