Ionic, covalent and metallic bonding — and exactly how the structure of each explains melting point, conductivity and solubility.
Bonding and structure is one of the most heavily examined topics in GCSE Chemistry, and for good reason — it underpins almost everything else. Reactivity, physical properties, uses of materials — all of these connect back to how atoms are bonded and how those bonds are arranged. This guide covers all three bonding types in full, with detailed explanations of how structure explains every physical property you'll be asked about.
Ionic bonds form between metals and non-metals. The metal atom loses one or more electrons to become a positively charged ion (cation). The non-metal atom gains those electrons to become a negatively charged ion (anion). The two oppositely charged ions attract each other through electrostatic forces — this attraction is the ionic bond.
The key is to think about which atoms lose electrons and which gain them. Metals in Group 1 lose 1 electron (forming 1+ ions). Group 2 metals lose 2 electrons (2+ ions). Non-metals in Group 7 gain 1 electron (1− ions). Group 6 non-metals gain 2 electrons (2− ions). The number of electrons lost or gained reflects the number of electrons needed to achieve a full outer shell.
Ionic compounds don't exist as individual pairs of ions — they form giant ionic lattices. Every positive ion is surrounded by negative ions and every negative ion is surrounded by positive ions, in a regular three-dimensional arrangement. Sodium chloride has a cubic lattice: each Na⁺ is surrounded by 6 Cl⁻ and vice versa.
Covalent bonds form between non-metals. Instead of transferring electrons, the atoms share pairs of electrons. Each shared pair constitutes one covalent bond. Both atoms count the shared electrons as their own, giving both a full outer shell.
Simple covalent molecules include H₂ (one bond), O₂ (double bond), N₂ (triple bond), H₂O (two bonds), CO₂ (two double bonds), NH₃ (three bonds), and CH₄ (four bonds). The bonding within these molecules is strong — but the forces between separate molecules (intermolecular forces) are very weak.
Simple covalent molecules have low melting and boiling points because it is the weak intermolecular forces between molecules that are overcome during melting — not the strong covalent bonds within them. Water boils at 100°C. Hydrogen (H₂) boils at −253°C.
Simple molecular substances do not conduct electricity — they have no free electrons or ions.
Some covalent substances form giant lattices — enormous networks of covalently bonded atoms. These have very different properties from simple molecules.
Diamond and graphite are both pure carbon — same element, completely different properties. The difference is entirely due to structure. Diamond has a 3D tetrahedral lattice (very hard, doesn't conduct). Graphite has flat layers with delocalised electrons (soft, conducts). This comparison is one of the most commonly tested points in GCSE bonding questions.
In metals, the outer electrons from each atom are released into a shared pool — a "sea" of delocalised electrons. The positively charged metal ions sit in this electron sea, attracted to the electrons surrounding them. This attraction between positive ions and delocalised electrons is the metallic bond.
Pure metals have layers of identical-sized atoms that slide easily over each other — they are relatively soft. Alloys contain atoms of different sizes (from different elements). These disrupt the regular arrangement of layers, making it harder for the layers to slide. Steel (iron + carbon) is much harder than pure iron. Brass (copper + zinc) is harder than pure copper. Exam questions often ask you to explain why alloys are harder — the disrupted layer structure answer scores full marks.
A common exam question gives you a set of physical properties and asks you to identify the bonding type. Here's a quick decision guide:
The full bonding and structure specification for AQA is at the AQA GCSE Chemistry specification page. Edexcel's is at the Edexcel GCSE Chemistry page.
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