Crystal

Although ionic solids follow similar patterns as described above for metals, the detailed arrangements are more complicated, because the positioning of two different types of ions, cations and anions, must be considered. In general, it is the larger ion (usually, the anion) that determines the overall packing and layering, while the smaller ion fits in the holes (spaces) that occur throughout the layers.

Two types of holes occupied by cations exist (in close-packed ionic structures.). These are named tetrahedral and octahedral. An ion in a tetrahedral site would be in contact with four ions of opposite charge, which, if linked by imaginary lines, produces a tetrahedron. An ion in an octahedral site would be in contact with six ions of opposite charge, producing an octahedron. The number of oppositely charged ions in contact with a given ion is called its coordination number (CN). Therefore, an ion in a tetrahedral site has a coordination number of four; an ion in an octahedral site has a coordination number of six. The total number of octahedral holes is the same as the number of close-packed ions, whereas there are twice as many tetrahedral holes as close-packed atoms. Because tetrahedral holes are smaller, they are occupied only when the ratio of the smaller ion's radius to the radius of the larger ion is very small. As the radiusratio of the smaller ion to the larger ion becomes greater, the smaller ion no longer fits into tetrahedral holes, but will fit into octahedral holes.

These principles can be illustrated by several examples. The repeating structural unit of crystalline sodium chloride (table salt) is the face-centered cubic unit cell. The larger chloride ions are cubic close-packed (ABCABC layering pattern). The radius ratio of sodium ion to chloride ion is about 0.6, so the smaller sodium ions occupy all the octahedral sites. Chloride and sodium ions both have coordination numbers of six. This structure occurs frequently among ionic compounds and is called the sodium chloride or rock salt structure.

In the sphalerite (or zinc blende) crystalline form of zinc sulfide, the larger sulfide ions are cubic close-packed (ABCABC layering), giving a face-centered cubic unit cell. The small zinc ions occupy tetrahedral sites. However, the number of tetrahedral holes is twice the number of sulfide ions, whereas the number of zinc ions is equal to the number of sulfide ions. Therefore, zinc ions occupy only half of the tetrahedral holes. In the wurtzite structure, another crystalline form of zinc sulfide, the sulfide ions are hexagonally close-packed, (ABAB layering), giving a hexagonal unit cell. Again, the zinc ions occupy half the tetrahedral sites.

Another common structure, the fluorite structure, is often observed for ionic compounds which have twice as many anions as cations, and in which the cations are larger than the anions. The structure is named after the compound, calcium fluoride, in which the calcium ions are cubic close-packed, with fluoride in all the tetrahedral sites.

As discussed for metals, many compounds have structures that do not involve close-packing. For example, in the cesium chloride structure, the larger chloride ions are arranged in primitive cubes, with cesium ions occupying positions at the cube centers.

Many other structures are observed for ionic compounds. These involve similar packing arrangements as described above, but vary in number and types of occupied holes, and the distribution of ions in compounds having more than two types of cation and/or anion.