Ionic, Covalent and Metallic Bonding: Revision Notes

Condensed for the final weeks. For the full explanation, work through the Ionic, Covalent and Metallic Bonding study guide first, then test yourself with the practice questions.

The four structures — one table to learn

	Ionic	Simple molecular	Giant covalent	Metallic
Particles	Ions	Molecules	Atoms	Ions + delocalised electrons
Held by	Electrostatic attraction	Weak forces between molecules	Covalent bonds throughout	Attraction of ions to electron sea
Melting point	High	Low	Very high	Usually high
Conducts solid	No	No	No (except graphite)	Yes
Conducts molten/aqueous	Yes	No	No	Yes
Solubility in water	Usually soluble	Usually insoluble	Insoluble	Insoluble

The ionic lattice, precisely

An ionic solid is a giant lattice — a regular, repeating 3D arrangement of alternating positive and negative ions, each ion surrounded by ions of the opposite charge. “Regular arrangement of alternating ions” is the phrase examiners look for; “ions stuck together” earns nothing. This lattice structure is exactly why melting an ionic compound takes so much energy — every ion is held by attractions in all directions at once, not just to one neighbour.

The rule that answers most questions

Properties follow from what must be broken.

• Melting a simple molecular substance overcomes only the weak forces between molecules — the covalent bonds inside stay intact. Hence low melting points.
• Melting ionic or giant covalent means breaking strong bonds throughout the lattice. Hence high melting points.
• Conduction needs charged particles free to move — ions freed by melting or dissolving, or delocalised electrons.

Bond formation

• Ionic — metal loses electrons (forms cation), non-metal gains (forms anion). Both reach a full outer shell.
• Covalent — two non-metals share pairs of electrons.
• Metallic — metal atoms lose outer electrons into a delocalised sea.

Dot-and-cross diagrams

Show only the outer shell, with dots for one atom’s electrons and crosses for the other’s, so the examiner can see where each electron came from.

• Ionic (electron transfer) — e.g. sodium chloride: Na’s one outer dot moves across to Cl, giving [Na]⁺ (empty outer shell) and [Cl]⁻ with 8 outer electrons. Magnesium oxide (0620 Extended/5070; Core covers only Group I with Group VII) is the same idea with two electrons transferred: [Mg]²⁺ and [O]²⁻, both ions carrying a 2+/2− charge.
• Covalent (electron sharing) — e.g. water, methane and hydrogen chloride: each shared pair sits between the two atoms, and every atom ends with a full outer shell (2 for hydrogen, 8 for everything else). Carbon dioxide (Extended/5070) needs double bonds — two shared pairs between carbon and each oxygen.

Describe vs explain — the tier trap

O Level 5070 and IGCSE 0620 Extended both require you to explain properties in terms of structure and bonding (why high melting point? — because breaking the lattice needs energy). IGCSE 0620 Core only requires you to describe the properties themselves, without the mechanism. If a question says “explain”, a bare description of the property with no reference to structure or bonding earns no marks, even if the fact stated is correct.

Carbon and silicon: the giant covalent family

	Diamond	Graphite
Bonds per carbon	4	3
Structure	Tetrahedral 3D network	Layers of hexagons
Conducts?	No	Yes — one delocalised electron per atom
Hardness	Hardest natural substance	Soft, slippery
Why	All 4 electrons in bonds	Weak forces between layers let them slide

Silicon(IV) oxide, SiO₂, is a third giant covalent structure, O Level 5070 required content and IGCSE 0620 Extended. It resembles diamond structurally — each silicon atom forms four covalent bonds and each oxygen forms two, and all the outer electrons are localised, either in a bonding pair or a lone pair — so, like diamond, it has a very high melting point and does not conduct electricity, because there are no delocalised electrons free to carry charge.

Exam traps

• Never say covalent bonds break when a simple molecular substance melts — only the intermolecular forces are overcome.
• Ionic solids do not conduct — ions are fixed in the lattice. Only molten or aqueous.
• Graphite conducts because of delocalised electrons, not because it is a metal.
• Say “electrostatic attraction between oppositely charged ions”, not just “the ions attract”.
• Metals are malleable because layers of ions slide while the electron sea maintains the bonding.
• Writing “electrons are shared” for an ionic bond, or “electrons are transferred” for a covalent bond — the two mechanisms are opposites and examiners specifically test the swap.
• If the question says “explain”, a description with no reference to structure or bonding scores zero, however correct the fact.

Self-test

1. Why does sodium chloride conduct when molten but not when solid?
2. Why does carbon dioxide have a much lower boiling point than silicon dioxide?
3. Explain why graphite is used as a lubricant and as an electrode.
4. Why are metals malleable when ionic solids are brittle?
5. Which particles are lost, gained or shared in each of the three bond types?
6. Describe the structure of an ionic lattice, using the phrase examiners expect.
7. SiO₂ and diamond are both giant covalent and very hard, but only one of the two elements involved (carbon) also forms a soft, slippery, conducting allotrope. Explain why graphite conducts but SiO₂ does not.

Answers: 1. Ions are fixed in the lattice when solid; melting frees them to move and carry charge. 2. CO₂ is simple molecular — only weak intermolecular forces need overcoming; SiO₂ is giant covalent, requiring strong covalent bonds to be broken. 3. Weak forces between layers let them slide (lubricant); one delocalised electron per carbon carries charge (electrode). 4. Metal layers slide while delocalised electrons maintain attraction; in ionic solids displacement brings like charges together, which repel and split the crystal. 5. Ionic — transferred; covalent — shared; metallic — delocalised into a sea. 6. A regular, repeating 3D arrangement of alternating positive and negative ions. 7. In graphite each carbon uses only 3 of its 4 outer electrons in bonds, leaving one delocalised per atom; in SiO₂, all the outer electrons are localised in bonding pairs or lone pairs, so none is delocalised or free to carry charge.