Chemical bonding explains how atoms join together. The three main types — ionic, covalent, and metallic — produce structures with very different properties.
Formed between metals and non-metals. Metals lose electrons → positive ions (cations). Non-metals gain electrons → negative ions (anions). Electrostatic attraction between oppositely charged ions. Dot-and-cross diagrams show electron transfer. Examples: NaCl (Na⁺ Cl⁻), MgO (Mg²⁺ O²⁻), CaCl₂. Giant ionic lattice: high melting points, conduct electricity when molten/dissolved (ions free to move), do not conduct as solids.
Formed between non-metals. Shared pairs of electrons. Single bond (1 pair), double bond (2 pairs). Simple molecular structures (H₂O, CO₂, CH₄): low melting/boiling points (weak intermolecular forces), do not conduct electricity. Giant covalent structures (diamond, graphite, SiO₂): very high melting points. Diamond: 4 bonds per carbon, hard. Graphite: layers, delocalised electrons → conducts electricity, slippery.
Metal atoms in a regular lattice, outer electrons delocalised ("sea of electrons"). Strong electrostatic attraction between positive metal ions and delocalised electrons. Properties: good conductors of heat and electricity (free electrons), malleable and ductile (layers can slide), high melting points (strong bonds). Alloys: mixture of metals (or metal + non-metal). Different-sized atoms disrupt layers → harder than pure metals.
Simple covalent molecules like water and methane have low boiling points because the intermolecular forces (forces between molecules) are weak. The covalent bonds within the molecules are strong, but they do not break during boiling — only the weak forces between molecules are overcome. Giant covalent substances like diamond have high melting points because many strong covalent bonds must be broken.
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