In simple terms
A friendly intro before the formal notes — no formulas yet.
Opposite Charges, Locked in a Grid
An ionic bond forms when a metal atom hands one or more electrons to a non-metal atom. This turns them into oppositely charged ions, and the electrostatic pull between those ions — extended in every direction — locks millions of them into a rigid, repeating grid called a giant ionic lattice.
Think of a stadium crowd where every red seat must sit next to blue seats and vice versa. No single red person is 'bonded' to one particular blue person — each is pulled equally by all its blue neighbours in every direction. That is why the whole block is so rigid: to move anyone you would have to fight the pull of all their neighbours at once. That non-directional, all-at-once attraction is exactly what makes ionic solids hard, high-melting and stiff.
- 1
Spot the metal and the non-metal. The metal (low ionisation energy) loses electrons; the non-metal (high electron affinity) gains them.
- 2
Transfer electrons until each ion has a full outer shell — a noble-gas configuration — and write the charge on each ion.
- 3
Balance the total positive and negative charge to deduce the formula (the compound is always neutral overall).
- 4
Explain every property from the lattice: strong attraction in all directions means high melting points, fixed ions mean no conduction until melted or dissolved, and shifting layers means brittleness.
Explore the concept
Use the live diagram, PhET or GeoGebra sim, and synced steps — play it, drag controls, or tap a step.
Step-synced diagram — highlights what to look for in the simulation above.
Step 1
Spot the metal and the non-metal. The metal (low ionisation energy) loses electrons; the non-metal (high electron affinity) gains them.
2 more simulations for this topic — run them in the Simulations section below
Simulations
Every simulation here runs the real model — try the steps on a card, then check what you see against the notes.
2 simulations
- PhETCoreIB S2.1
Coulomb’s Law
Set two charges on the atomic scale and drag them closer or further apart; the force between them updates live.
Try this
- On Atomic Scale set +1e and −1e (like Na⁺ and Cl⁻) and note the force.
- Change them to +2e and −2e (like Mg²⁺ and O²⁻) at the same separation.
- Now move the ions closer together, as for smaller ionic radii.
Look for Doubling both charges quadruples the attraction and a shorter distance raises it further: why MgO has a much larger lattice enthalpy and melting point than NaCl.
Simulation by PhET Interactive Simulations, University of Colorado Boulder · Licensed to MarkScheme (public licence CC BY-NC 4.0 since 2026-03-30)
- PhETJava · best on a laptopIB S2.1
Sugar and Salt Solutions
Shake salt or sugar into water, test the solution with a conductivity probe, then zoom in to the particles.
Try this
- On Macro put the conductivity tester in pure water, then shake in salt.
- Reset, then repeat with sugar.
- Open Micro and compare what sodium chloride and sucrose break up into.
Look for Salt solution lights the bulb because the lattice separates into mobile Na⁺ and Cl⁻ ions; sugar dissolves as neutral molecules and does not conduct. The solid lattice itself has no mobile ions.
Simulation by PhET Interactive Simulations, University of Colorado Boulder · Licensed to MarkScheme (public licence CC BY-NC 4.0 since 2026-03-30)
Full topic notes
Formal explanation with the rigour you need for the exam.
Forming ions by electron transfer
Ionic bonding occurs between elements with a large difference in electronegativity, almost always a metal and a non-metal. The metal has a low ionisation energy, so it loses its valence electrons readily; the non-metal has a high electron affinity, so it accepts them readily. Electrons transfer from metal to non-metal until each atom reaches a stable, full outer shell — the electron configuration of the nearest noble gas. Sodium, [Ne]3s¹, loses its single 3s electron to become Na⁺ (2,8), isoelectronic with neon; chlorine, [Ne]3s²3p⁵, gains one electron to become Cl⁻ (2,8,8), isoelectronic with argon.
Cations are positive ions formed when metal atoms LOSE valence electrons: Na → Na⁺ + e⁻; Mg → Mg²⁺ + 2e⁻.
Anions are negative ions formed when non-metal atoms GAIN electrons: Cl + e⁻ → Cl⁻; O + 2e⁻ → O²⁻.
The magnitude of the charge equals the number of electrons transferred to reach a noble-gas configuration.
An ion and its nearest noble gas are isoelectronic — same number of electrons — but the ion keeps its own nuclear charge, so Na⁺ is smaller than Ne and Cl⁻ is larger than a Cl atom.
Predicting ionic charge and formulae
Because the charge on a monatomic ion is set by how many electrons it must lose or gain to fill its outer shell, you can read the likely charge straight off the periodic table. Group 1 forms 1+, Group 2 forms 2+, Group 13 forms 3+; Group 15 forms 3−, Group 16 forms 2−, Group 17 forms 1−. To build a formula, remember that the compound must be electrically neutral: the total positive charge from the cations must exactly cancel the total negative charge from the anions. The quickest route is to find the lowest whole-number ratio of ions that makes the charges balance.
Write each ion with its charge, e.g. Ca²⁺ and Cl⁻.
Find the lowest ratio that makes total + charge equal total − charge. For Ca²⁺ and Cl⁻ you need two Cl⁻ per Ca²⁺, giving CaCl₂.
A useful shortcut is the 'cross-over' rule: the magnitude of each ion's charge becomes the subscript of the other — but always cancel down to the simplest ratio afterwards (Mg²⁺ with O²⁻ gives Mg₂O₂ → MgO, not Mg₂O₂).
Polyatomic ions (e.g. sulfate SO₄²⁻, nitrate NO₃⁻) balance the same way; keep the whole ion together and use brackets when more than one is needed, e.g. Ca(NO₃)₂.
The giant ionic lattice
Once formed, ions do not simply pair off. The electrostatic attraction of an ion is non-directional — a positive ion attracts every nearby negative ion equally in all directions, and vice versa. Ions therefore pack together in the arrangement that maximises attraction between opposite charges and minimises repulsion between like charges, producing a regular, repeating three-dimensional array called a giant ionic lattice. In sodium chloride each Na⁺ is surrounded by six Cl⁻ and each Cl⁻ by six Na⁺ (a coordination number of 6), and this pattern repeats over the entire crystal. There are no molecules: a formula such as NaCl gives only the simplest ratio of ions, called a formula unit.
Explaining the physical properties
Every characteristic property of an ionic compound follows from one idea: a giant lattice held by strong, non-directional electrostatic attraction, in which the ions are fixed in the solid but can move once the lattice is broken.
High melting and boiling points. Melting requires overcoming the strong electrostatic attraction between oppositely charged ions throughout the lattice, which needs a large amount of thermal energy — so these compounds are solids at room temperature.
Electrical conductivity. As a solid, the ions are locked in fixed lattice positions and cannot move, so it does not conduct. When molten or dissolved in water the ions become mobile and are free to move to the electrodes, so it conducts. The charge carriers are IONS, not electrons.
Brittleness. The lattice is hard, but a blow that shifts one layer of ions brings like charges opposite each other; the resulting repulsion pushes the layers apart and the crystal cleaves — brittle rather than malleable.
Solubility in water. Polar water molecules are attracted to the ions and surround them, pulling them out of the lattice (hydration). When the energy released on hydrating the ions is comparable to the lattice enthalpy, the compound dissolves and the ions become mobile in solution.
Use the exact vocabulary the mark scheme looks for. For melting point say 'strong electrostatic attraction between oppositely charged ions' — not just 'strong bonds'. For conductivity always name the STATE (molten or aqueous) AND the carriers ('ions are free to move'); never write 'free electrons' for an ionic compound, as that describes a metal and is penalised.
Melting point, ionic charge and radius
Not all ionic solids melt at the same temperature. The strength of the electrostatic attraction between two ions — and hence the energy needed to break the lattice — increases when the ionic charges are larger and when the ionic radii are smaller, because both bring more charge closer together. These are precisely the two factors that make a lattice enthalpy larger in magnitude. So magnesium oxide (Mg²⁺, O²⁻ — both doubly charged and small) has a much stronger lattice and a far higher melting point than sodium chloride (Na⁺, Cl⁻ — singly charged and larger). Down a group, as ionic radius grows, melting points of comparable compounds tend to fall (LiF > NaCl > KBr).
Common mistakes examiners penalise
Saying ionic melts/solutions conduct 'because of free electrons' — the charge carriers in an ionic compound are MOBILE IONS. 'Free/delocalised electrons' describes metals and loses the mark.
Explaining conduction without naming the state — you must say the ions conduct when MOLTEN or in AQUEOUS solution, and that in the solid the ions are FIXED. 'Ionic compounds conduct electricity' with no condition is wrong.
Calling a formula unit a 'molecule' — NaCl is the simplest ratio of ions in a giant lattice, not a discrete molecule. Do not write 'NaCl molecule'.
Forgetting to cancel a formula to its lowest ratio — Mg²⁺ with O²⁻ is MgO, not Mg₂O₂; cross-over subscripts must always be simplified.
Getting the charge sign or size wrong — metals form CATIONS (positive) by LOSING electrons; non-metals form ANIONS (negative) by GAINING electrons. The charge equals the number of electrons transferred.
Vague 'strong bonds' for high melting point — name them: 'strong electrostatic attraction between oppositely charged ions' throughout the giant lattice.
Attributing brittleness to breaking bonds directly — the crystal cleaves because shifting a layer aligns LIKE charges, which then REPEL; say the repulsion explicitly.
Model answer — marked the way our engine marks it
Structured 'explain' questions in Paper 2 are marked analytically: the marks are a checklist of distinct valid points, not a holistic markband. Each separate correct idea scores 1 mark, and the engine credits any wording that carries the same chemistry ('equivalent'). A mark tagged M is a reasoning point (the correct link, e.g. 'because the ions are fixed'); a mark tagged A is a correct conclusion or statement of fact. Error-carried-forward (ECF) means a correct consequence drawn from an earlier wrong statement can still score. Study how each of the four marks below is tied to one specific idea.
Where this leads
The ionic model is one of three bonding models in this unit. Next you will meet the covalent model (shared electrons and discrete molecules) and the metallic model (a lattice of cations in a sea of delocalised electrons), and then learn to place real bonds on a continuum of bonding using electronegativity difference. The same habit carries through: describe the structure, identify the forces, and let the forces explain the properties. Master that for ions here and the covalent and metallic models become variations on a method you already own.
Worked examples
See the formulas applied — reveal one step at a time, like the exam.
Deduce the formula of the ionic compound formed between magnesium and nitrogen, showing how the charges balance. [3]
- 1
Step 1 — deduce the ion each element forms. Magnesium is in Group 2, configuration [Ne]3s². It loses 2 electrons to reach a noble-gas configuration: Mg → Mg²⁺ + 2e⁻. Nitrogen is in Group 15, configuration [He]2s²2p³. It gains 3 electrons to fill its outer shell: N + 3e⁻ → N³⁻. [M1: correct ions Mg²⁺ and N³⁻]
Explain, in terms of structure and bonding, why magnesium oxide (m.p. 2852 °C) has a much higher melting point than sodium chloride (m.p. 801 °C). [3]
- 1
Step 1 — identify the ions and charges. Sodium chloride is a giant lattice of Na⁺ and Cl⁻ ions (charges 1+ and 1−). Magnesium oxide is a giant lattice of Mg²⁺ and O²⁻ ions (charges 2+ and 2−). [context]
Explain why sodium chloride has a high melting point and only conducts electricity when molten or dissolved in water. [4]
- 1
Model answer — full response.
How it all connects
The big idea sits in the middle — tap a linked idea to explore the link.
Tap a linked idea to see how it connects back to the main topic — that connection is what examiners reward.
Glossary
Key terms for this topic — skim now; the Check step will test them.
- Ionic bond
The electrostatic force of attraction between oppositely charged ions, formed by the transfer of one or more electrons from a metal atom to a non-metal atom. The attraction is non-directional.
- Cation
A positively charged ion, formed when an atom (usually a metal) LOSES one or more valence electrons. It has more protons than electrons. Example: Na → Na⁺ + e⁻.
- Anion
A negatively charged ion, formed when an atom (usually a non-metal) GAINS one or more electrons. It has more electrons than protons. Example: Cl + e⁻ → Cl⁻.
- Noble-gas configuration
The stable full-outer-shell arrangement ions reach. Na⁺ has the configuration of neon (2,8); Cl⁻ has the configuration of argon (2,8,8). Ions and their nearest noble gas are isoelectronic.
- Predicting ionic charge
From the group: Group 1 → 1+, Group 2 → 2+, Group 13 → 3+; Group 15 → 3−, Group 16 → 2−, Group 17 → 1−. The charge is the number of electrons lost or gained to reach a full shell.
- Giant ionic lattice
A regular, repeating three-dimensional array of alternating cations and anions held by strong electrostatic attraction in all directions. 'Giant' means the lattice extends over the whole crystal — there are no separate molecules.
- Coordination number
The number of nearest oppositely charged ions immediately surrounding a given ion. In sodium chloride each Na⁺ is surrounded by 6 Cl⁻ and each Cl⁻ by 6 Na⁺ (6:6 coordination).
- Solubility in water
Polar water molecules surround and pull the ions from the lattice (hydration). If the energy released by hydrating the ions is comparable to the lattice enthalpy, the compound dissolves; the ions become mobile in solution.
- Factors controlling melting point (lattice strength)
Attraction — and therefore melting point — rises with (1) larger ionic charges and (2) smaller ionic radii, because both increase the electrostatic force between ions. MgO (2+/2−, small ions) melts far higher than NaCl (1+/1−).
- Lattice enthalpy (idea)
The energy needed to separate one mole of a solid ionic lattice into gaseous ions. Larger charge and smaller radius give a more exothermic (larger-magnitude) lattice enthalpy and a stronger, higher-melting solid.
Name it
Read the meaning, then pick which of this lesson’s terms it describes. Miss one and you see what your choice really means.
Attraction — and therefore melting point — rises with (1) larger ionic charges and (2) smaller ionic radii, because both increase the electrostatic force between ions. MgO (2+/2−, small ions) melts far higher than NaCl (1+/1−).
Quick check
Write your answer first, then compare it with the model one — the gap is what you would have lost.
Teach it back
If you can explain it simply, you own it — gaps here are marks you’d lose.
Teach it back
Explain this topic as if teaching a friend. We name the gaps an examiner would still dock.
Revision flashcards
Guess first, then flip — retrieval beats re-reading.
Key takeaways
Review these before you close the topic — retrieval beats re-reading.
Cations are positive ions formed when metal atoms LOSE valence electrons: Na → Na⁺ + e⁻; Mg → Mg²⁺ + 2e⁻.
Anions are negative ions formed when non-metal atoms GAIN electrons: Cl + e⁻ → Cl⁻; O + 2e⁻ → O²⁻.
The magnitude of the charge equals the number of electrons transferred to reach a noble-gas configuration.
An ion and its nearest noble gas are isoelectronic — same number of electrons — but the ion keeps its own nuclear charge, so Na⁺ is smaller than Ne and Cl⁻ is larger than a Cl atom.
Practice — then mark it
The whole point: a real Cambridge question, marked mark-by-mark.
Get a Paper 2 structured answer marked: explain an ionic compound's properties from its structure
Get a Paper 2 structured answer marked: explain an ionic compound's properties from its structure
Extra simulations & links
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Frequently asked
Checkpoint
One marked question is worth ten re-reads — close the loop before you move on.
Reading it isn’t knowing it — prove it.
Before you move on: do Get a Paper 2 structured answer marked: explain an ionic compound's properties from its structure on paper, snap a photo, and get examiner-style feedback on exactly where you win and lose marks.
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