Chemical bonding (C2.2)

This topic is on every Gateway A Chemistry paper. Expect a dot-and-cross diagram, a 6-mark question on properties of giant ionic vs simple molecular substances, and a short answer on metallic bonding.

Why atoms bond

Atoms react to achieve a full outer shell (the noble-gas electron configuration of group 0). Three common ways:

1. Ionic — electrons are TRANSFERRED from a metal to a non-metal.
2. Covalent — electrons are SHARED between non-metals.
3. Metallic — outer electrons are DELOCALISED in a sea around metal cations.

Ionic bonding

Forms between a metal and a non-metal. The metal atom gives up one or more outer electrons to become a positive cation; the non-metal atom gains those electrons to become a negative anion.

Example — sodium chloride (NaCl):

• Na (1 outer electron) → Na⁺ + e⁻
• Cl (7 outer electrons) + e⁻ → Cl⁻
• Na⁺ and Cl⁻ are held together by strong electrostatic forces in a giant ionic lattice.

Properties of ionic compounds

• High melting and boiling points (lots of strong electrostatic bonds to break).
• Conduct electricity when molten or dissolved (ions free to move) — NOT as a solid.
• Often soluble in water; usually crystalline; brittle.

⚠ Common error: writing that NaCl conducts when solid. It does not — ions are locked in place.

Covalent bonding

Forms between two non-metals. The atoms share pairs of electrons.

• A single bond = one shared pair (e.g. H–H).
• A double bond = two shared pairs (e.g. O=O).
• A triple bond = three shared pairs (e.g. N≡N).

Simple molecular substances (e.g. H₂O, CO₂, CH₄, Cl₂)

• Strong covalent bonds INSIDE the molecule.
• Weak intermolecular forces BETWEEN molecules.
• → Low melting and boiling points (only weak forces to break).
• → Don't conduct electricity (no ions or free electrons).

Giant covalent substances (e.g. diamond, graphite, silicon dioxide)

• Lots of atoms covalently bonded in a 3-D network.
• Very high melting points (millions of strong covalent bonds to break).
• Generally don't conduct (exception: graphite, see below).

Carbon allotropes (split out into C2.3)

• Diamond — each C bonded to 4 others; very hard; doesn't conduct (no free electrons).
• Graphite — each C bonded to 3 others, in flat layers; layers slip → soft, slippery; one delocalised electron per carbon → conducts.
• Graphene — single layer of graphite; very strong, conducts.
• Fullerenes (e.g. buckminsterfullerene C₆₀) — molecular cages; nanotubes are cylindrical fullerenes.

Metallic bonding

Metals are giant lattices of cations surrounded by a "sea" of delocalised electrons.

• The delocalised electrons are no longer attached to any one atom — they're free to drift.
• The strong attraction between cations and electron sea is the metallic bond.

Properties of metals

• Conduct electricity (delocalised electrons move).
• Conduct heat (delocalised electrons transfer kinetic energy fast).
• Malleable and ductile (layers of cations can slide without breaking the bond — the electron sea repositions).
• High melting points (strong electrostatic forces).

Alloys

A mix of a metal with another element (often another metal). The added atoms are different sizes and disrupt the regular lattice — layers cannot slide as easily, so alloys are HARDER than pure metals.

Common Gateway-paper mistakes

1. Saying NaCl conducts when solid (it doesn't — only molten or aqueous).
2. Confusing simple molecular and giant covalent ("CO₂ is a giant covalent" — wrong).
3. Forgetting that diamond doesn't conduct because each C uses ALL FOUR outer electrons in bonds.
4. Saying metallic bonds are between ions and ions (they're between cations and the electron sea).
5. Writing that graphite is soft because the bonds are weak — the bonds within layers are strong; it's the intermolecular forces between layers that are weak.