Resonance: Patterns, Charges, and Allyl Carbocations

Resonance and Bonding: Key Concepts, Patterns, and Examples

• Goal of resonance study

  • Resonance helps explain molecule stability and reactivity by showing how electrons (not nuclei) are distributed in overlapping orbitals (pi bonds and lone pairs).
  • Individual resonance structures are not real; the real molecule is the resonance hybrid of these structures.
  • Resonance structures are not constitutional isomers; they do not change connectivity, only electron placement.
  • Use curved arrows to represent electron movement (curved arrows indicate the flow of electrons, not nuclei).
  • The resonance concept is foundational for predicting reactivity, acid-base behavior, and intermolecular interactions.

• Quick recap: formal charges and valence patterns

  • Formal charge (FC) is a bookkeeping method to assess electron ownership:
  • FC = V - N{nonbonding} - \frac{B}{2} where V = valence electrons for the atom, N{nonbonding} = electrons in lone pairs, B = bonding electrons (shared).
  • Relative valence counts used in class examples:
  • Oxygen: typically 6 valence electrons; common patterns include neutral O, O⁻ (alkoxide-like), and O⁺ (e.g., oxonium-like) depending on lone pairs and bonds.
  • Carbon: valence = 4. Formal charge examples:
    • Carbocation (C⁺): carbon with three sigma bonds and no lone pair; sp² hybridized; typically trigonal planar.
    • Carbanion (C⁻): carbon with a lone pair; typically sp³ hybridized and tetrahedral around carbon; formal charge of −1.
    • Neutral carbon: typically has four substituents (or two substituents plus a lone pair in special cases) with no formal charge; can be sp³, sp², or sp depending on bonding.
  • Nitrogen: valence = 5. Can be N⁻ (amide-like or nitride contexts), neutral amine, or N⁺ (e.g., ammonium). Negatives are highly nucleophilic and basic; positives lack a lone pair.
  • Important practical takeaway: changes in formal charge accompany changes in electron ownership and/or bonding patterns; resonance often distributes charge across multiple atoms.

• Oxygen patterns and context for resonance

  • Oxygen commonly participates in:
  • Neutral patterns (e.g., carbonyl O in C=O, alcohol O–H) with two lone pairs and two bonds.
  • Negative patterns (e.g., alkoxide O⁻) with three lone pairs and one bond, contributing to nucleophilicity.
  • Positive patterns (e.g., oxonium-like O⁺) when oxygen bears a formal positive charge due to bonding arrangements.
  • In resonance, lone pairs on oxygen and adjacent pi bonds participate in delocalization, shifting electron density without breaking sigma bonds.
  • Induction vs resonance (conceptual link): electronegative atoms (like O) pull electron density toward themselves, creating partial negative/positive regions that influence reactivity; resonance delocalization further distributes charge.

• Carbon patterns and key species

  • Carbocation (C⁺):
  • Valence electrons owned by carbon reduces from four to three; carbon is sp² hybridized; trigonal planar geometry.
  • Can be primary (one carbon substituent), secondary (two substituents), or tertiary (three substituents); the number of substituents affects stability (tertiary > secondary > primary).
  • Carbanion (C⁻):
  • Carbon with a lone pair; charged center can act as a strong base/nucleophile; typically sp³ hybridized with tetrahedral geometry.
  • Neutral carbon as shown in alkanes/alkenes can be sp³ (single bonds) or sp² (in alkenes) depending on bonding.
  • Hydrogens: when counting hydrogens in resonance-related considerations, remember that hydrogens are sigma-bonded and generally do not participate in forming or breaking sigma bonds during resonance.

• Nitrogen patterns

  • Neutral nitrogen (amines, e.g., R₃N) has valence 5 and typically bears a lone pair.
  • Negative nitrogen (amides or amide-like anions) is highly nucleophilic and basic; it has a filled lone pair and can participate in resonance with adjacent pi systems or carbonyls in some contexts.
  • Positive nitrogen (e.g., ammonium, R₄N⁺) bears four sigma bonds and has no lone pair; it is not nucleophilic in that state.
  • In resonance contexts, nitrogen’s lone pairs can delocalize but must satisfy octet rules for second-row elements; violations indicate incorrect resonance forms.

• Allyl carbocation: a canonical resonance pattern

  • Definition: a carbocation directly neighboring a double bond (allylic position).
  • Key structural features:
  • Three carbons involved with a delocalized π system across C1–C2–C3; all three carbons are sp² hybridized (bound to three groups each).
  • The π bond is delocalized: the double bond can shift among C1=C2 or C2=C3 without moving hydrogens, changing where the positive charge appears.
  • Resonance forms distribute the positive charge between adjacent carbons (e.g., + on C1 in one form, + on C3 in another) while keeping overall charge constant.
  • Conceptual consequence: resonance stabilizes allyl carbocations by delocalizing the charge; this is a classic example of why conjugation stabilizes reactive intermediates.
  • Important visualization point: the resonance hybrid is more stable than any single resonance structure due to charge delocalization.

• How to draw resonance structures (rules and workflow)

  • Rule 1: Do not delocalize sigma bonds. Only pi bonds and lone pairs participate in resonance; breaking a sigma bond would imply a chemical reaction, not resonance.
  • Rule 2: Never exceed an octet for second-row elements (C, N, O, F). Some early-stage resonance forms violate octets; those forms are invalid.
  • Rule 3: Use curved arrows to indicate electron movement. A double-headed curved arrow represents movement of two electrons; a single-headed curved arrow represents movement of one electron.
  • Rule 4: Arrows originate from electron-rich sites (lone pairs or pi bonds) and terminate at electron-deficient sites (positive charges or empty orbitals).
  • Rule 5: In resonance, hydrogens and sigma-bonded connections do not move. The changes involve the redistribution of electrons (lone pairs and π electrons) only.
  • Practical approach:
  • Start with a skeletal structure, then add hydrogens as needed to assess octets.
  • Add resonance forms by shifting electrons with curved arrows, then re-check formal charges on all atoms.
  • Remember the total charge and the octet constraint must be consistent across all valid resonance forms.

• Pattern recognition: five general resonance patterns

  • Pattern A: Lone pair adjacent to a pi bond (or a positive center). The lone pair can donate into the π system, creating a new π bond and redistributing charge.
  • Pattern B: Allylic lone pair adjacent to a double bond (allylic system) can delocalize into the π system.
  • Pattern C: Lone pair adjacent to a carbocation (lone-pair donation to stabilize a positive center).
  • Pattern D: Pi bond between two atoms with different electronegativities (conjugated systems) enables delocalization of electrons across the bond; the difference in electronegativity helps determine the direction of electron density shift.
  • Pattern E: Conjugated or alternating double bonds (polyene systems) allow multiple valid resonance contributors. When multiple double bonds are conjugated, electrons can delocalize across the entire system, and multiple resonance forms exist; each contributor has a portion of the overall electron density.

• Practical implications: induction, conjugation, and reactivity

  • Induction vs resonance: induction refers to electron density shifts due to electronegativity (a static effect), while resonance refers to actual delocalization of electrons across multiple atoms in a conjugated system.
  • In resonance-stabilized systems, partial charges are spread over several atoms, reducing localized charge density and stabilizing the molecule.
  • Understanding resonance helps predict sites of nucleophilic attack, electrophilic attack, and overall acidity/basicity in acid-base chemistry.
  • For carbonyl-containing systems, resonance can place partial negative character on the oxygen and partial positive character on the carbonyl carbon, affecting reactivity and intermolecular interactions.

• Important visual and conceptual takeaways

  • Resonance structures are not interchangeable real structures; they are a set of contributing pictures whose average is the actual structure (the resonance hybrid).
  • The resonance can help you identify reactive sites: for example, allyl carbocations have delocalized positive charge across the allylic carbons; nucleophiles will attack electron-poor centers, while electrophiles seek electron-rich sites.
  • The idea of “alkenes and carbonyls” in resonance gives rise to common reactivity patterns you will encounter in subsequent topics (e.g., acid-base chemistry, carbonyl chemistry, and conjugated systems).
  • When drawing resonance forms for practice, consider common mistakes to avoid:
  • Breaking a sigma bond in a resonance form.
  • Violating the octet rule for second-row elements.
  • Forgetting to redistribute formal charges appropriately after electron movement.

• A quick, concrete worked example: resonance in the allyl cation (outline)

  • Start with: CH2=CH–CH2⁺ (allyl cation) – carbocation on terminal carbon, π bond between C1=C2.
  • Step 1: Move the π electrons from C1=C2 to form a new π bond between C2 and C3; this shifts the positive charge to C1.
  • Step 2: The resulting structure places the positive charge on C1 with a remaining π bond between C2 and C3.
  • Step 3: The second canonical form can reposition the π bond back to C1=C2 with the positive charge shifted to C3.
  • Result: two main resonance contributors with the charge distributed across C1 and C3; the actual molecule is the resonance hybrid, showing stabilization through delocalization across all three carbons.
  • Hybridization in this system: all three carbons are sp² hybridized in these forms; each carbon is involved in a planarity that accommodates the delocalized π system.

• Connections to broader concepts and exam-oriented tips

  • Resonance is foundational for understanding conjugated systems, carbonyl chemistry, and acid-base equilibria.
  • Expect to identify potential resonance structures by spotting patterns of lone pairs next to pi bonds, allylic cations, and conjugated double bonds.
  • Practice recognizing which electrons are moving (lone pairs vs. π electrons) and which bonds or lone pairs remain unchanged (sigma bonds generally do not move in resonance).
  • When preparing for assessments, build a set of practice resonance problems from examples in your book or lecture slides and verify that all resulting resonance forms obey the octet rule and charge conservation.

• Quick practical reminders for study sessions

  • Use the hydrogen labeling primarily to keep track of octets but remember that hydrogens do not participate in the electron-shifting events in resonance.
  • Distinguish resonance from actual chemical reactions; do not interpret resonance arrows as reaction steps.
  • The resonance framework will connect to later topics like aromaticity, carbonyl chemistry, and reaction mechanisms, so a solid grasp here will pay off in multiple chapters.

• Summary takeaway

  • Resonance describes how electrons are distributed across conjugated systems and lone pairs, creating multiple valid contributing structures that collectively stabilize the molecule. By applying the five recognition patterns, honoring the octet rule, and using curved arrows correctly, you can predict where charge resides and how reactivity may proceed in many organic molecules.