Silicon is a dark-gray, metalloid element, the second most abundant element in Earth's crust (25.7% by mass). Silicon occurs in various forms including silicate minerals and quartz. It has a diamond-like crystal structure, although it can also exist in an amorphous state. Silicon is used in alloys and to make semiconductor components.
Many organosilicon compounds are known and there has been much speculation over the years on the possibility of silicon-based life.
|relative atomic mass||28.086|
|atomic radius||117 pm|
|oxidation states||2, 4, - 4|
|melting point||1,410°C (2,570°F)|
|boiling point||2,355°C (4,271°F)|
Chemistry of silicon
Silicon is quite inert at low temperatures, but when strongly heated in air the surface becomes covered with a layer of oxide. Silicon is insoluble in water and resists the action of most acids, but not hydrofluoric. When boiled with alkaline hydroxides, such as sodium hydroxide (caustic soda), sodium silicate is formed:
Si + 2NaOH + H2O → Na2SiO3 + 2H2
Silicon itself is not very hard, but silicon carbide (known commercially as carborundum), which is obtained by heating a mixture of silicon dioxide (silica) and coke in an electric furnace, is almost as hard as diamond. Crystals of carborundum are hard, chemically inactive, and not decompose until heated to about 2,200°C. Crushed crystals of this compound can be mixed with a binder such as clay and molded into various shapes for use as grindstones and grinding wheels. The blocks and wheels have to be baked subsequently, so that the individual crystals fuse together.
As silicon lies immediately below carbon in the same family in the periodic table, one might expect the compounds of the elements to be similar. This is true to some extent, but there is very little in common physically between the oxide of carbon and silicon. Carbon dioxide is a gas at normal temperatures, whereas silicon dioxide is a hard solid which melts at 1,730°C.
The stability of silica (silicon dioxide) in its crystalline state is attributed to the structure of the molecules. Carbon dioxide, even in the solid state, comprises CO2 units in which two oxygen atoms are joined by double bonds to each carbon. In contrast, the silicon atoms in silicon dioxide are joined by single bonds to four oxygen atoms. The other bond of each oxygen atom is linked to a different silicon atom. Thus each oxygen atom is shared between two silicon atoms. The existence of this macromolecule, as such a giant network of silicon and oxygen atoms is known, accounts for the extraordinary stability of silica (e.g. sand).
In addition to the use of sand and various silica-containing stones in the building trade, silica has many other uses, including the manufacture of glass. A special type of glass can be obtained by heating quartz until it melts, and then working the molten mass in much the same way as glass. Quartz glass fibers are used for suspensions in delicate electrical apparatus. Crucibles, evaporating dishes, and similar apparatus made from fused silica are useful for certain types of reactions. At high temperatures silica will combine with bases and metallic oxides to yield silicates, so there is a limit to the reactions for which silica is suitable.
Silicon tetrachloride (SiCl4) is a colorless fuming liquid, made by reacting chlorine with a mixture of silica and carbon. It is the starting point for preparing a whole range of silicon-containing organic compounds called the silicones.
The basic unit of silicones is a – Si – O – Si – chain of linkages with two organic groups (methyl, ethyl, phenyl, etc) attached to each silicon atom in the chain. Most of the silicones are very stable compounds and are not wetted by water. hence they are used for water-proofing or as water repellents. Depending on the structure of the particular molecules, the silicones may be oils, greases, or resins. Some are used as lubricants (see lubrication) in situations where there are large temperature variations which would render ordinary oils and greases unsuitable.
See also silane.