**For a square lattice as shown, the coordination number is
4 (the number of circles touching any individual).**

**The empty spaces between the atoms are interstitial sites. Since
each is touched by 4 circles, the interstitial sites are also 4-coordinate,
and the geometry is also square.** **In this example there are
the same number of interstitial sites as circles.**

**Also because the atoms attract one another, there is a tendency to
squeeze out as much empty space as possible.** **The packing efficiency
is the fraction of the crystal (or unit cell) actually occupied by the
atoms. It must always be less than 100% because it is impossible to pack
spheres (atoms are usually spherical) without having some empty space between
them.**

**P.E. = (area of contents) / (area of unit cell)**

**Regarding our square lattice of circles, we can calculate the packing
efficiency (PE) for this particular lattice as follows:**

**The interstitial sites must occupy 100% - 78.54% = 21.46%.**

**Let us compare this with a hexagonal lattice of circles. We
note that :**

**The larger coordination number (more bonds) and greater packing efficiency
suggest that this would be a more stable lattice than the square one.**

If we compare the square and hexagonal lattices, we see that both are made of columns of circles. However, in the hexagonal lattice every other column is shifted allowing the circles to nestle into the empty spaces. Thus, the higher packing efficiency.

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**Structure of Crystals**
**Crystal Lattices**
**Unit Cells**
**From Unit Cell to Lattice**
**From Lattice to Unit Cell**
**Stoichiometry**
**Packing
& Geometry**
**Simple Cubic Metals**
**Close Packed Structures**
**Body Centered Cubic**
**Cesium Chloride**
**Sodium Chloride**
**Rhenium Oxide**
**Niobium Oxide**

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