Some solids, like aluminum and titanium, are light and strong; they conduct electricity and heat well, are a shiny silver or gold color when pure, can be bent without breaking, and are easily malleable. We call them metals. Others, like sulfur or graphite, are crumbly or brittle, do not conduct electricity or heat well, are white, yellow, black, orange, or other colors, and break rather than bend. They are the non-metals. Other solids, like diamond, have other properties entirely, such as extreme hardness, no electrical conductivity, and high thermal conductivity. What causes such wide differences between these solids?
The answer lies in looking at the molecular arrangement of these substances. Based on their crystal structures, chemists have sorted solids into five categories:
Metallic crystals: These crystals are composed of atoms locked into a dense and rigid matrix. Their electrons, however, are free to move around in the crystal, leading to the notion of metals as atoms awash in a "sea of electrons." The mobile electrons allow both heat and electric current to be rapidly distributed through the crystal, explaining these properties. Due to the fairly strong attraction between atoms and electrons, these solids have usually high melting points, but some do not (like mercury, a liquid metal at room temperature). These crystals, due again to the mobile electrons, are fairly flexible and can be bent or shaped without destroying the structure. Some are very hard, such as titanium, while others are soft (like lead) or brittle (metals are obvious examples of metallic crystals. They can be represented simply by the name of the element composing the crystal, such as Ti(s) for titanium.
Ionic crystals: Ionic crystals are held together by the attraction between positively-charged atoms or molecules (cations) and a negatively-charged ones (anions). In most ionic solids, an alternating pattern of anion-cation-anion-cation is followed in all three dimensions (like a checkerboard in 3D), so each cation ends up being surrounded by six anions, and vice versa. The attractions holding the crystal together are very strong; the energy required to remove a unit from the crystal (the lattice energy) is large; this energy is also given off when the crystal forms. Melting the crystal is hard; ionic solids have very high melting points. Conversely, most ionic solids are very soluble in water.
Since the electrons are held by the cation, ionic solids do not conduct (they do conduct when melted, however). They are brittle because just a slight shock to the crystal can offset the arrangements, causing repulsion instead of attraction between the charged atoms or molecules and shattering the crystal, as shown in the Flash movie to the right.
An example of an ionic solid is table salt, NaCl. Since each atom or molecule is joined with many others, discernable units of ionic solids are not present (we can't pick out one specific Na-Cl pair, since there are so many), so the crystal is represented by the smallest unit that can be repeated to "fill in" the structure; for example, Na-Cl can be repeated over and over to form the 3D-checkerboard pattern. This "smallest identifiable" unit, as you may remember, is called the empirical formula.
Molecular crystals: These solids are held together by weak attractions called van der Waals forces, including dipole interactions, hydrogen bonding, and dispersion forces (all of which will be explained a little later). Most molecular crystals have low melting points (in fact, many are liquids or gases at room temperature) and are not very resistant to bending or deformation. They almost never good conductors of electricity or heat, and can have a variety of physical appearances. Ice, dry ice, and solid methanol are examples of molecular crystals. Since they are composed of one type of molecule, simply identifying that molecule will identify the solid (such as H2O(s) for ice).
Network crystals: Also known as covalent crystals, these are solids in which every atom is bonded to several others. These crystals may be one-dimensional, in which case all the bonds are along one axis; two-dimensional, when all the bonds are in a plane; and three-dimensional, when the bonds extend through all three dimensions. One-dimensional network solids are rare. A typical two-dimensional crystal is graphite, which forms "sheets" that are only weakly attracted to each other, and so can slide easily. A famous three-dimensional network solid is diamond, in which the four bonds to each carbon atom are arranged in an extremely strong 3D matrix, making diamond the strongest substance known. Some network solids conduct electricity, and some have excellent heat conductivity. Almost all three-dimensional network solids are very hard, due to their stable, strong bond matrices. Therefore, they have extremely high melting points. Some examples of 3D network solids are diamond and silicates (including silicon dioxide, which makes up sand and quartz). As in ionic solids, no single set of bonding atoms is possible; therefore, an empirical formula must be used.
Amorphous solids: These solids have no overall crystal structure, and simply stick in random orientations. A very common example is glass, which forms no crystal matrix. You can see this by the random cracks that form in glass; if it were a crystal, the fractures would form along the edges of the crystal formations. This is particularly strange, because glass is made up of silicon dioxide, which normally forms a network solid! Amorphous substances have a wide range of chemical properties; in particular, they often melt gradually, slowly softening into liquids.