What an ionic solid really is
When sodium reacts with chlorine, each sodium atom loses one electron and each chlorine atom gains one. You get Na⁺ ions and Cl⁻ ions. But the solid you can hold in your hand is not a box of separate “NaCl pairs”. It is a single, continuous structure.
- An ionic solid is a repeating three-dimensional array of cations and anions. The pattern repeats in every direction until it reaches the surface of the crystal.
- Every cation is surrounded by anions, and every anion is surrounded by cations.
- This arrangement makes the attractions between opposite charges as large as possible and keeps ions of the same charge as far apart as possible.
- There are no molecules. The formula NaCl tells you the ratio of ions (one Na⁺ for every Cl⁻). One Na⁺ and one Cl⁻ together are called a formula unit, but a formula unit is a counting unit, not a particle that exists on its own.
Why does this arrangement form? Think of it as an energy argument. Every time a cation sits next to an anion, the potential energy drops (Topic 2.2). Every time two like charges sit close together, the energy rises. The lowest-energy arrangement is the one where each ion has as many opposite-charge neighbours as possible and no like-charge neighbours touching it. Repeating that local pattern in three dimensions gives a crystal.
Scope note. The course does not assess specific crystal structures. You will not be asked to name a structure, count neighbours in a unit cell or recognise the “rock-salt” arrangement. The diagrams on this page are illustrations of the general idea.
Drawing a particulate model of an ionic solid
Exam questions often ask you to draw or choose a particulate diagram. A model that is “consistent with Coulomb’s law and the properties of the ions” follows four rules.
- Alternate the charges. Each ion should touch only ions of the opposite charge. Two cations or two anions should never be drawn touching.
- Show sensible relative sizes. A cation is smaller than its parent atom (it has lost its outer shell or some electrons), and an anion is larger than its parent atom. In most common ionic compounds the cation is drawn smaller than the anion. If radii are given, follow them: in KF, for example, K⁺ (138 pm) is slightly larger than F⁻ (133 pm).
- Keep the ratio from the formula. In CaF₂ there must be twice as many F⁻ ions as Ca²⁺ ions in the region you draw. Check that the total charge in your drawing is zero (or close to zero for a small, cut-off section).
- Label the charges. Use + and − signs (or the ion symbols) on every ion or in a key. Do not rely on colour alone.
Figure 1 shows a two-dimensional slice. A real crystal has layers above and below, and in those layers a Na⁺ ion sits above each Cl⁻ ion and vice versa, so opposite charges also face each other between layers. For the exam a clear 2-D drawing is usually enough, as long as it obeys the four rules.
Ion sizes you can use
Ionic radii are measured from the distances between ion centres in crystals. Different data tables split the distance between cation and anion in slightly different ways, so you may see different numbers in other books. The trends are the same in every table. This guide uses these effective radii (in picometres, pm; 1 pm = 10⁻¹² m):
| Cations | Radius (pm) | Anions | Radius (pm) |
|---|---|---|---|
| Li⁺ | 76 | F⁻ | 133 |
| Na⁺ | 102 | Cl⁻ | 181 |
| K⁺ | 138 | Br⁻ | 196 |
| Mg²⁺ | 72 | I⁻ | 220 |
| Ca²⁺ | 100 | O²⁻ | 140 |
Two patterns from Topic 1.7 explain the table. Going down a group, ions get larger because they have more electron shells (Li⁺ < Na⁺ < K⁺; F⁻ < Cl⁻ < Br⁻ < I⁻). Across a period, ions with the same number of electrons get smaller as the number of protons rises: Na⁺ (11 protons) is larger than Mg²⁺ (12 protons), though both have 10 electrons.
Coulomb’s law sets the strength of the array
In Topic 2.2 you used Coulomb’s law for a single cation–anion pair. The same idea applies to the whole crystal, because every attraction in the array is a cation–anion interaction.
The size of the attraction between two ions is proportional to the product of their charges and inversely related to the distance between their centres:
- Potential energy of attraction ∝ (q₊ × q₋) / d, where d is the distance between the centres of a cation and a neighbouring anion.
- For touching ions, d ≈ r₊ + r₋ (the sum of the two radii).
So you can rank ionic solids using a simple relative measure, |q₊ × q₋| / d. It is not an energy in joules, but it puts compounds in the right order when you compare similar solids. In the table, “relative attraction” is each compound’s |q₊ × q₋| / d divided by the value for NaCl.
| Compound | Charges | d = r₊ + r₋ (pm) | Relative attraction (NaCl = 1.00) | Melting point (°C) |
|---|---|---|---|---|
| KCl | +1, −1 | 138 + 181 = 319 | 0.89 | 770 |
| NaCl | +1, −1 | 102 + 181 = 283 | 1.00 | 801 |
| NaF | +1, −1 | 102 + 133 = 235 | 1.20 | 993 |
| MgO | +2, −2 | 72 + 140 = 212 | 5.34 | 2852 |
Two rules come out of the table:
- Charge has the biggest effect. Doubling both charges multiplies the product q₊q₋ by 4. That is why MgO melts more than 1,800 °C higher than NaF, even though the ions are only a little closer together.
- For equal charges, smaller ions win. NaF > NaCl > KCl, because the ion centres get further apart from left to right in that list.
The relative measure does not predict melting points exactly (MgO’s value is about 5 times NaCl’s, but its melting point in °C is not 5 times larger). Melting points also depend on other factors, such as the exact arrangement of the ions in each crystal. Use the measure for ranking, not for calculating a melting point.
From the model to the properties
The particulate model explains why ionic solids behave as they do. You will study these properties fully in Topic 3.2, but it helps to see the link now.
- High melting points. Melting means disrupting many strong cation–anion attractions throughout the array, which needs a lot of energy.
- Brittleness. If a force pushes one layer of ions along by one ion, like charges end up facing each other. The repulsion splits the crystal along that plane (Figure 2).
- Electrical conductivity only when the ions can move. In the solid, ions vibrate about fixed positions and cannot carry a current. When the solid is melted or dissolved in water, the ions are free to move, so the liquid or solution conducts.
Worked example 1: evaluating particulate models of potassium chloride
Question. Three students draw a section of solid potassium chloride, KCl. Radii: K⁺ 138 pm, Cl⁻ 181 pm.
- Student A draws separate K–Cl pairs, each pair a little apart from the next pair.
- Student B draws a grid of alternating + and − circles, but makes every K⁺ circle larger than every Cl⁻ circle.
- Student C draws a grid of alternating + and − circles, with the Cl⁻ circles larger, and 8 K⁺ and 8 Cl⁻ ions in total.
Which model is best, and what is wrong with the others?
- Check the structure. Student A’s drawing shows molecules. In an ionic solid every ion is attracted to all of its neighbours, not to one partner, so there are no pairs. A is wrong.
- Check relative sizes. K⁺ (138 pm) is smaller than Cl⁻ (181 pm); the ratio of radii is 138 ÷ 181 = 0.76. Student B reverses the sizes. B shows the arrangement correctly but not the ion properties.
- Check the ratio and the charge. KCl is 1 : 1. Student C has 8 K⁺ and 8 Cl⁻, so the total charge is 8(+1) + 8(−1) = 0. The charges alternate and the sizes are sensible.
Answer. Student C’s model is best. A shows molecules instead of an extended array; B draws the cation larger than the anion.
Why it matters. Exam diagrams are marked on exactly these features: an extended array, alternating charges, sensible relative sizes and the formula ratio.
Worked example 2: ranking the strength of attraction with Coulomb’s law
Question. Use the radii in the table above to rank NaF, KCl and MgO in order of increasing strength of attraction between their ions. Explain the ranking, and say what you would predict about their melting points.
- Distances. d(KCl) = 138 + 181 = 319 pm; d(NaF) = 102 + 133 = 235 pm; d(MgO) = 72 + 140 = 212 pm.
- Charge products. KCl and NaF: |(+1)(−1)| = 1. MgO: |(+2)(−2)| = 4.
- Relative measure |q₊q₋| / d.
- KCl: 1 ÷ 319 = 0.00313 pm⁻¹
- NaF: 1 ÷ 235 = 0.00426 pm⁻¹
- MgO: 4 ÷ 212 = 0.01887 pm⁻¹
- Compare. NaF is 0.00426 ÷ 0.00313 = 1.36 times KCl, because its ions are smaller. MgO is 0.01887 ÷ 0.00426 = 4.43 times NaF. The charge product alone gives a factor of 4; the slightly shorter distance (235 ÷ 212 = 1.11) adds the rest.
Answer. KCl < NaF < MgO. KCl and NaF have the same charges, but Na⁺ and F⁻ are smaller than K⁺ and Cl⁻, so the ion centres are closer and the attraction is stronger. MgO has ions with double the charge, so the product of charges is four times larger, and its ions are also small. Predicted melting points follow the same order, which matches the data: 770 °C, 993 °C and 2852 °C.
Check. A charge difference beats a size difference: never rank a 2+/2− compound below a 1+/1− compound unless the 2+/2− ions are far larger.
Common misconceptions
- “Solid NaCl is made of NaCl molecules.” No. It is an extended array of ions. The formula is a ratio.
- “Each Na⁺ is bonded only to the Cl⁻ that took its electron.” No. Once the ions form, each Na⁺ is attracted to every nearby Cl⁻. The electron transfer does not create partners.
- “Bigger ions give stronger attraction because they have more electrons.” No. Bigger ions keep the centres further apart, which weakens the attraction. What matters is charge and distance.
- “Only charge matters” (or “only size matters”). Both appear in Coulomb’s law. Charge usually dominates, but for ions with equal charges, size decides.
- “Ionic solids conduct electricity because they contain charged particles.” Charged particles must be free to move to carry a current. In the solid they are locked in place.
- Drawing cations and anions the same size, or touching like charges. Both make a diagram inconsistent with the ions’ properties and with Coulomb’s law.
- “The formula tells you the crystal structure.” The formula tells you only the ratio of ions. Structures are not assessed in this course.
Where this leads
The ionic array is one of four kinds of solid you will compare. Next, Topic 2.4, Structure of Metals and Alloys, builds a different model, positive ions in a sea of electrons, and explains why metals bend where ionic crystals shatter. In Unit 3 you will use these models to explain the properties of solids (Topic 3.2) and why ionic compounds dissolve in water (Topic 3.10). Try the practice questions now, then use the revision notes and the checklist to consolidate.
Quick questions
Is NaCl a molecule?
No. Solid sodium chloride is a continuous array of Na⁺ and Cl⁻ ions. The formula NaCl gives the ratio of ions (1 : 1), called a formula unit, not a separate particle.
Do I need to learn crystal structures such as the rock-salt or cesium chloride structure?
No. The course states that specific crystal structures are not assessed. You need a model that alternates charges, uses sensible sizes and keeps the correct ratio.
Why are cations usually drawn smaller than anions?
A cation has lost its outer electrons, so it is smaller than its atom; an anion has gained electrons, so it is larger. In most common salts, such as NaCl, KBr and MgO, the cation is the smaller ion.