Chemical Bonding - Quick Notes
Bonding overview
- Atoms are most stable with a full outer electron shell; bonding forms to increase stability.
- Noble gases have full outer shells and typically do not form bonds.
- Warm-up: electron configurations (examples): \text{Ne} = 2,8\qquad \text{Ar} = 2,8,8
- Metals: outer shell electrons are delocalised and free to move around metal cations as a “sea” of electrons.
- Result: mobile electrons allow conductivity and a lattice of positive metal ions.
Ionic Bonding
- Formed between metals (cations) and non-metals (anions) via transfer of electrons.
- Ions are fixed in a solid 3D lattice held by strong electrostatic attraction; overall neutral ionic compound.
- Crystalline lattice structure helps limit repulsion between like charges.
Covalent Bonding
- Occurs between non-metals when electrons are shared to achieve a more stable electron configuration.
- Molecules are formed when atoms share electrons.
- Hydrogen sharing requires 2 electrons, other atoms typically seek 8 (octet).
Simple Covalent Molecules
- Examples of covalent molecules: \text{H}2\text{O}, \text{CH}4, \text{CO}2, \text{H}2, \text{NH}_3
Bonding Strength
- Covalent bonds inside molecules are strong primary bonds.
- Between covalently bonded molecules, weaker secondary forces occur; these determine whether a substance is a gas or a liquid at room temperature.
Checkpoint: Bond Type by Partners
- Metal + Metal: Metallic bonding
- Metal + Non-metal: Ionic bonding
- Non-metal + Non-metal: Covalent bonding
Compound Bonding Examples
- CH₄ (methane) → Covalent
- MgCl₂ (magnesium chloride) → Ionic
- C₆H₁₂O₆ (glucose) → Covalent
- NaOH (sodium hydroxide) → Ionic
- Co (cobalt) → Metallic
- NH₃ (ammonia) → Covalent
- Fe₂O₃ (iron(III) oxide) → Ionic
- V (vanadium) → Metallic
- CH₃COOH (ethanoic acid) → Covalent