Overview
This lecture introduced the concept of chemical bonding, focusing on why atoms form bonds, types of chemical bonds, limitations of the octet theory, and the basics of valence bond theory (VBT).
Why Do Atoms Form Bonds?
- Atoms form bonds to minimize their energy and gain stability by achieving duplet (2 electrons) or octet (8 electrons) configurations.
- Inert gases like helium, neon, and argon do not form bonds as they already have stable configurations.
- Hydrogen, oxygen, and nitrogen form diatomic molecules (H₂, O₂, N₂) to complete their duplet/octet.
Potential Energy and Bond Formation
- As two atoms approach each other, their potential energy decreases until a minimum (most stable) point (bond formation).
- Over-compressing atoms increases potential energy and causes instability, possibly breaking the bond (bond dissociation).
- The bond forms at an optimal inter-nuclear distance, R₀, where stability is maximized.
Electron Pairing and Magnetic Effects
- Electrons in bonding orbitals must have opposite spins (anti-spin) to reduce repulsion and allow bond formation (Pauli exclusion principle).
- Spinning electrons create tiny magnets; only pairing opposite spins creates a stable bond.
Types of Chemical Bonds
- Ionic bond: Electron(s) completely transferred from one atom (metal) to another (non-metal), e.g., NaCl.
- Covalent bond: Electrons are shared between atoms, e.g., H₂, O₂, N₂.
- Coordinate (dative) bond: Both shared electrons come from one atom, e.g., NH₄⁺.
- Other interactions: Hydrogen bonds and Van der Waals (London) forces are weaker and not true bonds.
Lewis (Leavis) Dot Structures
- Lewis structures represent bonding and lone pair electrons as dots/crosses, completing each atom's duplet or octet.
- Practice drawing Lewis structures for F₂, Cl₂, HF, NH₃, O₃, CO₂, SO₂, H₂O₂, NO, NO₂, CO₃²⁻, PCl₅, etc.
- Some molecules show limitations in Lewis theory (odd electrons, expanded or incomplete octets).
Limitations of Octet Theory
- Cannot explain stability of odd-electron species (e.g., NO).
- Fails for electron-deficient molecules (e.g., BCl₃, AlCl₃).
- Cannot explain expanded octets (e.g., PCl₅, SO₃, HClO₄).
- Does not predict molecule shapes.
Introduction to Valence Bond Theory (VBT)
- VBT describes covalent bond formation as the overlap of valence orbitals containing unpaired electrons.
- Only orbitals with unpaired electrons and opposite spins participate in bonding.
- Greater orbital overlap leads to stronger bonds; poor orientation results in weaker (π) bonds.
- Sigma (σ) bonds have head-to-head overlap (strong), while pi (π) bonds have side-to-side overlap (weaker).
Key Terms & Definitions
- Octet Rule — Atoms tend to form bonds to have eight electrons in their valence shell.
- Duplet — A stable configuration of two electrons (as in helium).
- Lewis Structure — Diagram showing bonding and lone pairs as dots.
- Coordinate Bond — Covalent bond where both electrons come from the same atom.
- Bond Dissociation — Breaking of a chemical bond via energy input.
- Sigma (σ) Bond — Strong covalent bond with head-to-head orbital overlap.
- Pi (π) Bond — Weaker covalent bond with side-to-side orbital overlap.
- Valence Bond Theory (VBT) — Theory explaining covalent bonding through orbital overlap.
- Expanded Octet — More than 8 electrons in the outer shell (common in period 3 and beyond).
Action Items / Next Steps
- Draw Lewis structures for carbonate, NO₂, PCl₅, and any molecules not completed in class.
- Review limitations of the octet theory.
- Read textbook sections on VBT and orbital overlaps.