Prepare for AP Chemistry Topic 2.6 with this free practice quiz. Test your understanding of resonance structures and use formal charge calculations to determine the most stable molecular arrangement.
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The Pre-Quiz Review
Resonance Structures
Resonance is used when a single Lewis structure cannot adequately represent the actual distribution of electrons in a species. The different valid Lewis structures are called resonance structures or resonance contributors.
- Resonance structures have the same arrangement of atoms but different arrangements of electrons.
- Only electrons move; nuclei and σ-bonded atoms do not move.
- Resonance commonly involves π electrons, lone pairs, and adjacent empty p orbitals.
- The actual species is a resonance hybrid, not a molecule rapidly switching between structures.
- All resonance structures represent the same species and therefore have the same atom connectivity.
- Resonance does not mean that atoms physically move back and forth between structures.
For AP Chemistry, first draw a correct Lewis structure before deciding whether resonance exists.
A common pattern is:
- A π bond next to a lone pair can produce resonance.
- A π bond next to a positive charge can produce resonance.
- A π bond next to an atom that can accommodate a charge can produce resonance.
- In conjugated systems, adjacent p orbitals allow electrons to become delocalized across several atoms.
For example, nitrate, NO₃⁻, has three equivalent resonance structures. In each contributor, one N–O bond is represented as a double bond and the other two as single bonds, with formal charges assigned appropriately. The real nitrate ion has three equivalent N–O bonds rather than one permanently double-bonded N–O bond.
Carbonate, CO₃²⁻, similarly has three equivalent resonance contributors.
Rules for Drawing Resonance Structures
When converting one resonance structure into another:
- Keep the atoms in exactly the same positions.
- Move only electrons.
- Move lone pairs or π electrons, not σ bonds.
- Maintain the total number of valence electrons.
- Maintain the overall charge of the species.
- Check formal charges after moving the electrons.
A curved arrow shows electron movement. The arrow must begin at an electron source, such as a lone pair or bond, and point toward where those electrons will form a bond or become a lone pair.
For AP MCQs, an option that changes the position of an atom or breaks a σ bond is not simply another resonance structure. It represents a different structure or potentially a different species.
Equivalent vs. Nonequivalent Resonance Structures
Equivalent resonance structures have the same energy and equivalent charge distributions. They contribute equally to the resonance hybrid.
For example, the three major contributors of NO₃⁻ are equivalent because the three oxygen atoms occupy equivalent positions.
Nonequivalent resonance structures have different energies and therefore do not contribute equally.
When deciding which contributor is more important, examine:
- Formal charge.
- Charge separation.
- Whether atoms have complete octets.
- Placement of formal charges.
- Electronegativity of atoms bearing charge.
A resonance contributor with complete octets and smaller or fewer formal charges is generally more important.
Delocalized Electrons
Delocalized electrons are electrons that are spread over multiple atoms rather than being confined to one bond or one atom.
Resonance is a Lewis-structure representation of electron delocalization.
A useful distinction:
- Localized electrons: associated primarily with one atom or one bond.
- Delocalized electrons: distributed across several adjacent atoms.
π electrons are especially important in delocalization because overlapping p orbitals can allow electrons to be distributed across a larger region.
In a resonance hybrid, electron density is distributed over the atoms involved in resonance. A bond represented as a single bond in one contributor and a double bond in another may have an intermediate bond order in the actual species.
For example, in NO₃⁻, all three N–O bonds are equivalent in the resonance hybrid. The Lewis structures use one N=O bond at a time for bookkeeping, but the actual electron density is delocalized over all three N–O bonds.
Resonance and Bond Properties
Delocalization can affect measurable properties such as:
- Bond length.
- Bond strength.
- Stability.
- Charge distribution.
When electrons are delocalized over multiple atoms, the resulting species is often more stable than a comparable structure in which the electrons are localized.
For equivalent resonance structures, equivalent bonds have equivalent average properties.
A bond that is represented as single in one contributor and double in another may have a bond order between 1 and 2 in the resonance hybrid. For AP Chemistry, focus on the qualitative consequence: greater resonance delocalization generally produces greater stabilization and can make equivalent bonds have intermediate characteristics.
Formal Charge
Formal charge is a bookkeeping method used to determine how valence electrons are assigned in a Lewis structure.
The formula is:
Formal charge = valence electrons − nonbonding electrons − ½(bonding electrons)
For a typical atom in a Lewis structure:
- Nonbonding electrons are the electrons in lone pairs.
- Bonding electrons are the electrons shared in bonds.
- Half of the bonding electrons are assigned to the atom for formal-charge purposes.
An alternative form that is often faster:
Formal charge = valence electrons − dots − bonds
Here, “dots” means nonbonding electrons and “bonds” means the number of bond lines around the atom.
For example, oxygen normally has 6 valence electrons.
- O with three lone pairs and one single bond: formal charge = −1.
- O with two lone pairs and two single bonds: formal charge = 0.
- O with one lone pair and three single bonds: formal charge = +1.
Nitrogen normally has 5 valence electrons.
- N with three bonds and one lone pair: formal charge = 0.
- N with four bonds and no lone pair: formal charge = +1.
- N with two bonds and two lone pairs: formal charge = −1.
Using Formal Charge in Lewis Structures
The sum of all formal charges must equal the overall charge of the species.
For a neutral molecule:
Total formal charge = 0
For an ion:
Total formal charge = ion charge
This is one of the fastest ways to detect an incorrect Lewis structure in an AP MCQ.
If a proposed structure for SO₄²⁻ has formal charges that add to −1, it cannot represent sulfate correctly.
Formal charge is not the same as actual charge density. It is a formal electron-accounting system based on an equal division of bonding electrons.
Do not assume that an atom with a formal charge of −1 literally possesses a full unit of negative charge localized entirely on that atom. Resonance and electronegativity affect the actual electron distribution.
Formal Charge and Resonance
Formal charge is especially useful when comparing resonance contributors.
Consider a resonance structure in which an atom has an incomplete octet versus one in which all relevant atoms have complete octets. The complete-octet structure is generally favored, provided no serious charge or bonding problem is introduced.
For second-row elements such as C, N, O, and F, an incomplete octet is generally unfavorable.
Carbon, nitrogen, oxygen, and fluorine are second-period elements and cannot expand their valence shells beyond an octet in standard AP Chemistry Lewis-structure reasoning.
Structural Stability
When comparing possible Lewis structures or resonance contributors, use a consistent hierarchy.
1. Satisfy the octet rule when possible
Structures in which major second-row atoms have complete octets are generally favored.
Hydrogen is an exception: it follows a duet rule and can have only two electrons around it.
2. Minimize formal charges
Structures with fewer formal charges are generally more stable than structures with extensive charge separation, assuming comparable octet completion.
A structure with several positive and negative formal charges is usually less favorable than one with fewer formal charges.
3. Consider charge separation
Charge separation means placing opposite formal charges on different atoms.
A contributor with substantial charge separation is generally less stable than one with fewer separated charges.
For example, a neutral structure with all atoms having formal charge 0 is generally preferred over an alternative neutral structure containing +1 and −1 formal charges, if both satisfy octet requirements.
4. Place negative charge on more electronegative atoms
When formal charges cannot be avoided, negative formal charge is generally more favorable on the more electronegative atom.
For common AP Chemistry elements:
F > O > N > C
Thus, a negative formal charge is generally more favorable on oxygen than on carbon.
Likewise, a positive formal charge is generally more favorable on a less electronegative atom than on a highly electronegative atom.
This principle is particularly useful when identifying the major resonance contributor among nonequivalent structures.
Formal Charge and Oxidation State Are Different
Formal charge and oxidation state use different electron-assignment rules.
Formal charge assigns bonding electrons equally between bonded atoms.
Oxidation state assigns bonding electrons to the more electronegative atom.
Therefore, the same atom can have a formal charge of 0 but a nonzero oxidation state.
Do not substitute oxidation-state calculations for formal-charge calculations when evaluating Lewis structures or resonance contributors.
Resonance, Charge, and Electron Movement
A useful AP strategy is to track electrons rather than memorizing individual resonance patterns.
If a lone pair forms a π bond, the atom supplying the lone pair may become more positively charged.
If a π bond becomes a lone pair, the atom receiving those electrons may become more negatively charged.
For example, when a lone pair adjacent to a π bond participates in resonance:
- The lone pair can form a new π bond.
- The existing π bond must move elsewhere.
- The total number of electrons remains unchanged.
- Formal charges redistribute.
The overall charge of the species never changes simply because you draw a different resonance contributor.
Recognizing Resonance in AP MCQs
When an AP question presents several structures, check them systematically.
First, compare atom connectivity. If the atoms are connected differently, the structures are not resonance structures.
Next, compare the location of electrons. Valid resonance structures differ only in the placement of electrons.
Then calculate or inspect formal charges.
Finally, determine whether the structures are equivalent or nonequivalent.
A particularly common AP question asks which diagram represents resonance. The correct choice will show identical atomic connectivity with different π bonds, lone pairs, or formal charges.
Another common question asks which resonance contributor is most stable. Look for complete octets, fewer formal charges, less charge separation, and appropriate placement of charge based on electronegativity.
Resonance and Molecular Stability
Resonance stabilization occurs because electron density can be distributed over multiple atoms.
Greater effective delocalization generally lowers the energy of the actual resonance hybrid relative to a hypothetical localized structure.
The resonance hybrid is therefore typically more stable than any single resonance contributor.
However, do not say that the molecule “switches” between resonance structures. The contributors are drawings used to represent one real electronic structure.
For AP Chemistry, distinguish between:
- Resonance structures: individual Lewis representations.
- Resonance hybrid: the actual delocalized electronic structure represented collectively by the contributors.
- Delocalized electrons: electrons distributed across multiple atoms.
- Formal charge: electron bookkeeping used to evaluate Lewis structures.
High-Value Formal Charge Patterns
Memorizing common patterns saves time on MCQs.
| Atom | Typical neutral arrangement |
|---|---|
| H | 1 bond |
| C | 4 bonds |
| N | 3 bonds + 1 lone pair |
| O | 2 bonds + 2 lone pairs |
| F | 1 bond + 3 lone pairs |
Useful charge patterns:
- N with 4 bonds → +1 formal charge.
- N with 2 bonds and 2 lone pairs → −1 formal charge.
- O with 3 bonds and 1 lone pair → +1 formal charge.
- O with 1 bond and 3 lone pairs → −1 formal charge.
- Halogen with 1 bond and 3 lone pairs → 0 formal charge.
These patterns are especially useful when evaluating nitrate, nitrite, carbonate, ozone, sulfate, and other polyatomic species.
AP-Level Reasoning
When a question asks why a resonance-stabilized structure is favored, the strongest explanation usually involves electron delocalization and lower electronic energy.
When a question asks why one resonance contributor is more important, evaluate its formal charges, octets, charge separation, and electronegativity.
When a question asks about equivalent bonds, remember that resonance can make bonds equivalent even though individual Lewis contributors draw them differently.
When a question asks whether two structures are resonance forms, ignore the apparent visual differences and compare the positions of the nuclei first. If the atom framework changes, they are not resonance structures.
When a question gives a proposed Lewis structure, calculate formal charges rather than relying only on appearance. The total formal charge provides an immediate consistency check.
Finally, remember that resonance is fundamentally about electron delocalization. Lewis structures are representations; the actual molecule or ion has one electronic structure whose electron density may be distributed across several atoms.