|Name, Symbol, Number||silicon, Si, 14|
|Group, Period, Block||14, 3, p|
|Appearance||as coarse powder,|
|Standard atomic weight||28.0855(3) g·mol−1|
|Electron configuration||[Ne] 3s2 3p2|
|Electrons per shell||2, 8, 4|
|Density (near r.t.)||2.33 g·cm−3|
|Liquid density at m.p.||2.57 g·cm−3|
|Melting point||1687 K
(1414 °C, 2577 °F)
|Boiling point||3538 K
(3265 °C, 5909 °F)
|Heat of fusion||50.21 kJ·mol−1|
|Heat of vaporization||359 kJ·mol−1|
|Heat capacity||(25 °C) 19.789 J·mol−1·K−1|
|Crystal structure||Face-centered cubic|
|Electronegativity||1.90 (Pauling scale)|
|1st: 786.5 kJ·mol−1|
|2nd: 1577.1 kJ·mol−1|
|3rd: 3231.6 kJ·mol−1|
|Atomic radius||110 pm|
|Atomic radius (calc.)||111 pm|
|Covalent radius||111 pm|
|Van der Waals radius||210 pm|
|Thermal conductivity||(300 K) 149 W·m−1·K−1|
|Thermal expansion||(25 °C) 2.6 µm·m−1·K−1|
|Speed of sound (thin rod)||(20 °C) 8433 m/s|
|Young's modulus||150 GPa|
|Bulk modulus||100 GPa|
|CAS registry number||7440-21-3|
|Band gap energy at 300 K||1.12 eV|
Silicon (chemical element symbol Si, atomic number 14) is a member of a group of chemical elements classified as metalloids. It is less reactive than its chemical analog carbon. It is the eighth most common element in the universe (by mass) and is the second most abundant element (after oxygen) in the Earth's crust, making up 25.7 percent of the crust by mass. It occasionally occurs as the pure free element in nature, but is more widely distributed in dusts, planetoids, and planets as various forms of silicon dioxide or silicate.
Various biological systems contain silicon as an essential element. Although only tiny traces of it appear to be required by animals, it is much more important for the metabolism of plants, particularly many grasses. Also, silicic acid (a family of chemical compounds of silicon, hydrogen, and oxygen) forms the basis of the array of protective shells of diatoms.
Silicon has many industrial uses. Elemental silicon is the principal component of most semiconductor devices, particularly integrated circuits or "microchips." Given its importance in semiconductors and high-tech devices, its name has been used for the high-tech region known as Silicon Valley in California. In the form of silica and silicates, silicon forms useful glasses, cements, and ceramics. It is also a component of silicones, a group of various synthetic plastic substances made of silicon, oxygen, carbon, germanium, and hydrogen.
Given that some properties of silicon are similar to those of carbon, some individuals have proposed the possibility of silicon-based living organisms. This possibility, however, seems remote for a variety of reasons, including the absence of a "silicon cycle" (analogous to the carbon cycle), the absence of an appropriate solvent for silicon compounds (analogous to water that dissolves organic compounds), and the inability of silicon to form the diversity of compounds required for living systems.
Measured by mass, silicon makes up 25.7 percent of the Earth's crust and is the second most abundant element on Earth, after oxygen. Pure silicon crystals are only occasionally found in nature; they can be found as inclusions with gold and in volcanic exhalations. Silicon is usually found in the form of silicon dioxide (also known as silica), and silicate.
Silica occurs in minerals consisting of (practically) pure silicon dioxide in different crystalline forms. Sand, amethyst, agate, quartz, rock crystal, chalcedony, flint, jasper, and opal are some of the forms in which silicon dioxide appears. They are known as "lithogenic" (as opposed to "biogenic") silicas.
Silicon also occurs as silicates (various minerals containing silicon, oxygen, and one or other metal). These minerals occur in clay, sand, and various types of rock such as granite and sandstone. Asbestos, feldspar, clay, hornblende, and mica are a few of the many silicate minerals.
Silicon is a principal component of aerolites, which are a class of meteoroids, and also is a component of tektites, a natural form of glass.
Etymology and history
The name silicon is derived from the Latin word, silex, meaning "flint" or "hard stone," corresponding to the materials now called "silica" or "silicates." It was first identified by Antoine Lavoisier in 1787, as a component of silex, but Humphry Davy (in 1800) mistook it as a compound. In 1811, Gay-Lussac and Louis Jacques Thénard probably prepared impure amorphous silicon through the heating of potassium with silicon tetrafluoride. The first person to identify it as an element was Jöns Jakob Berzelius, in 1823. In the following year, Berzelius prepared amorphous silicon using approximately the same method as Gay-Lussac. He also purified the product by repeated washing.
In the periodic table, silicon is located in group 14 (former group 4A), between carbon and germanium. In addition, it lies in period 3, between aluminum and phosphorus. Elemental silicon has a gray color and a metallic luster, which increases with the size of the crystal.
The electronic configuration in the outermost shell of a silicon atom is the same as that of a carbon atom—both types of atoms have four bonding electrons. Consequently, both elements are tetravalent (each atom binding up to four other atoms) and share some chemical properties. Both are semiconductors, readily donating or sharing their four outer electrons, allowing for various forms of chemical bonding.
Silicon is similar to glass in that it is strong but brittle and prone to chipping. Although it is a relatively inert element, silicon reacts with halogens and dilute alkalis. Most acids (except for some hyper-reactive combinations of nitric acid and hydrofluoric acid) do not affect it.
Silicon is widely used in semiconductors because it remains a semiconductor at higher temperatures than the semiconductor germanium, and because its native oxide is easily grown in a furnace and forms a better semiconductor/dielectric interface than almost all other material combinations. The electrical resistance of single-crystal silicon significantly changes under the application of mechanical stress, due to what is called the "piezoresistive effect."
Silicon has many known isotopes, with mass numbers ranging from 22 to 44. Of these, the stable isotopes are 28Si (the most abundant isotope, at 92.23 percent), 29Si (4.67 percent), and 30Si (3.1 percent). In addition, 32Si is a radioactive isotope produced by argon decay. Its half-life has been determined to be approximately 170 years (0.21 MeV), and it decays by beta emission to 32P (which has a half-life of 14.29 days), and then to 32S.
Examples of silicon compounds:
- Silane (SiH4)
- Silicic acid (H4SiO4)
- Silicon carbide (SiC)
- Silicon dioxide (SiO2)
- Silicon tetrachloride (SiCl4)
- Silicon tetrafluoride (SiF4)
- Trichlorosilane (HSiCl3)
Silicon is commercially prepared by the reaction of high-purity silica with wood, charcoal, and coal, in an electric arc furnace using carbon electrodes. At temperatures over 1900 °C, the carbon reduces the silica to silicon according to the chemical equation
- SiO2 + C → Si + CO2.
Liquid silicon collects in the bottom of the furnace, and is then drained and cooled. The silicon produced via this process is called "metallurgical grade silicon" and is at least 98 percent pure. Using this method, silicon carbide, SiC, can form. However, provided the amount of SiO2 is kept high, silicon carbide may be eliminated, as explained by this equation:
- 2 SiC + SiO2 → 3 Si + 2 CO.
In 2005, metallurgical grade silicon cost about $ 0.77 per pound ($1.70/kg).
The use of silicon in semiconductor devices demands a much greater purity than afforded by metallurgical grade silicon. Historically, a number of methods have been used to produce high-purity silicon.
Early silicon purification techniques were based on the fact that if silicon is melted and re-solidified, the last parts of the mass to solidify contain most of the impurities. The earliest method of silicon purification, first described in 1919, and used on a limited basis to make radar components during World War II, involved crushing metallurgical grade silicon and then partially dissolving the silicon powder in an acid. When crushed, the silicon cracked so that the weaker impurity-rich regions were on the outside of the resulting grains of silicon. As a result, the impurity-rich silicon was the first to be dissolved when treated with acid, leaving behind a more pure product.
In zone melting, also called zone refining, the first silicon purification method to be widely used industrially, rods of metallurgical grade silicon are heated to melt at one end. Then, the heater is slowly moved down the length of the rod, keeping a small length of the rod molten as the silicon cools and re-solidifies behind it. Since most impurities tend to remain in the molten region rather than re-solidify, when the process is complete, most of the impurities in the rod will have been moved into the end that was the last to be melted. This end is then cut off and discarded, and the process repeated if a still higher purity is desired.
Today, silicon is instead purified by converting it to a silicon compound that can be more easily purified than silicon itself, and then converting that silicon element back into pure silicon. Trichlorosilane is the silicon compound most commonly used as the intermediate, although silicon tetrachloride and silane are also used. When these gases are blown over silicon at high temperature, they decompose to high-purity silicon.
At one time, DuPont produced ultra-pure silicon by reacting silicon tetrachloride with high-purity zinc vapors at 950°C, producing silicon according to the chemical equation
- SiCl4 + 2 Zn → Si + 2 ZnCl2.
However, this technique was plagued with practical problems (such as the zinc chloride byproduct solidifying and clogging lines) and was eventually abandoned in favor of the Siemens process.
In the Siemens process, high-purity silicon rods are exposed to trichlorosilane at 1150°C. The trichlorosilane gas decomposes and deposits additional silicon onto the rods, enlarging them according to chemical reactions like
- 2 HSiCl3 → Si + 2 HCl + SiCl4.
Silicon produced from this and similar processes is called polycrystalline silicon. Polycrystalline silicon typically has impurity levels of less than 10−9.
In 2006, Renewable Energy Corporation (REC) announced construction of a plant based on fluidized bed technology using silane.
- 3SiCl4 + Si + 2H2 → 4HSiCl3
- 4HSiCl3 → 3SiCl4 + SiH4
- SiH4 → Si + 2H2
The majority of silicon crystals grown for device production are produced by the Czochralski process (CZ-Si), because it is the cheapest method available and is capable of producing large crystals. However, silicon single-crystals grown by the Czochralski method contain impurities because the crucible that contains the melt dissolves. For certain electronic devices, particularly those required for high-power applications, silicon grown by the Czochralski method is not pure enough. For these applications, float-zone silicon (FZ-Si) can be used instead. It is worth mentioning, though, that it is difficult to grow large crystals using the float-zone method. Today, all the dislocation-free silicon crystals used in semiconductor industry with diameter 300mm or larger are grown by the Czochralski method, with purity level significantly improved.
As the second most common element on earth, silicon is a very useful element that is vital to many human industries and impacts much of modern life. For instance, it is a major component of glass, concrete, and cements of many kinds. In addition, one of its most valuable applications lies in that it forms the fundamental substrate in manufacturing electronics devices such as integrated circuits and power transistors. Further, the element and its compounds find widespread use in explosives and pyrotechnics.. Silicon is also used in mechanical seals, caulking compounds, and high-temperature, silicon-based greases.
- The largest application of pure (metallurgical grade) silicon is in aluminum-silicon alloys, often called "light alloys," to produce cast parts, mainly for automotive industry. (This represents about 55% of the world consumption of pure silicon.)
- The second largest application of pure silicon is as a raw material in the production of silicones (about 40% of the world consumption of silicon)
- Pure silicon is also used to produce ultra-pure silicon for electronic and photovoltaic applications:
- Semiconductor: Ultrapure silicon can be doped with other elements to adjust its electrical response by controlling the number and charge (positive or negative) of current carriers. Such control is necessary for transistors, solar cells, microprocessors, semiconductor detectors and other semiconductor devices which are used in electronics and other high-tech applications.
- Photonics: Silicon can be used as a continuous wave Raman laser to produce coherent light. (Though it is ineffective as a light source.)
- LCDs and solar cells: Hydrogenated amorphous silicon is widely used in the production of low-cost, large-area electronics in applications such as LCDs. It has also shown promise for large-area, low-cost thin-film solar cells.
- Steel and cast iron: Silicon is an important constituent of some steels, and it is used in the production process of cast iron. It is introduced as ferrosilicon or silicocalcium alloys.
- Construction: Silicon dioxide or silica in the form of sand and clay is an important ingredient of concrete and brick and is also used to produce Portland cement.
- Pottery/Enamel is a refractory material used in high-temperature material production and its silicates are used in making enamels and pottery.
- Glass: Silica from sand is a principal component of glass. Glass can be made into a great variety of shapes and with a many different physical properties. Silica is used as a base material to make window glass, containers, insulators, and many other useful objects.
- Abrasives: Silicon carbide is one of the most important abrasives.
- Medical materials: Silicones are flexible compounds containing silicon-oxygen and silicon-carbon bonds; they are widely used in applications such as artificial breast implants and contact lenses. Silicones are also used in many other applications.
- Silly Putty was originally made by adding boric acid to silicone oil. Now name-brand Silly Putty also contains significant amounts of elemental silicon. (Silicon binds to the silicone and allows the material to bounce 20 percent higher.)
Different forms of silicon
One can notice the color change in silicon nanopowder. This is caused by the quantum effects which occur in particles of nanometric dimensions. See also Potential well, Quantum dot, and Nanoparticle.
Given that silicon is similar to carbon, particularly in its valency, some have pondered over the possibility of silicon-based life. For instance, A. G. Cairns-Smith has proposed that the first living organisms may have been forms of clay minerals, which were probably based around the silicon atom.
Although there are no known forms of life that rely entirely on silicon-based chemistry, there are some that rely on silicon minerals for specific functions. Some bacteria and other forms of life, such as the protozoa radiolaria, have silicon dioxide skeletons, and the sea urchin has spines made of silicon dioxide. These forms of silicon dioxide are known as biogenic silica. Silicate bacteria use silicates in their metabolism.
Yet, life as it is known today could not have developed based on a silicon biochemistry. The main reason is that life on Earth depends on the carbon cycle: Autotrophic organisms use carbon dioxide to synthesize organic compounds with carbon, which is then used as food by heterotrophic organisms, which produce energy and carbon dioxide from these compounds. If carbon were to be replaced by silicon, there would be a need for a silicon cycle, involving the participation of silicon dioxide. However, unlike carbon dioxide, silicon dioxide is a solid that does not dissolve in water and cannot be transported through living systems by common biological means. Consequently, another solvent would be necessary to sustain silicon-based life forms. It would be difficult (if not impossible) to find another common compound with the unusual properties of water that make it an ideal solvent for carbon-based life.
Larger silicon compounds (silanes) that are analogous to common hydrocarbon chains are generally unstable, owing to the larger atomic radius of silicon and the correspondingly weaker silicon-silicon bond. Silanes decompose readily and often violently in the presence of oxygen, making them unsuitable for an oxidizing atmosphere such as our own. Moreover, unlike carbon, silicon does not have the tendency to form double and triple bonds.
Some silicon rings (cyclosilanes) have been synthesized and are analogous to the cycloalkanes formed by carbon, but the cyclosilanes are rare whereas the cycloalkanes are common. Synthesis of the cyclosilanes suffers from the difficulties inherent in producing any silane compound. On the other hand, carbon will readily form five-, six-, and seven-membered rings by a variety of pathways, even in the presence of oxygen.
Silicon's inability to readily form multiple bonds, long silane chains, and rings severely limits the diversity of compounds that can be synthesized from it. Under known conditions, silicon chemistry simply cannot begin to approach the diversity of organic chemistry, a crucial factor in carbon's role in biology.
Some have construed silicon-based life as existing under a computational substrate. This concept, yet to be explored in mainstream technology, receives ample coverage by science fiction authors.
- ↑ Science Gateway, Phosphorus-32. Retrieved August 11, 2007.
- ↑ U.S. Geological Survey, Mineral Commodity Summaries, Silicon. Retrieved August 11, 2007.
- ↑ Renewable Energy Corporation, Analyst Silicon Field Trip, March 28, 2007. Retrieved August 11, 2007.
- ↑ E.-C. Koch, D. Clement. Special Materials in Pyrotechnics: VI. Silicon—An Old Fuel with New Perspectives. Retrieved August 11, 2007.
- Cotton, F. Albert and Geoffrey Wilkinson. 1980. Advanced Inorganic Chemistry. New York: Wiley. ISBN 0471027758
- Chang, Raymond. 2006. Chemistry. New York: McGraw-Hill Science/Engineering/Math. ISBN 0073221031
- Greenwood, N. N. and A. Earnshaw. 1998. Chemistry of the Elements. Burlington, MA: Butterworth-Heinemann, Elsevier Science. ISBN 0750633654
- Jutzi, Peter and Ulrich Schubert. 2003. Silicon Chemistry: From the Atom to Extended Systems. New York: Wiley. ISBN 3527306471
- Los Alamos National Laboratory. Silicon. Retrieved August 11, 2007.
- Siffert, P. and E. Krimmel, eds. 2004. Silicon: Evolution and Future of a Technology. Berlin: Springer. ISBN 3540405461
- Silicon. WebElements.com. Retrieved August 22, 2007.
- The Mineral Silicon. Amethyst Galleries. Retrieved August 22, 2007.
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