Overview
What you’ll learn
- Describe ionic bonding as the transfer of electrons from metal to non-metal, forming a giant ionic lattice held by electrostatic attraction.
- Deduce the formulae of ionic compounds from the charges on their ions.
- Describe covalent bonding as electron sharing, and draw simple molecules (describing dot-and-cross diagrams in words).
- Compare simple molecular and giant molecular structures, including the allotropes of carbon.
- Describe metallic bonding as a lattice of positive ions in a sea of delocalised electrons.
- Relate each structure type to melting point, boiling point and electrical conductivity.
In 30 seconds
- 1Ionic bonds form when a metal gives electrons to a non-metal: oppositely charged ions then attract in all directions, building a giant ionic lattice.
- 2Ionic formulae come from charge balance: Al³⁺ and O²⁻ give Al₂O₃ because 2(3+) = 3(2−) = 6 charge units cancelled.
- 3Covalent bonds are shared pairs of electrons between non-metals; the share gives each atom a full outer shell.
- 4Simple molecular substances (water, carbon dioxide) have strong bonds inside molecules but weak forces between molecules — hence low melting and boiling points.
- 5Giant molecular substances (diamond, silicon dioxide) bond through a continuous network: very high melting points and no mobile particles in any state.
- 6Diamond: each carbon bonds to 4 others, tetrahedrally, very hard. Graphite: each carbon bonds to 3, in layers that slide, conducting along the surface.
- 7Metallic bonding: positive ions fixed in a lattice, surrounded by delocalised electrons that carry both charge and heat — and let layers slide without shattering.
- 8Conductivity verdicts: ionic solids no (ions fixed), ionic melts/solutions yes (mobile ions), molecular no (no charged particles), graphite yes (one spare electron per atom), metals yes (electron sea).
Notes · 1 of 5
Ionic bonding: transfer and lattice
Metal loses, non-metal gains — then electrostatic attraction locks the ions into a lattice.
When sodium meets chlorine, sodium (2.8.1) hands its single outer electron to chlorine (2.8.7). Both finish with full outer shells: Na⁺ and Cl⁻. But the story continues in three dimensions — each Na⁺ attracts six neighbouring Cl⁻ and vice versa, building a giant ionic lattice in which every attraction is an ionic bond.
Because each ion pulls on all its neighbours equally in every direction, the lattice is held together strongly throughout — that is why ionic compounds are solids at room temperature with high melting points. A dot-and-cross diagram of NaCl shows the electron transfer: the electron sodium started with appears, drawn as a cross, in chlorine’s outer shell after the arrow.
| Ions | Ratio needed | Formula |
|---|---|---|
| Na⁺ and Cl⁻ | 1 : 1 | NaCl |
| Ca²⁺ and Cl⁻ | 1 : 2 | CaCl₂ |
| Mg²⁺ and O²⁻ | 1 : 1 | MgO |
| Al³⁺ and O²⁻ | 2 : 3 | Al₂O₃ |
| Ca²⁺ and NO₃⁻ | 1 : 2 | Ca(NO₃)₂ |
Notes · 2 of 5
Covalent bonding: sharing to fill shells
Non-metals pool electrons; each shared pair is one bond.
Two non-metals both need electrons and neither can hand them over, so they share. A single covalent bond is one shared pair; double and triple bonds share two and three pairs. In words, a dot-and-cross diagram of water shows oxygen sharing one electron from each of two hydrogen atoms — each H holds a pair with O, and oxygen reaches its full 8.
| Molecule | Bonds in it | Why the numbers work |
|---|---|---|
| H₂ | 1 single bond | Each H contributes 1 electron: one shared pair completes each shell (2) |
| H₂O | 2 single bonds | Oxygen (2.6) needs 2 more: shares a pair with each of 2 H |
| NH₃ | 3 single bonds | Nitrogen (2.5) needs 3: one pair with each of 3 H |
| CH₄ | 4 single bonds | Carbon (2.4) needs 4: one pair with each of 4 H |
| O₂ | 1 double bond | Each O needs 2: two shared pairs serve both |
| CO₂ | 2 double bonds | Carbon needs 4: two double bonds, one to each O |
Inside a molecule the shared pair is held tightly by both nuclei — covalent bonds are strong. The weaknesses of a molecular substance live between its molecules, in the feeble intermolecular forces, and that distinction runs every melting-point answer in this topic.
Notes · 3 of 5
Giant molecular structures and the carbon allotropes
When covalent bonding refuses to stop at a molecule: diamond, graphite and silica.
Some covalent substances never finish their bonding: each atom links to neighbours, which link to more neighbours, forming a giant covalent (macromolecular) lattice. The classic examples are diamond, graphite and silicon (IV) oxide (silica, the main ingredient of sand).
| Feature | Diamond | Graphite |
|---|---|---|
| Bonds per carbon atom | 4 (to 4 neighbours) | 3 (to 3 neighbours) |
| Structure | Rigid tetrahedral network | Flat hexagonal layers |
| Melting point | Very high (network must be broken) | Very high (network must be broken) |
| Electrical conductivity | No — all 4 outer electrons are bonding | Yes — the unbonded 4th electron of each atom is delocalised along the layer |
| Soft/hard | Hardest natural substance | Soft and slippery — layers slide over each other |
| Typical uses | Cutting tools | Pencils, lubricant, electrodes |
Silica (SiO₂) copies diamond’s tetrahedral network with Si and O atoms alternating — hence sand’s high melting point. Compare all this with simple molecular substances: even though water’s O–H bonds are strong, the forces between neighbouring water molecules are weak, so ice melts at 0 °C. Melting or boiling a molecular substance never breaks covalent bonds; it only overcomes intermolecular forces.
Notes · 4 of 5
Metallic bonding and the properties of metals
Positive ions in a sea of delocalised electrons explains almost everything metals do.
A metal is a lattice of positive ions (the atoms have surrendered their outer electrons) sitting in a sea of delocalised electrons that belong to the whole crystal. The attraction between the positive ions and the negative electron sea is the metallic bond.
- Conduct electricity — the delocalised electrons drift when a potential difference is applied (electrons in, electrons out).
- Conduct heat — fast-moving electrons pass kinetic energy along the metal quickly.
- Malleable and ductile — layers of ions can slide over one another without breaking the bonding, because the electron sea re-adjusts instantly.
- High melting points (generally) — strong attraction between ions and the dense electron sea; more delocalised electrons per ion (e.g. in magnesium vs sodium) make it stronger.
Notes · 5 of 5
Structure decides properties
The four-lattice summary table that answers most exam questions in one line each.
| Structure | Particles held | Melting point | Conducts when… | Example |
|---|---|---|---|---|
| Giant ionic | Oppositely charged ions | High | Molten or dissolved (mobile ions); never solid | NaCl, MgO |
| Simple molecular | Molecules (strong in, weak between) | Low | Never (no charged particles) | H₂O, CO₂, I₂ |
| Giant molecular | Atoms in a network | Very high | Only graphite (mobile electrons within layers) | Diamond, SiO₂ |
| Giant metallic | Ions in an electron sea | High (varies) | Always solid and liquid (mobile electrons) | Cu, Fe, Al |
Solubility rounds out the picture: ionic compounds dissolve well in water (polar solvent prising ions away) but not in organic solvents; simple molecular substances are the reverse — hexane and ethanol dissolve them far better than water does. When a question asks "explain in terms of structure and bonding", pick the particles that would have to move, and say whether the structure lets them.
Revise
Key terms
The words examiners expect you to use precisely.
- Ionic bond
- The electrostatic attraction between oppositely charged ions, acting in all directions in a giant lattice.
- Covalent bond
- A shared pair of electrons between two atoms.
- Metallic bond
- The attraction between positive ions and a sea of delocalised electrons.
- Giant ionic lattice
- The regular 3-D array of ions in an ionic compound.
- Intermolecular forces
- The weak attractions between separate molecules; overcome on melting and boiling.
- Giant covalent (macromolecular) structure
- A network of atoms bonded covalently throughout, e.g. diamond or silica.
Show 4 more key termsShow fewer key terms
- Allotrope
- One of several forms of the same element in the same state, e.g. diamond and graphite.
- Delocalised electrons
- Electrons not fixed to any one atom, free to move through the structure.
- Malleable
- Able to be hammered into shape without breaking — metals, because layers slide.
- Dot-and-cross diagram
- A bonding drawing where each atom’s outer electrons are shown as dots or crosses.
Revise
Common mistakes
Where students most often lose marks in this chapter — and what to do instead.
Mistake: Saying covalent bonds break when water boils.
Boiling separates water molecules from each other; every O–H bond stays intact.
Instead: State changes overcome intermolecular forces, not covalent bonds.
Mistake: Explaining low molecular melting points as "weak covalent bonds".
The covalent bonds are strong; it is the forces between molecules that are weak.
Instead: Little energy is needed to separate molecules, not to break bonds inside them.
Show 10 more mistakesShow fewer mistakes
Mistake: Saying ionic solids conduct because they contain ions.
In the solid the ions are locked in the lattice — present but not mobile.
Instead: Ionic compounds conduct only when molten or dissolved, when ions are free to move.
Mistake: Writing NaCl₂ or CaCl as an ionic formula.
Formulae must balance charge: Na⁺ needs one Cl⁻; Ca²⁺ needs two Cl⁻.
Instead: Total positive charge = total negative charge in every ionic formula.
Mistake: Saying metals conduct "because their ions move".
The ions stay in the lattice; the electrons move.
Instead: Delocalised electrons carry charge through a metal.
Mistake: Claiming graphite is soft because its C–C bonds are weak.
The C–C bonds are strong; the weak point is between layers.
Instead: Layers slide because the forces between them are weak, while each layer is strong internally.
Mistake: Saying diamond conducts because it has delocalised electrons.
All four outer electrons of each carbon are bonding — nothing is delocalised.
Instead: Diamond has no mobile charged particles; graphite has one delocalised electron per atom.
Mistake: Drawing an ionic bond as one arrow in all explanations.
The transfer arrow only creates the ions; the bond itself is the lattice-wide electrostatic attraction.
Instead: Describe attraction between oppositely charged ions acting in all directions.
Mistake: Giving AlO or Al₃O₂ as the formula of aluminium oxide.
Charges 3+ and 2− do not cancel in a 1:1 or 3:2 ratio.
Instead: 2 Al³⁺ (6+) balance 3 O²⁻ (6−): Al₂O₃.
Mistake: Saying covalent molecules conduct when dissolved in water.
Dissolving a molecular substance makes a solution of molecules — no ions form (sugar water does not conduct).
Instead: Only substances that produce mobile ions conduct; pure molecular solutions do not.
Mistake: Explaining malleability as "weak metallic bonds".
Metallic bonding is strong; the ions can still slide because the electron sea adjusts.
Instead: Layers of positive ions slide while the delocalised electrons hold the structure together.
Mistake: Mixing up "atoms" and "molecules" when describing structures, e.g. calling diamond "molecular".
Diamond is one continuous network — there are no separate molecules in it.
Instead: Name the structure first: giant ionic, simple molecular, giant molecular or metallic.
Check yourself
Quick quiz
20 questions with instant explanations. About 5 minutes.
Question 1 of 3
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Define bonding types
An ionic bond is best described as
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Practise
Topical questions
Exam-style practice in three tiers, from recall to data questions.
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- 3Challenging
12 questions worth 54 marks, each with a model answer, mark scheme and examiner tips.
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Practice paper
A 30-mark paper set out like a school paper, with a separate mark scheme. Give yourself 40 min.
Question paper
Chapter 4 Practice Paper
30 marks · 40 min · 4 pages
5 multiple-choice questions and structured questions with answer lines. Chapter 4 (ionic, covalent and metallic bonding; structures and properties).
Mark scheme
PROChapter 4 Practice Paper — mark scheme
Answers · Marking points · 3 pages
Model answers with a mark for each creditworthy point — for marking at home or in class.
Where this chapter is tested
Each school sets its own scope — check your assessment outline.
Questions
Frequently asked questions
What does Sec 3 Chemistry Chapter 4 cover?
It covers sub-topic 2.3 of O-Level Pure Chemistry 6092: ionic, covalent and metallic bonding; the giant ionic, simple molecular, giant molecular and metallic structures; the allotropes of carbon; and explaining melting point, conductivity and solubility from structure.
How do I deduce an ionic formula quickly?
Write the two ions with charges, then choose the smallest whole-number ratio that makes total positive charge equal total negative charge. Al³⁺ with O²⁻: 3 × 2+ = 6+ and 2 × 2− = 6−, so Al₂O₃. Brackets are needed when several polyatomic ions are used, as in Ca(NO₃)₂.
Why do examiners keep asking "in terms of structure and bonding"?
Because the answer must name the structure, identify the particles that would have to move (ions, molecules, atoms, electrons), and say whether the structure lets them move — with the energy cost compared. Saying "it has strong bonds" without saying which attractions break earns nothing.
Which solids conduct, and why?
Metals (delocalised electrons) and graphite (one delocalised electron per atom along its layers) conduct as solids. Ionic compounds conduct only when molten or dissolved — mobile ions, not electrons. Simple and giant molecular substances never conduct because they have no charged particles at all.
When is this chapter usually tested in school?
Bonding follows atomic structure in Sec 3 Term 2, so it appears in Weighted Assessment 2 and the end-of-year examination — and every later topic (electrolysis, organic chemistry, metals) leans on it. Scope varies by school, so check your school’s assessment outline.
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Chapter 5: Chemical Calculations — The Mole Concept and Stoichiometry