an octahedral interstice
Between the tightly packed anions of a ceramic there are gaps, and the roomier of the two kinds is the octahedral interstice. Picture a small cation sitting at the centre with six anions around it: four in a square around its waist, one directly above, one directly below. Join the six centres and you trace an octahedron, a shape like two square pyramids glued base to base, which is where the name comes from. The cation is said to have a coordination number of 6.
How big a cation fits? Geometry gives a clean answer: an octahedral hole exactly cradles a cation whose radius is 0.414 times the anion radius. So when the radius ratio r(cation)/r(anion) lands between about 0.414 and 0.732, six-fold octahedral coordination is favoured; below that the cation rattles and prefers a smaller tetrahedral hole, above it the cation is large enough to gather eight neighbours. In a face-centred-cubic anion array the octahedral holes sit at the middle of each cell edge and one at the body centre, giving exactly four per cell, one for every anion.
Whether those octahedral holes are all filled, partly filled, or empty is the fingerprint of a structure. Fill every one in an FCC oxygen array and you get rock salt (MgO, NiO); fill two-thirds in an HCP oxygen array and you get corundum (Al2O3); fill half in FCC oxygen and you are on the way to spinel. Honest caveat: the 0.414 rule assumes hard charged spheres; strong covalent or directional bonding, and the real softness of ions, mean the boundaries are fuzzy and exceptions are common.
In magnesium oxide each Mg2+ ion sits in an octahedral hole ringed by six O2- ions, and each O2- is likewise ringed by six Mg2+, the tidy 6:6 coordination that defines the rock-salt structure.
Six anions around a cation: coordination number 6.
Do not confuse the interstice with the ion in it. The octahedral hole is the empty pocket defined by six surrounding anions; a cation may fill it, or it may stay vacant, as many do in corundum and spinel.