A covalent bond is a chemical bond that involves the sharing of pairs of electrons between atoms. This sharing results in a stable balance of attractive and repulsive forces between those atoms. Covalent solids are a class of extended-lattice compounds in which each atom is covalently bonded to its nearest neighbors. This means that the entire crystal is, in effect, one giant molecule. The extraordinarily strong binding forces that join all adjacent atoms account for the extreme hardness of these solids. They cannot be broken or abraded without breaking a large number of covalent chemical bonds. Similarly, a covalent solid cannot "melt" in the usual sense, since the entire crystal is one giant molecule. When heated to very high temperatures, these solids usually decompose into their elements.
Another property of covalent network solids is poor electrical conductivity, since there are no delocalized electrons. When molten, unlike ionic compounds, the substance is still unable to conduct electricity, since the macromolecule consists of uncharged atoms rather than ions. (This is also contrary to most forms of metallic bonds. )
A Case Study: Allotropes of Carbon
Graphite is an allotrope of carbon. In this allotrope, each atom of carbon forms three covalent bonds, leaving one electron in each outer orbital delocalized, creating multiple "free electrons" within each plane of carbon. This grants graphite electrical conductivity. Its melting point is high, due to the large amount of energy required to rearrange the covalent bonds. It is also quite hard because of the strong covalent bonding throughout the lattice. However, because of the planar bonding arrangements between the carbon atoms, the layers in graphite can be easily displaced, so the substance is malleable. This explains the use of graphite in pencils, where the layers of carbon are "shedded" on paper (pencil "lead" is typically a mixture of graphite and clay, and was invented for this use in 1795). Graphite is generally insoluble in any solvent due to the difficulty of solvating a very large molecule.
Diamond is also an allotrope of carbon. Diamond is the hardest material known, defining the upper end of the 1-10 scale known as Moh's hardness scale. Diamond cannot be melted; above 1700 °C it is converted to graphite, the more stable form of carbon. The diamond unit cell is face-centered cubic and contains eight carbon atoms.
Boron nitride (BN) is similar to carbon because it exists as a diamond-like cubic polymorph as well as in a hexagonal form similar to graphite. Cubic boron nitride is the second-hardest material after diamond, and it is used in industrial abrasives and cutting tools. Recent interest in boron nitride has centered on its carbon-like ability to form nanotubes and related nanostructures.
Silicon carbide (SiC) is also known as carborundum. Its structure is very much like that of diamond, with every other carbon replaced by silicon. When heated at atmospheric pressure, it decomposes at 2700 °C, but it has never been observed to melt. Structurally, silicon carbide is very complex; at least 70 crystalline forms have been identified. Its extreme hardness and ease of synthesis have led to a diversity of applications -- in cutting tools and abrasives, high-temperature semiconductors and other high-temperature applications, the manufacturing of specialty steels and jewelry, and many more. Tungsten carbide (WC) is probably the most widely encountered covalent solid, owing to its use in carbide cutting tools and as the material used to make the rotating balls in ball-point pens. It has a high melting point (2870 °C) and a structure similar to that of diamond, although it is slightly less hard. In many of its applications, it is embedded in a softer matrix of cobalt or coated with titanium compounds.