Basically, there is more than one way to categorize a crystal. The two most common methods are to group them according to their crystalline structure and to to group them according to their chemical/physical properties.
Crystal Categories by Shape There are seven crystal lattice systems that classifying crystals according to their atomic lattice or structure.
The atomic lattice is a three dimensional network of atoms that are arranged in a symmetrical pattern. The shape of the lattice determines not only which crystal system the stone belongs to, but all of its physical properties and appearance. In some crystal healing practices the axial symmetry of a crystal is believed to directly influence its metaphysical properties. For example crystals in the Cubic System are believed to be grounding, because the cube is a symbol of the element Earth.
1. Cubic or Isometric - not always cube shaped! You'll also find octahedrons (eight faces) and dodecahedrons (10 faces).
2. Tetragonal - similar to cubic crystals, but longer along one axis than the other, forming double pyramids and prisms.
3. Orthorhombic - like tetragonal crystals except not square in cross section (when viewing the crystal on end), forming rhombic prisms or dipyramids (two pyramids stuck together).
4. Hexagonal - six-sided prisms. When you look at the crystal on-end, the cross section is a hexagon.
5. Trigonal - possess a single 3-fold axis of rotation instead of the 6-fold axis of the hexagonal division.
6. Triclinic - usually not symmetrical from one side to the other, which can lead to some fairly strange shapes.
7. Monoclinic - like skewed tetragonal crystals, often forming prisms and double pyramids.
Crystal Categories by Properties Generally, there are four main categories of crystals, as grouped by their chemical and physical properties.
1.Ionic Crystals Ionic crystals are hard, have high melting points, and are brittle. When they melt, the resulting liquids conduct electricity well. These properties reflect the strong attractive forces between ions of opposite charge as well as the repulsions that occur when ions of like charge are placed near each other.
They are brittle and tend to shatter into smaller crystals when stressed. When a crystal is hammered or stressed, ions with like charges are forced into close proximity. The crystal then literaly self destructs because of electrostatic repulsion.
2. Molecular Crystals Molecular crystals are solids in which the lattice sites are occupied either by atoms - as in solid argon or krypton - or by molecules - as in solid CO2, SO2, or H2O. Such solids tend to be soft and have low melting points because the particles in the solid experience relatively weak intermolecular attractions. The crystals are soft because little effort is needed to separate the particles or to move them past each other. The solid melts at low temperatures because the particles need little kinetic energy to break away from the solid.
If the crystals contain only individual atoms, as in solid argon or krypton, or if they are composed of non-polar molecules, as in naphthalene, the only attractions between the molecules are the London forces. In crystals containing polar molecules, such as sulphur dioxide, the major forces that hold the particles together are dipole-dipole attractions. In crystals such as water the primary forces of attractions are due to hydrogen bonding.
3. Covalent Crystals Covalent crystals are solids in which the lattice points are occupied by atoms that are covalently bonded to other atoms at neighbouring lattice sites. The result is a crystal that is essentially one gigantic molecule. These solids are sometimes called network solids because of the interlocking network of covalent bonds extending throughout the crystal in all directions.
Covalent crystals tend to be hard and to have very high melting points because of the strong attractions between covalently bonded atoms. They do not conduct electricity because the electrons are bound too tightly to the bonds. Other examples of covalent crystals are quartz (SiO2 - typical grains of sand) and silicon carbide (SiC - a common abrasive used in sandpaper).
4. Metallic crystals Metallic crystals have properties that are quite different from those of the other three types of crystals above. They conduct heat and electricity well, and they have the lustre that we characteristically associate with metals. A number of different models have been developed to describe metals. The simplest one views the crystal as having positive ions at the lattice positions which are surrounded by electrons in a cloud that spreads throughout the entire solid.
The electrons in this cloud belong to no single positive ion, but rather to the crystal as a whole. Because the electrons aren't localized on any one atom, they are free to move easily, which accounts for the electrical conductivity of metals. By their movement, the electrons can also transmit kinetic energy rapidly thorough the solid, so metals are also good conductors of heat. Because the electrons are free to move easily, even the lattice points are moveable or deformable and this helps explain the malleability of metals.
Even though the lattice points and electrons are free to move there are some strong attractive forces at work and these help explain ductility in some metals. This model also explains the lustre of metals. When light shines on the metal, the loosely held electrons vibrate easily and readily reemit the light with essentially the same frequency and intensity.
It is not possible to make simple generalizations about the melting points of metals. Some have very high melting points, like tungsten, whereas others like mercury have quite low melting points. To some degree the melting point depends on the charge of the positive ions in the metallic crystals. The ions of Group IA tend to exist as cations with a +1 charge, and they are only weakly attracted to the "electron sea" that surrounds them. Atoms of Group IIA metals, however, tend to form cations with a +2 charge.
These more highly charged ions are attached more strongly to the surrounding electron sea, so the Group IIA metals have higher melting points than their neighbours in Group IA. For example magnesium melts at 650oC but sodium melts at 98oC. Tungsten which has a very high melting point, must have very strong attractions between their atoms, which suggests that there probably is some covalent bonding in them as well.