Chapter 3: Chemical Bonding

Chapter 3: Chemical Bonding

Types of Chemical Bonds

- Ionic bonding: Electrostatic attraction between oppositely charged ions (metal + non-metal).

- Covalent bonding: Shared pair of electrons between non-metal atoms.

- Dative (coordinate) bonding: One atom donates both electrons in a shared pair.

Ionic Bonding

- Formed when electrons are transferred from a metal to a non-metal.

- Results in the formation of positive (cation) and negative (anion) ions.

- Strong electrostatic forces → high melting/boiling points.

- Example: NaCl

Covalent Bonding

- Atoms share electrons to achieve a noble gas configuration.

- Can be single, double, or triple bonds.

- Molecules may be polar or non-polar depending on the electronegativity difference.

- Example: H₂O, O₂, CO₂

Dative Bonding

- A lone pair from one atom is donated to an empty orbital of another.

- Common in complex ions and molecules like NH₄⁺ and H₃O⁺.

Bond Polarity and Electronegativity

- Electronegativity: Ability of an atom to attract bonding electrons.

- Greater difference → more polar bond.

- Polar molecules have uneven charge distribution → dipole moments.

Shapes of Molecules (VSEPR Theory)

- Electron pairs repel → determine molecular shape.

- Lone pairs repel more than bonding pairs.

- Common shapes:

  - Linear (e.g., CO₂)

  - Trigonal planar (e.g., BF₃)

  - Tetrahedral (e.g., CH₄)

  - Pyramidal (e.g., NH₃)

  - Bent (e.g., H₂O)

Intermolecular Forces

- Van der Waals forces: Weak, present in all molecules.

- Permanent dipole-dipole interactions: Between polar molecules.

- Hydrogen bonding: Strongest; occurs when H is bonded to N, O, or F.

- Affects boiling/melting points and solubility.

Metallic Bonding

- Positive metal ions in a sea of delocalized electrons.

- Explains the conductivity, malleability, and ductility of metals.