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Atoms in covalent solids are covalently bonded with their neighbors, creating, in effect, one giant molecule.
Discuss the properties of covalent crystals or network solids
Covalent (or network) solids are extended-lattice compounds, in which each atom is covalently bonded to its nearest neighbors. Because there are no delocalized electrons, covalent solids do not conduct electricity.
The rearranging or breaking of covalent bonds requires large amounts of energy; therefore, covalent solids have high melting points.
Covalent bonds are extremely strong, so covalent solids are very hard. Generally, covalent solids are insoluble due to the difficulty of solvating very large molecules.
Diamond is the hardest material known, while cubic boron nitride (BN) is the second-hardest. Silicon carbide (SiC) is very structurally complex and has at least 70 crystalline forms.
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-latticecompounds 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. )
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 and Graphite: Two Allotropes of Carbon
These two allotropes of carbon are covalent network solids which differ in the bonding geometry of the carbon atoms. In diamond, the bonding occurs in the tetrahedral geometry, while in graphite the carbons bond with each other in the trigonal planar arrangement. This difference accounts for the drastically different appearance and properties of these two forms of carbon.
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.
Diamond Unit Cell
The diamond unit cell is face-centered cubic and contains 8 carbon atoms. The four darkly shaded atoms are contained within the cell and are completely bonded to other members of the cell. Two of the carbons nearest the viewer are shown as open circles in order to more clearly reveal the bonding arrangement. The other carbon atoms (six in the faces and four at the corners) have some bonds that extend to atoms in other cells.
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. Silicon carbide exists in about 250 crystalline forms. It is used mostly in its synthetic form because it is extremely rare in nature. It is found in a certain type of meteorite that is thought to originate outside of our solar system. The first light-emitting diodes (LEDs) used in high-efficiency lighting were based on SiC
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.
Silicon carbide is an extremely rare mineral, and in nature is is mostly found in a certain type of meteorite.