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Look at grains of table salt under a magnifying glass and you’ll see tiny cubes. That cubic shape isn’t a coincidence: it reflects the perfectly ordered way sodium and chloride ions are stacked inside the crystal. Sodium chloride was one of the very first crystal structures ever worked out, and it remains the model for understanding ionic lattices. This article explains how the ions are arranged, how chemists describe the arrangement, why other ionic compounds adopt different patterns, and how the structure explains the properties of salts.
What is an ionic lattice?
An ionic compound is made of positive and negative ions held together by electrostatic attraction (see ionic bonding explained). This attraction acts equally in all directions, so each ion attracts as many oppositely charged ions around it as will fit. The result isn’t a molecule but a giant ionic lattice: a regular, repeating 3D arrangement of ions that extends throughout the crystal.
A few key points:
- There are no molecules in an ionic lattice. The formula NaCl shows only the ratio of ions (1 : 1).
- Each ion is surrounded by ions of the opposite charge, and its nearest same-charge neighbours are further away. This maximises attraction and minimises repulsion.
- The same pattern repeats billions of times in every direction. A grain of salt 0.5 mm across contains several times 10¹⁸ ions.
The sodium chloride (rock-salt) structure
In sodium chloride:
- Each Na⁺ ion is surrounded by six Cl⁻ ions, arranged at the corners of an octahedron (above, below, left, right, front, back).
- Each Cl⁻ ion is surrounded by six Na⁺ ions, also octahedrally.
The number of nearest oppositely charged neighbours is called the coordination number. In NaCl it’s 6 : 6 — six for each type of ion.
You can picture the structure as a 3D chessboard: ions alternate Na⁺, Cl⁻, Na⁺, Cl⁻ along every row, column and stack. Because the arrangement is cubic, crystals tend to grow and break along cubic faces — hence the little cubes.
This arrangement is so common that it’s called the rock-salt structure. Many other compounds share it, including KCl, NaF, LiCl, MgO, CaO and AgCl.
The unit cell
Crystallographers describe a lattice by its unit cell: the smallest repeating box that, stacked in three dimensions, builds up the whole crystal.
The unit cell of NaCl is a cube with:
- Cl⁻ ions at each corner and at the centre of each face (a face-centred cubic pattern);
- Na⁺ ions at the centre of each edge and at the very centre of the cube (another face-centred cubic pattern, offset by half a cube edge).
Counting ions in the unit cell requires care, because ions on corners, edges and faces are shared with neighbouring cells:
| Position | Shared between | Fraction in this cell |
|---|---|---|
| Corner | 8 cells | ⅛ |
| Edge | 4 cells | ¼ |
| Face | 2 cells | ½ |
| Body centre | 1 cell | 1 |
Chloride ions: 8 corners × ⅛ + 6 faces × ½ = 1 + 3 = 4 Sodium ions: 12 edges × ¼ + 1 centre × 1 = 3 + 1 = 4
So each unit cell contains 4 Na⁺ and 4 Cl⁻ — a 1 : 1 ratio, matching the formula NaCl. ✓
Worked example: density from the unit cell
The edge length of the NaCl unit cell is 564 pm (5.64 × 10⁻⁸ cm).
- Volume = (5.64 × 10⁻⁸ cm)³ = 1.794 × 10⁻²² cm³
- Mass of 4 NaCl units = 4 × 58.44 g mol⁻¹ ÷ 6.022 × 10²³ mol⁻¹ = 3.882 × 10⁻²² g
- Density = 3.882 × 10⁻²² ÷ 1.794 × 10⁻²² = 2.16 g cm⁻³
The measured density of sodium chloride is 2.17 g cm⁻³. A structure worked out from X-rays predicts a bulk property almost exactly. (You can check the molar mass of NaCl with the molar mass calculator.)
How X-rays revealed the structure
In 1913, William Henry Bragg and his son William Lawrence Bragg shone X-rays through crystals of sodium chloride and other salts. Because the spacing between layers of ions is similar to the wavelength of X-rays, the crystal diffracts them into a pattern of spots. From the angles, the Braggs worked out the arrangement of ions. They showed that there were no NaCl molecules — each sodium was equally surrounded by six chlorines. This was a striking confirmation of the ionic model, and it earned them the 1915 Nobel Prize in Physics. X-ray crystallography has since been used to determine the structures of everything from minerals to DNA and proteins.
Why different compounds have different lattices
Not every ionic compound has the rock-salt structure. The arrangement depends mainly on:
- The ratio of ions (set by their charges).
- The relative sizes of the ions, expressed as the radius ratio r₊/r₋ (see ionic radius).
A larger cation can fit more anions around it without them touching each other, so it can have a higher coordination number.
| Structure | Example | Coordination | Typical radius ratio |
|---|---|---|---|
| Caesium chloride | CsCl, CsBr | 8 : 8 (cubic) | above about 0.73 |
| Rock salt | NaCl, MgO, KBr | 6 : 6 (octahedral) | about 0.41–0.73 |
| Zinc blende | ZnS, CuCl | 4 : 4 (tetrahedral) | about 0.23–0.41 |
| Fluorite (1 : 2 ratio) | CaF₂ | 8 : 4 | — |
In caesium chloride, the Cs⁺ ion is large enough to be surrounded by eight Cl⁻ ions at the corners of a cube, with a Cs⁺ at the centre. In fluorite, CaF₂, each Ca²⁺ is surrounded by eight F⁻ ions, but each F⁻ is surrounded by only four Ca²⁺ — because there are twice as many fluoride ions as calcium ions.
The radius-ratio rules are guidelines, not laws. Real compounds sometimes break them, especially when bonding has some covalent character.
How the lattice explains properties
| Property | Explanation from the lattice |
|---|---|
| High melting point (NaCl 801 °C) | Every ion is held by many strong attractions in all directions |
| Hard | Ions are held firmly in place |
| Brittle — shatters when hit | A blow shifts one layer of ions; like charges line up and repel, splitting the crystal |
| Cleaves along flat faces | The cubic lattice has natural planes |
| Doesn’t conduct as a solid | Ions can’t move out of their lattice positions |
| Conducts when molten or dissolved | Ions become free to move and carry charge |
Compounds with higher-charged or smaller ions have stronger attractions and higher melting points: MgO (Mg²⁺ and O²⁻, same rock-salt structure) melts at about 2,850 °C, over 2,000 °C higher than NaCl. The energy released when the lattice forms from gaseous ions is the lattice energy (see lattice energy).
Common misconceptions
- “Salt is made of NaCl molecules.” It’s a lattice of ions; NaCl is the ratio.
- “Each Na⁺ is bonded to one Cl⁻.” Each is attracted equally to six neighbours.
- “All ionic compounds have the same structure.” Ion sizes and charges lead to different lattices.
- “Ionic solids conduct because they contain charged particles.” The ions must be free to move.
Key takeaways
- An ionic lattice is a regular 3D arrangement of oppositely charged ions; there are no molecules.
- In NaCl, each ion has six nearest neighbours of opposite charge (6 : 6, octahedral).
- The NaCl unit cell contains 4 Na⁺ + 4 Cl⁻; its dimensions predict density accurately.
- Ion size and charge decide the structure: CsCl (8 : 8), NaCl (6 : 6), ZnS (4 : 4), CaF₂ (8 : 4).
- The lattice explains high melting points, brittleness and conductivity only when molten or dissolved.
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