chemistry/chemical_bonding/intermediate.md

🔗 Chemical Bonding — Intermediate#

Scope — Bonding models beyond simple labels: Lewis structures, VSEPR geometry, hybridization, resonance, and how these explain molecular shape and reactivity.

Key concepts#

• Lewis structures — electron-dot diagrams that show valence electrons, lone pairs, and bonding pairs; used to predict formal charges.
• VSEPR theory — electron pair repulsion determines molecular geometry; lone pairs alter bond angles.
• Hybridization — mixing of atomic orbitals (e.g., sp, sp², sp³) to explain observed bond angles and molecular shapes.
• Resonance — delocalization of electrons across multiple structures stabilizes molecules and affects reactivity.

Seed Q&A triads#

• Q: How do you decide the best Lewis structure when multiple resonance forms exist?
  A: Favor structures with full octets, minimal formal charges, and negative formal charges on more electronegative atoms.

• Q: Why is methane tetrahedral while ethene is planar?
  A: Methane uses sp³ hybridization giving tetrahedral geometry; ethene uses sp² hybridization with a remaining p orbital forming a π bond, producing planarity.

• Q: How does VSEPR predict the shape of ammonia (NH₃) versus water (H₂O)?
  A: Both have tetrahedral electron-pair geometry; ammonia has one lone pair so trigonal pyramidal shape, water has two lone pairs so bent shape with smaller bond angle.

Short exercises#

• Draw Lewis structures and predict geometry for: SO₂, NO₃⁻, NH₄⁺.
• Explain how resonance in NO₃⁻ leads to equal N–O bond lengths experimentally observed.