There is no such thing as a molecule of sodium chloride. A grain of salt is one enormous, repeating three-dimensional grid of ions, and almost every physical property ionic compounds have — high melting point, brittleness, conducting only when molten — falls straight out of that one structural fact.
📘 What you need to know
An ionic lattice is a giant, regular, repeating 3D arrangement of alternating cations and anions.
Each ion is attracted to all the oppositely charged ions around it, in every direction.
The lattice is electrically neutral overall, and the formula is empirical — a ratio, not a molecule.
Lattice enthalpy is the energy needed to separate one mole of a solid lattice into gaseous ions. It is always endothermic and therefore positive.
Lattice enthalpy increases with larger ionic charge and smaller ionic radius.
Ionic solids have high melting points, are hard but brittle, and conduct only when molten or dissolved.
They dissolve in water because ion–dipole attractions to polar water molecules release energy.
Why there is no NaCl molecule
In a covalent substance like water you can point at a molecule: two hydrogens and one oxygen, a discrete unit with a beginning and an end. In an ionic solid you cannot. Each Na+ pulls on every Cl− near it, and each Cl− pulls back on every Na+. There is no way to draw a line around one “molecule” that would not cut through bonds.
So the formula NaCl does not mean “one sodium joined to one chlorine”. It means “sodium and chloride ions in a 1 : 1 ratio”, which is exactly what an empirical formula is.
Only 27 ions are drawn. The real lattice carries on in all six directions until it reaches the surface of the crystal, which is why ionic solids form neat cubic grains.
Coordination number. In sodium chloride each Na+ has six Cl− as immediate neighbours — one above, one below, and four around the sides — and each Cl− has six Na+. That is described as 6 : 6 coordination.
Lattice enthalpy
How strong is an ionic lattice? Put a number on it by asking how much energy it takes to pull the whole thing completely apart into free gaseous ions.
Two things follow from the definition. First, you are breaking attractions, so it always costs energy — lattice dissociation enthalpy is always endothermic and always positive. Second, the size of that number is a direct measure of how strongly the ions are held.
What makes it bigger? The same two things that make any electrostatic attraction stronger:
Larger charges. Doubling both charges roughly quadruples the attraction, which is why 2+/2− compounds sit in a completely different league from 1+/1− ones.
Smaller ions. Smaller radii let the centres of charge sit closer together, and attraction falls off sharply with distance.
Compound
Ions
Lattice enthalpy / kJ mol−1
Melting point / °C
KCl
K+ and Cl−
711
770
NaCl
Na+ and Cl−
790
801
CaO
Ca2+ and O2−
3401
2572
MgO
Mg2+ and O2−
3791
2852
Read that table across, not down. KCl to NaCl is a small change, because only the cation size differs. NaCl to MgO is a jump of nearly five times, because both charges have doubled. Charge beats size, every time — and the melting points track the lattice enthalpies almost perfectly.
The properties, and why
High melting and boiling points
Melting means giving the ions enough energy to break free of their fixed positions. Every ion is held by strong electrostatic attractions acting in all directions at once, so a great deal of energy is needed. Sodium chloride melts at 801 °C; magnesium oxide, with its doubled charges, needs 2852 °C.
Hard, but brittle
These two sound contradictory and are not. Hard means difficult to scratch, because the ions are locked tightly in place. Brittle means it shatters rather than bends, and that is a consequence of the alternating pattern.
Compare this with a metal, where the layers slide over a sea of delocalised electrons and nothing repels. That is why metals bend and ionic solids crack.
Conducting only when molten or dissolved
Electrical conductivity needs charged particles that are free to move. In an ionic solid the ions are charged but locked in the lattice, so nothing moves and nothing conducts. Melt it, or dissolve it in water, and the lattice breaks up — now the ions can move to the electrodes and the substance conducts.
Two conditions, both needed. Charged particles, and freedom to move. Ionic solids fail the second; that is the whole explanation, and it is worth writing out in full in an exam.
Generally soluble in water
Water is a polar molecule, with a partial negative charge on the oxygen and partial positives on the hydrogens. Those partial charges are attracted to the ions in the lattice.
The ions end up hydrated — each one wrapped in its own shell of water molecules. This is what the state symbol (aq) actually means.
Solubility is a balance. Breaking the lattice costs energy; hydrating the ions releases it. When the hydration energy roughly covers the lattice enthalpy, the compound dissolves. When the lattice is exceptionally strong — as in magnesium oxide — hydration cannot pay for it and the compound stays insoluble.
Comparing the structure types
Exam questions rarely ask about ionic lattices in isolation. They give you a table of properties and expect you to identify which structure each substance has.
Property
Giant ionic
Giant metallic
Simple molecular
Giant covalent
Melting point
high
moderately high to high
low
very high
Conducts as a solid
no
yes
no
no, except graphite
Conducts when molten
yes
yes
no
no
Solubility in water
generally soluble
insoluble, some react
usually insoluble unless polar
insoluble
Hardness
hard but brittle
hard and malleable
soft
very hard, except graphite
Particles present
ions
cations in a sea of electrons
small molecules
atoms
Example
NaCl
copper
Br2
silicon(IV) oxide
🧩 Identifying a structure from data
Look at the melting point first. Below about 300 °C means simple molecular; anything in the hundreds or thousands means a giant structure.
Check conduction as a solid. Only metals and graphite do this, so a “yes” here almost always means metallic.
Check conduction when molten. A “no” as a solid but “yes” when molten is the fingerprint of giant ionic.
Use solubility to confirm. Soluble in water supports ionic; insoluble supports metallic or giant covalent.
State the evidence, not just the answer. Marks are awarded for the reasoning far more often than for the label.
Worked examples
WORKED EXAMPLE
Comparing melting points
Magnesium oxide melts at 2852 °C but sodium chloride melts at 801 °C. Explain the difference.
Step 1: compare the chargesMgO has Mg²⁺ and O²⁻. NaCl has Na⁺ and Cl⁻.product of charges: 2 × 2 = 4 versus 1 × 1 = 1Step 2: compare the sizesMg²⁺ and O²⁻ are both smaller than Na⁺ and Cl⁻, so the ions sit closer together.Step 3: link to energyBigger charges and shorter distances both give stronger electrostatic attraction, so the lattice enthalpy is far larger.More energy is needed to separate the ions, so MgO melts much highercharge is the dominant factor here — mention it first
WORKED EXAMPLE
Ordering lattice enthalpies
Place KCl, NaCl and CaO in order of increasing lattice enthalpy, and justify the order.
Step 1: separate by charge firstCaO is 2+ / 2−, the other two are 1+ / 1−, so CaO must be the largest by a long way.Step 2: separate the remaining two by sizeK⁺ is larger than Na⁺, so the ions in KCl are further apart and held less strongly.KCl < NaCl < CaO711, 790 and 3401 kJ mol⁻¹ — the jump to CaO is enormous
WORKED EXAMPLE
Identifying structures from properties
Deduce the type of structure present in each substance.
Substance
Melting point / °C
Conducts as a solid
Conducts when molten
A
801
no
yes
B
−7
no
no
C
1085
yes
yes
Substance AHigh melting point means a giant structure. It cannot conduct as a solid but can when molten, which needs charged particles that become mobile on melting.giant ionicSubstance BMelts below room temperature, so only weak forces are being broken. No conduction at all means no free charged particles.simple molecularSubstance CConducts as a solid, so it has delocalised electrons that are already free to move.giant metallic“conducts as a solid” is the single most decisive line in a table like this
WORKED EXAMPLE
Explaining brittleness
Explain why solid sodium chloride shatters when it is struck, rather than deforming like a metal.
Step 1: describe the arrangementThe lattice has alternating positive and negative ions in fixed positions.Step 2: describe what the blow doesA force shifts one layer of ions by one position relative to the next.Step 3: the consequenceIons of the same charge are now adjacent, so attraction is replaced by repulsion.The layers repel and the crystal splits along that planethe word “repulsion” is almost always the mark-scoring word here
💡 Exam tip
Say “electrostatic attraction between oppositely charged ions” in full. “Strong bonds” earns nothing.
For melting point comparisons, always cite both charge and ionic radius, and say which matters more.
For conductivity, always give both conditions: charged particles present and free to move.
Use the word lattice, not “molecule”, anywhere near an ionic compound. It signals you understand the structure.
Lattice enthalpy values are in the data booklet, so you can quote them in comparison questions rather than guessing.
In identify-the-structure tables, quote the specific data you used. “High melting point of 801 °C suggests a giant structure” beats “it is ionic”.
⚠ Common mix-up
Talking about molecules of NaCl. There are none. The formula is an empirical ratio.
Saying ionic solids conduct because the ions are charged. They are charged in the solid too, and it does not conduct. Mobility is the missing half.
Confusing hard with strong, or brittle with weak. Ionic solids are hard and brittle at the same time, for the same structural reason.
Saying lattice enthalpy is exothermic. Lattice dissociation is endothermic and positive; only lattice formation is negative.
Ranking by ionic radius before charge. Charge dominates, so compare charges first and only then use size.
Assuming all ionic compounds dissolve. When the lattice enthalpy is very large, as in MgO, hydration cannot compete.
Up next: The Covalent Model — what happens when neither atom is willing to hand an electron over, and sharing becomes the only option.
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