Ionic compounds don’t come in molecules. Because the attraction between ions pulls in every direction at once, the ions stack into a giant repeating lattice — and that single structural fact explains every physical property examiners ask you about.
📚 What you need to know
An ionic lattice is a giant crystalline structure of alternating cations and anions in a regular, repeating 3D pattern.
The lattice is held by strong electrostatic attractions acting in all directions, and is electrically neutral overall.
Ionic compounds are represented by empirical formulae — the simplest whole-number ratio of ions.
High melting and boiling points, hard but brittle, non-volatile.
Conduct electricity only when molten or dissolved — never as a solid.
Generally soluble in water because of ion–dipole interactions.
Attractions get stronger with greater ionic charge and smaller ionic radius.
What a giant ionic lattice is
A cation doesn’t just attract the one anion it took an electron from. Electrostatic attraction is non-directional — it pulls equally in every direction — so each ion surrounds itself with as many oppositely charged neighbours as will fit. Repeat that billions of times and you get a giant ionic lattice.
Sodium chloride does not exist as NaCl molecules — it is one continuous lattice of alternating ions. NaCl is its empirical formula: the 1:1 ratio, not a molecule count.
The word giant matters: there is no molecule to point to and no edge to the bonding. In sodium chloride, each Na+ is surrounded by six Cl− ions and each Cl− by six Na+ ions, all the way through the crystal.
This is exactly why ionic formulae are empirical. “NaCl” is not one molecule of sodium chloride — it is the statement that the lattice contains Na+ and Cl− in a 1:1 ratio.
What makes some lattices stronger than others
The strength of the electrostatic attraction between two ions depends on two things, and both come from Coulomb’s law:
Strength of ionic attraction
attraction ∝ (charge on ion 1 × charge on ion 2) ÷ (distance between ion centres)
Bigger charges → stronger attraction. Mg2+O2− attracts far more strongly than Na+Cl− — hence melting points of about 2850 °C for MgO versus 800 °C for NaCl.
Smaller ions → stronger attraction, because the charge centres sit closer together. LiF melts higher than LiI for exactly this reason.
Charge is the dominant factor. If two compounds differ in charge, compare charges first; only use size as a tie-breaker.
WORKED EXAMPLE
Explain why magnesium oxide has a much higher melting point than sodium fluoride, given that the ions are of similar size.
Compare the chargesMgO contains Mg²⁺ and O²⁻. NaF contains Na⁺ and F⁻.2 × 2 = 4 vs 1 × 1 = 1Stronger electrostatic attraction between the ionsMore energy needed to break the lattice → higher melting pointSizes are similar, so charge is the deciding factor.
HL extension — lattice enthalpy. The same reasoning is quantified by the lattice enthalpy (ΔHlatt), the enthalpy change when one mole of an ionic solid separates into gaseous ions. It is always endothermic (positive) because bonds are being broken, and it increases with greater ionic charge and smaller ionic radius. For example, NaCl(s) → Na+(g) + Cl−(g), ΔHlatt = +790 kJ mol−1.
Properties of ionic compounds
Every property below is an argument that starts from the same premise: strong electrostatic attractions in all directions, with ions fixed in place.
High melting and boiling points
Melting means giving the ions enough energy to break free of their positions. Because the attractions are strong and act in every direction, that takes a great deal of thermal energy. The same reasoning makes ionic solids non-volatile — they don’t evaporate at room temperature.
Hard but brittle
Ionic crystals resist scratching, yet shatter when struck. A blow shifts one layer of ions relative to the next; suddenly like charges are aligned, the repulsion is enormous, and the crystal splits cleanly along that plane.
Displacing one layer brings like charges into contact. The repulsion that follows is why ionic crystals cleave instead of bending.
Electrical conductivity
Conduction needs mobile charge carriers. Ionic compounds have plenty of charge — the question is whether it can move:
Solid: ions are locked in fixed lattice positions → does not conduct.
Molten: the lattice has broken down and ions can move freely → conducts.
In aqueous solution: ions are separated and hydrated, free to move → conducts.
Say ions, not electrons. Ionic compounds conduct because ions move; metals conduct because delocalised electrons move. Mixing those up costs marks every year.
Solubility in water
Water is polar, so it can attack the lattice from both sides. The δ− oxygen is attracted to cations and the δ+ hydrogens to anions. These ion–dipole interactions pull ions away from the lattice and surround, or hydrate, them.
Water is polar, so it can orient itself around either charge. These ion–dipole attractions release energy and are what make most ionic compounds soluble.
Dissolving is a balance: energy is needed to break the lattice, and energy is released when the ions are hydrated. Where the lattice is exceptionally strong — typically with high charges — the compound may be insoluble, which is why some ionic compounds simply don’t dissolve. Solubility also drops sharply in less polar solvents, which cannot form strong enough interactions with the ions.
Comparing giant ionic with other structures
Paper 1 loves to give you a data table and ask you to identify the structure. This is the comparison to have at your fingertips:
Giant ionic
Giant metallic
Simple molecular
Giant covalent
Particles
Cations and anions
Cations in a sea of delocalised electrons
Small molecules
Atoms
Forces
Electrostatic attraction between ions
Attraction between cations and delocalised electrons
Conducts as a solid? Only metallic (or graphite) does that.
High melting point, conducts only when molten or aqueous?Giant ionic.
Very high melting point but never conducts?Giant covalent.
WORKED EXAMPLE
Deduce the structure of each substance from the data below.
Substance
Melting point / °C
Conducts as solid
Conducts when molten
Solubility in water
W
801
No
Good
Soluble
X
−7
No
No
Almost insoluble
Y
1085
Good
Good
Insoluble
Z
1610
No
No
Insoluble
W — high m.p., conducts only when molten, solublegiant ionicX — very low m.p., never conducts, insolublesimple molecularY — conducts as a solid, so mobile electronsgiant metallicZ — very high m.p. but no conduction at allgiant covalentConducting as a solid is the giveaway that separates Y from W.
💡 Exam tip
“Explain” questions want the chain: strong electrostatic attraction → between oppositely charged ions → a lot of energy needed to overcome → therefore the property.
Never say “ionic bonds are broken when NaCl dissolves or melts” without naming the force — write electrostatic attractions between ions.
For comparisons, lead with charge, then ionic radius.
Remember ionic formulae are empirical — there is no such thing as an NaCl molecule.
⚠️ Common mix-up
Melting a molecular solid breaks intermolecular forces, not covalent bonds. Melting an ionic solid does break the electrostatic attractions holding the lattice — the two cases are not parallel.
Brittle is not weak. Ionic solids are hard; they shatter because of charge repulsion when layers slip, not because the bonding is feeble.
“Ionic compounds conduct electricity” is only half a statement — always add when molten or in aqueous solution.
That completes Ionic Bonding. Next you’ll move into Covalent Bonding, where atoms share electrons instead of transferring them — and where a completely different set of properties follows.
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