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Valency & Lewis Structure: Valency defines an atom’s combining power. Lewis symbols represent valence electrons. Key points for Lewis structures:
- Total valence electrons are summed.
- Anions add electrons; cations subtract.
- Central atom: typically less electronegative.
Octet Rule: Atoms strive for a full valence shell (8 electrons). Exceptions include H (2 electrons), BF3, AlCl3, PF5, SF6, and inert gas compounds.
Ionic Bond: Formed by full electron transfer, typically between metals and non-metals, resulting in cations and anions. Properties include:
- High melting points and crystalline structures.
- Conduct electricity when dissolved or molten.
Covalent Bond: Atoms share electrons (bond pairs). Characteristics include:
- Generally low melting and boiling points.
- Poor electrical conductors, soluble in non-polar solvents.
Co-ordinate Covalent Bond: Formed when one atom donates both electrons in the shared pair. Examples include ammonium ion (NH4+).
Metallic Bond: Electrons form a sea around positively charged kernels, explaining properties like conductivity and malleability.
Dipole Moment & Molecular Geometry: Measure of bond polarity. Calculated as $\boldsymbol{ ext{μ} = e imes d}$ (where $e$ is charge and $d$ is distance).
Hybridization: Mixing of atomic orbitals to form equivalent hybrid orbitals (e.g. sp, sp², sp³).

Formed when H is bonded to F, O, or N, creating strong dipoles.
Affects properties like boiling point, solubility, and viscosity:
- Water’s unique properties stem from hydrogen bonding (higher density than ice).

States molecular geometry is determined by minimizing electron pair repulsions. Examples include:
- Linear: 2 electron pairs (e.g., BeCl2)
- Trigonal Planar: 3 electron pairs (e.g., BF3)
- Tetrahedral: 4 electron pairs (e.g., CH4)

VSEPR does not account for polar covalent bonds or transition metal complexes.
Valence Bond Theory does not explain magnetism or consider lone pairs in overlaps.