Chemical Bonding
World A — Chemistry Foundations
R06.01 Octet Rule
- Atoms seek 8 valence electrons.
- Drives chemical reactivity.
- Attaining noble gas stability.
Na
Cl
R06.02 Ionic Bonding
- Complete transfer of electrons.
- Electrostatic attraction (+/-).
- Forms crystalline lattices.
R06.03 Covalent Bonding
- Mutual sharing of electrons.
- Between non-metals.
- Forms distinct molecules.
R06.04 Metallic Bonding
- Positively charged metal ions.
- Delocalized "sea of electrons".
- Explains conductivity & malleability.
C
R06.05 Lewis Structures
- Visualizes valence electrons as dots.
- Shows lone pairs and bond pairs.
- Predicts connectivity in molecules.
R06.06 VBT & Hybridization
- Valence Bond Theory (overlapping orbitals).
- Atomic orbitals mix (e.g., sp³).
- Creates equivalent hybrid orbitals.
R06.07 VSEPR Theory
- Valence Shell Electron Pair Repulsion.
- Electron pairs push apart maximally.
- Determines 3D molecular geometry.
R06.08 Molecular Orbital Theory
- Atomic orbitals combine into MOs.
- Bonding (lower E) & Antibonding (higher E).
- Explains paramagnetism of O₂.
R06.09 Hydrogen Bonding
- Strong dipole-dipole interaction.
- H bonded to highly electronegative N, O, F.
- Gives water its unique high boiling point.
→
R06.10 Coordinate Bond
- Dative covalent bond.
- One atom provides BOTH shared electrons.
- Common in complex metal ions.
R06.11 Resonance
- Multiple valid Lewis structures.
- Actual structure is a hybrid average.
- Delocalized pi-electrons (e.g., Benzene).
δ+
δ-
R06.12 Dipole Moment
- Unequal sharing of electrons.
- Creates partial charges (polarity).
- Vector sum determines overall molecular polarity.