Silicon

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Silicon,  ₁₄Si

General properties
Pronunciation	/ˈsɪlɪkən/ ​(SIL-ik-ən)
Appearance	crystalline, reflective with bluish-tinged faces
Standard atomic weight (Ar, standard)	[28.084, 28.086] conventional: 28.085
Silicon in the periodic table
Hydrogen Helium Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon Sodium Magnesium Aluminium Silicon Phosphorus Sulfur Chlorine Argon Potassium Calcium Scandium Titanium Vanadium Chromium Manganese Iron Cobalt Nickel Copper Zinc Gallium Germanium Arsenic Selenium Bromine Krypton Rubidium Strontium Yttrium Zirconium Niobium Molybdenum Technetium Ruthenium Rhodium Palladium Silver Cadmium Indium Tin Antimony Tellurium Iodine Xenon Caesium Barium Lanthanum Cerium Praseodymium Neodymium Promethium Samarium Europium Gadolinium Terbium Dysprosium Holmium Erbium Thulium Ytterbium Lutetium Hafnium Tantalum Tungsten Rhenium Osmium Iridium Platinum Gold Mercury (element) Thallium Lead Bismuth Polonium Astatine Radon Francium Radium Actinium Thorium Protactinium Uranium Neptunium Plutonium Americium Curium Berkelium Californium Einsteinium Fermium Mendelevium Nobelium Lawrencium Rutherfordium Dubnium Seaborgium Bohrium Hassium Meitnerium Darmstadtium Roentgenium Copernicium Nihonium Flerovium Moscovium Livermorium Tennessine Oganesson C ↑ Si ↓ Ge aluminium ← silicon → phosphorus
Atomic number (Z)	14
Group	group 14 (carbon group)
Period	period 3
Element category	metalloid
Block	p-block
Electron configuration	[Ne] 3s2 3p2
Electrons per shell	2, 8, 4
Physical properties
Phase at STP	solid
Melting point	1687 K ​(1414 °C, ​2577 °F)
Boiling point	3538 K ​(3265 °C, ​5909 °F)
Density (near r.t.)	2.3290 g/cm3
when liquid (at m.p.)	2.57 g/cm3
Heat of fusion	50.21 kJ/mol
Heat of vaporization	383 kJ/mol
Molar heat capacity	19.789 J/(mol·K)
Vapor pressure P (Pa) 1 10 100 1 k 10 k 100 k at T (K) 1908 2102 2339 2636 3021 3537
Atomic properties
Oxidation states	4, 3, 2, 1[1] −1, −2, −3, −4 ​(an amphoteric oxide)
Electronegativity	Pauling scale: 1.90
Ionization energies	1st: 786.5 kJ/mol 2nd: 1577.1 kJ/mol 3rd: 3231.6 kJ/mol (more)
Atomic radius	empirical: 111 pm
Covalent radius	111 pm
Van der Waals radius	210 pm
Spectral lines
Miscellanea
Crystal structure	​face-centered diamond-cubic
Speed of sound thin rod	8433 m/s (at 20 °C)
Thermal expansion	2.6 µm/(m·K) (at 25 °C)
Thermal conductivity	149 W/(m·K)
Electrical resistivity	2.3×103 Ω·m (at 20 °C)[2]
Band gap	1.12 eV (at 300 K)
Magnetic ordering	diamagnetic[3]
Magnetic susceptibility	−3.9·10−6 cm3/mol (298 K)[4]
Young's modulus	130–188 GPa[5]
Shear modulus	51–80 GPa[5]
Bulk modulus	97.6 GPa[5]
Poisson ratio	0.064–0.28[5]
Mohs hardness	6.5
CAS Number	7440-21-3
History
Naming	after Latin 'silex' or 'silicis', meaning flint
Prediction	Antoine Lavoisier (1787)
Discovery and first isolation	Jöns Jacob Berzelius[6][7] (1823)
Named by	Thomas Thomson (1817)
Main isotopes of silicon
Iso­tope Abun­dance Half-life (t1/2) Decay mode Pro­duct 28Si 92.2% stable 29Si 4.7% stable 30Si 3.1% stable 31Si trace 2.62 h β− 31P 32Si trace 153 y β− 32P

Silicon is a chemical element with symbol Si and atomic number 14. A hard and brittle crystalline solid with a blue-grey metallic lustre, it is a tetravalent metalloid and semiconductor. It is a member of group 14 in the periodic table, along with carbon above it and germanium, tin, and lead below. It is rather unreactive, though less so than germanium, and has a very large chemical affinity for oxygen; as such, it was first prepared and characterized in pure form only in 1823 by Jöns Jakob Berzelius. Its melting and boiling points of 1414 °C and 3265 °C respectively are the second-highest among all the metalloids and nonmetals, being only surpassed by boron (carbon sublimes rather than melts at atmospheric pressure, albeit at a higher temperature than boron).

Silicon is the eighth most common element in the universe by mass, but very rarely occurs as the pure element in the Earth's crust. It is most widely distributed in dusts, sands, planetoids, and planets as various forms of silicon dioxide (silica) or silicates. Over 90% of the Earth's crust is composed of silicate minerals, making silicon the second most abundant element in the Earth's crust (about 28% by mass) after oxygen.

Most silicon is used commercially without being separated, and often with little processing of the natural minerals. Such use includes industrial construction with clays, silica sand, and stone. Silicates are used in Portland cement for mortar and stucco, and mixed with silica sand and gravel to make concrete for walkways, foundations, and roads. They are also used in whiteware ceramics such as porcelain, and in traditional quartz-based soda-lime glass and many other specialty glasses. Silicon compounds such as silicon carbide are used as abrasives and components of high-strength ceramics. Silicon is the basis of the widely used synthetic polymers called silicones.

Elemental silicon also has a large impact on the modern world economy. Most free silicon is used in the steel refining, aluminium-casting, and fine chemical industries (often to make fumed silica). Even more visibly, the relatively small portion of very highly purified elemental silicon used in semiconductor electronics (< 10%) is essential to integrated circuits — most computers, cell phones, and modern technology depend on it.

Silicon is an essential element in biology, although only traces are required by animals. However, various sea sponges and microorganisms, such as diatoms and radiolaria, secrete skeletal structures made of silica. Silica is deposited in many plant tissues, such as in the bark and wood of Chrysobalanaceae and the silica cells and silicified trichomes of Cannabis sativa, horsetails and many grasses.^[8]

History

Jöns Jacob Berzelius, discoverer of silicon

In 1787 Antoine Lavoisier suspected that silica might be an oxide of a fundamental chemical element,^[9] but the chemical affinity of silicon for oxygen is high enough that he had no means to reduce the oxide and isolate the element.^[10] After an attempt to isolate silicon in 1808, Sir Humphry Davy proposed the name "silicium" for silicon, from the Latin silex, silicis for flint, and adding the "-ium" ending because he believed it to be a metal.^[11] Most other languages use transliterated forms of Davy's name, sometimes adapted to local phonology (e.g. German Silizium, Turkish silisyum). A few others use instead a calque of the Latin root (e.g. Russian кремний, from кремень "flint"; Greek πυριτιο from πυρ "fire"; Finnish pii from piikivi "flint").^[12]

In 1811, Gay-Lussac and Thénard are thought to have prepared impure amorphous silicon, through the heating of recently isolated potassium metal with silicon tetrafluoride, but they did not purify and characterize the product, nor identify it as a new element.^[13] Silicon was given its present name in 1817 by Scottish chemist Thomas Thomson. He retained part of Davy's name but added "-on" because he believed that silicon was a nonmetal similar to boron and carbon.^[14] In 1823, Jöns Jacob Berzelius prepared amorphous silicon using approximately the same method as Gay-Lussac (reducing potassium fluorosilicate with molten potassium metal), but purifying the product to a brown powder by repeatedly washing it.^[15] As a result, he is usually given credit for the element's discovery.^[16][17] The same year, Berzelius became the first to prepare silicon tetrachloride; silicon tetrafluoride had already been prepared long before in 1771 by Carl Wilhelm Scheele by dissolving silica in hydrofluoric acid.^[10]

Silicon in its more common crystalline form was not prepared until 31 years later, by Deville.^[18][19] By electrolyzing a mixture of sodium chloride and aluminium chloride containing approximately 10% silicon, he was able to obtain a slightly impure allotrope of silicon in 1854.^[20] Later, more cost-effective methods have been developed to isolate several allotrope forms, the most recent being silicene in 2010.^[21][22] Meanwhile, research on the chemistry of silicon continued; Friedrich Wöhler discovered the first volatile hydrides of silicon, synthesising trichlorosilane in 1857 and silane itself in 1858, but a detailed investigation of the silanes was only carried out in the early 20th century by Alfred Stock, despite early speculation on the matter dating as far back as the beginnings of synthetic organic chemistry in the 1830s.^[23] Similarly, the first organosilicon compound, tetraethylsilane, was synthesised by Charles Friedel and James Crafts in 1863, but detailed characterisation of organosilicon chemistry was only done in the early 20th century by Frederick Kipping.^[10]

Starting in the 1920s, the work of William Lawrence Bragg on X-ray crystallography successfully elucidated the compositions of the silicates, which had previously been known from analytical chemistry but had not yet been understood, together with Linus Pauling's development of crystal chemistry and Victor Goldschmidt's development of geochemistry. The middle of the 20th century saw the development of the chemistry and industrial use of siloxanes and the growing use of silicone polymers, elastomers, and resins. In the late 20th century, the complexity of the crystal chemistry of silicides was mapped, along with the solid-state chemistry of doped semiconductors.^[10]

Because silicon is an important element in high-technology semiconductor devices, many places in the world bear its name. For example, Santa Clara Valley in California acquired the nickname Silicon Valley since the element is the base material used in the semiconductor industry located there. Since then, many other locations have been nicknamed for similar reasons.^[24]

Characteristics

Physical and atomic

Silicon crystallizes in a diamond cubic crystal structure

A silicon atom has fourteen electrons. In the ground state, they are arranged in the electron configuration [Ne]3s²3p². Of these, four are valence electrons, occupying the 3s orbital and two of the 3p orbitals. Like the other members of its group, the lighter carbon and the heavier germanium, tin, and lead, it has the same number of valence electrons as valence orbitals: hence, it can complete its octet and obtain the stable noble gas configuration of argon by forming sp³ hybrid orbitals, forming tetrahedral SiX₄ derivatives where the central silicon atom shares an electron pair with each of the four atoms it is bonded to.^[25] The first four ionisation energies of silicon are 786.3, 1576.5, 3228.3, and 4354.4 kJ/mol respectively; these figures are high enough to preclude the possibility of simple cationic chemistry for the element. Following periodic trends, its single-bond covalent radius of 117.6 pm is intermediate between those of carbon (77.2 pm) and germanium (122.3 pm). The hexacoordinate ionic radius of silicon may be considered to be 40 pm, although this must be taken as a purely notional figure given the lack of a simple Si⁴⁺ cation in reality.^[26]

At standard temperature and pressure, silicon is a shiny semiconductor with a bluish-grey metallic lustre; as typical for semiconductors, its resistivity drops as temperature rises. This arises because silicon has a small energy gap between its highest occupied energy levels (the valence band) and the lowest unoccupied ones (the conduction band). The Fermi level is about halfway between the valence and conduction bands and is the energy at which a state is as likely to be occupied by an electron as not. Hence pure silicon is an insulator at room temperature. However, doping silicon with a pnictogen such as phosphorus, arsenic, or antimony introduces one extra electron per dopant and these may then be excited into the conduction band either thermally or photolytically, creating an n-type semiconductor. Similarly, doping silicon with a group 13 element such as boron, aluminium, and gallium results in the introduction of acceptor levels that trap electrons that may be excited from the filled valence band, creating a p-type semiconductor. Joining n-type silicon to p-type silicon creates a p-n junction with a common Fermi level; electrons flow from n to p, while holes flow from p to n, creating a voltage drop. This p-n junction thus acts as a diode that can rectify alternating current that allows current to pass more easily one way than the other. A transistor is an n-p-n junction, with a thin layer of weakly p-type silicon between two n-type regions. Biasing the emitter through a small forward voltage and the collector through a large reverse voltage allows the transistor to act as a triode amplifier.^[27]

Silicon crystallises in a giant covalent structure at standard conditions, specifically in a diamond cubic lattice. It thus has a high melting point of 1414 °C, as a lot of energy is required to break the strong covalent bonds and melt the solid. It is not known to have any allotropes at standard pressure, but several other crystal structures are known at higher pressures. The general trend is one of increasing coordination number with pressure, culminating in a hexagonal close-packed allotrope at about 40 gigapascals known as Si–VII (the standard modification being Si–I). Silicon boils at 3265 °C: this, while high, is still lower than the temperature at which its lighter congener carbon sublimes (3642 °C) and silicon similarly has a lower heat of vaporisation than carbon, consistent with the fact that the Si–Si bond is weaker than the C–C bond.^[27]

Isotopes

Naturally occurring silicon is composed of three stable isotopes, ²⁸Si (92.23%), ²⁹Si (4.67%), and ³⁰Si (3.10%).^[28] Out of these, only ²⁹Si is of use in NMR and EPR spectroscopy,^[29] as it is the only one with a nuclear spin (I = 1/2).^[30] All three are produced in stars through the oxygen-burning process, with ²⁸Si being made as part of the alpha process and hence the most abundant. The fusion of ²⁸Si with alpha particles by photodisintegration rearrangement in stars is known as the silicon-burning process; it is the last stage of stellar nucleosynthesis before the rapid collapse and violent explosion of the star in question in a type II supernova.^[31]

Twenty radioisotopes have been characterized, with the two stablest being ³²Si with a half-life of about 150 years, and ³¹Si with a half-life of 2.62 hours.^[28] All of the remaining radioactive isotopes have half-lives that are less than seven seconds, and the majority of these have half-lives that are less than one tenth of a second.^[28] Silicon does not have any known nuclear isomers.^{[28] 32}Si undergoes low-energy beta decay to ³²P and then stable ³²S. ³¹Si may be produced by the neutron activation of natural silicon and is thus useful for quantitative analysis; it can be easily detected by its characteristic beta decay to stable ³¹P, in which the emitted electron carries up to 1.48 MeV of energy.^[30]

The known isotopes of silicon range in mass number from 22 to 44.^[28] The most common decay mode of the isotopes with mass numbers lower than the three stable isotopes is inverse beta decay, primarily forming aluminium isotopes (13 protons) as decay products.^[28] The most common decay mode for the heavier unstable isotopes is beta decay, primarily forming phosphorus isotopes (15 protons) as decay products.^[28]

Chemistry and compounds

Crystalline bulk silicon is rather inert, but becomes more reactive at high temperatures. Like its neighbour aluminium, silicon forms a thin, continuous surface layer of silicon dioxide (SiO₂) that protects the metal from oxidation. Thus silicon does not measurably react with the air below 900 °C, but formation of the vitreous dioxide rapidly increases between 950 °C and 1160 °C and when 1400 °C is reached, atmospheric nitrogen also reacts to give the nitrides SiN and Si₃N₄. Silicon reacts with gaseous sulfur at 600 °C and gaseous phosphorus at 1000 °C. This oxide layer nevertheless does not prevent reaction with the halogens; fluorine attacks silicon vigorously at room temperature, chlorine does so at about 300 °C, and bromine and iodine at about 500 °C. Silicon does not react with most aqueous acids, but is oxidised and fluorinated by a mixture of concentrated nitric acid and hydrofluoric acid; it readily dissolves in hot aqueous alkali to form silicates. At high temperatures, silicon also reacts with alkyl halides; this reaction can be catalysed by copper to directly synthesise organosilicon chlorides as precursors to silicone polymers. Upon melting, silicon becomes extremely reactive, alloying with most metals to form silicides, and reducing most metal oxides because the heat of formation of silicon dioxide is so large. As a result, containers for liquid silicon must be made of refractory, unreactive materials such as zirconium dioxide or group 4, 5, and 6 borides.^[27]

Tetrahedral coordination is a major structural motif in silicon chemistry just as it is for carbon chemistry. However, the 3p subshell is rather more diffuse than the 2p subshell and does not hybridise as well with the 3s subshell. As a result, the chemistry of silicon and its heavier congeners shows significant differences from that of carbon,^[32] and thus octahedral coordination is also significant.^[27] For example, the electronegativity of silicon (1.90) is much less than that of carbon (2.55), because the valence electrons of silicon are further from the nucleus than those of carbon and hence experience smaller electrostatic forces of attraction from the nucleus. The poor overlap of 3p orbitals also results in a much lower tendency towards catenation (formation of Si–Si bonds) for silicon than for carbon due to the concomitant weakening of the Si–Si bond compared to the C–C bond:^[33] the average Si–Si bond energy is approximately 226 kJ/mol, compared to a value of 356 kJ/mol for the C–C bond.^[34] This results in multiply bonded silicon compounds generally being much less stable than their carbon counterparts, an example of the double bond rule. On the other hand, the presence of 3d orbitals in the valence shell of silicon suggests the possibility of hypervalence, as seen in five- and six-coordinate derivatives of silicon such as SiX⁻
₅ and SiF²⁻
₆.^[33] Lastly, because of the increasing energy gap between the valence s and p orbitals as the group is descended, the divalent state grows in importance from carbon to lead, so that a few unstable divalent compounds are known for silicon; this lowering of the main oxidation state, in tandem with increasing atomic radii, results in an increase of metallic character down the group. Silicon already shows some incipient metallic behaviour, particularly in the behaviour of its oxide compounds and its reaction with acids as well as bases (though this takes some effort), and is hence often referred to as a metalloid rather than a nonmetal.^[33] However, metallicity does not become clear in group 14 until germanium and dominant until tin, with the growing importance of the lower +2 oxidation state.^[10]

Silicon shows clear differences with carbon. For example, organic chemistry has very few analogies with silicon chemistry, while silicate minerals have a structural complexity unseen in oxocarbons.^[10] Silicon tends to resemble germanium far more than it does carbon, and this resemblance is enhanced by the d-block contraction resulting in the size of the germanium atom being much closer to that of the silicon atom than periodic trends would predict.^[26] Nevertheless, there are still some differences because of the growing importance of the divalent state in germanium compared to silicon, that result in germanium being significantly more metallic than silicon. Additionally, the lower Ge–O bond strength compared to the Si–O bond strength results in the absence of "germanone" polymers that would be analogous to silicone polymers.^[34]

Silicides

Phase diagram of the Fe–Si system

Many metal silicides are known, most of which have formulae that cannot be explained through simple appeals to valence: their bonding ranges from metallic to ionic and covalent. Some known stoichiometries are M₆Si, M₅Si, M₄Si, M₁₅Si₄, M₃Si, M₅Si₂, M₂Si, M₅Si₃, M₃Si₂, MSi, M₂Si₃, MSi₂, MSi₃, and MSi₆. They are structurally more similar to the borides than the carbides, in keeping with the diagonal relationship between boron and silicon, although the larger size of silicon than boron means that exact structural analogies are few and far between. The heats of formation of the silicides are usually similar to those of the borides and carbides of the same elements, but they usually melt at lower temperatures.^[35] Silicides are known for all stable elements in groups 1–10, with the exception of beryllium: in particular, uranium and the transition metals of groups 4–10 show the widest range of stoichiometries. Except for copper, the metals in groups 11–15 do not form silicides. Most instead form eutectic mixtures, although the heaviest post-transition metals mercury, thallium, lead, and bismuth are completely immiscible with liquid silicon.^[35]

Silicides are usually prepared by direct reaction of the elements. For example, the alkali metals and alkaline earth metals react with silicon or silicon oxide to give silicides. Nevertheless, even with these highly electropositive elements true silicon anions are not obtainable, and most of these compounds are semiconductors. For example, the alkali metal silicides (M⁺)
₄(E⁴⁻
₄) contain pyramidal tricoordinate silicon in the Si⁴⁻
₄ anion, isoeelctronic with white phosphorus, P₄.^[35][36] Metal-rich silicides tend to have isolated silicon atoms (e.g. Cu₅Si); with increasing silicon content, catenation increases, resulting in isolated clusters of two (e.g. U₃Si₂) or four silicon atoms (e.g. [K⁺]₄[Si₄]⁴⁻) at first, followed by chains (e.g. CaSi), layers (e.g. CaSi₂), or three-dimensional networks of silicon atoms spanning space (e.g. α-ThSi₂) as the silicon content rises even higher.^[35]

The silicides of the group 1 and 2 metals are usually more reactive than the transition metal silicides. The latter usually do not react with aqueous reagents, except for hydrofluoric acid; however, they do react with much more aggressive reagents like liquid potassium hydroxide, or gaseous fluorine or chlorine when red-hot. The pre-transition metal silicides instead readily react with water and aqueous acids, usually producing hydrogen or silanes:^[35]

Na₂Si + 3 H₂O → Na₂SiO₃ + 3 H₂

Mg₂Si + 2 H₂SO₄ → 2 MgSO₄ + SiH₄

Products often vary with the stoichiometry of the silicide reactant. For example, Ca₂Si is polar and non-conducting and has the anti-PbCl₂ structure with single isolated silicon atoms, and reacts with water to produce calcium hydroxide, hydrated silicon dioxide, and hydrogen gas. CaSi with its zigzag chains of silicon atoms instead reacts to give silanes and polymeric SiH₂, while CaSi₂ with its puckered layers of silicon atoms does not react with water, but will react with dilute hydrochloric acid: the product is a yellow polymeric solid with stoichiometry Si₂H₂O.^[35]

Silanes

Speculation on silicon hydride chemistry started from the 1830s, contemporary with the development of synthetic organic chemistry. Silane itself, as well as trichlorosilane, were first synthesised by Friedrich Wöhler and Heinrich Buff in 1857 by reacting aluminium–silicon alloys with hydrochloric acid, and characterised as SiH₄ and SiHCl₃ by Charles Friedel and Albert Ladenburg in 1867.  Disilane (Si₂H₆) followed in 1902, when it was first made by Henri Moissan and Samuel Smiles by the protonolysis of magnesium silicides. Further investigation had to wait until 1916 because of the great reactivity and thermal instability of the silanes; it was then that Alfred Stock began the study of silicon hydrides in earnest with new greaseless vacuum techniques, as they were found as contaminants of his focus, the boron hydrides. The names silanes and boranes are due to him, based on analogy with the alkanes.^[23][37][38] Moissan and Smiles' method of preparation of silanes and silane derivatives via protonolysis of metal silicides is still used, although the yield is lowered by the hydrolysis of the products that occurs simultaneously, and thus the preferred route today is to treat substituted silanes with hydride reducing agents such as lithium aluminium hydride in etheric solutions at low temperatures. Direct reaction of HX or RX with silicon, possibly with a catalyst such as copper, is also a viable method to produce substituted silanes.^[23]

The silanes comprise a homologous series of silicon hydrides with the general formula Si_nH_2n + 2. They are all strong reducing agents. Unbranched and branched trains are known up to n = 8, and the cycles Si₅H₁₀ and Si₆H₁₂ are also known. The first two, silane and disilane, are colourless gases; the heavier members of the series are volatile liquids. All silanes are very reactive and catch fire or explode spontaneously in air. They become less thermally stable with room temperature, so that only silane is indefinitely stable at room temperature, although disilane does not decompose very quickly (only 2.5% of a sample decomposes after eight months have passed).^[23] They decompose to form polymeric polysilicon hydride and hydrogen gas.^[39][40] As expected from the difference in atomic weight, the silanes are less volatile than the corresponding alkanes and boranes, but more volatile than the corresponding germanes. They are much more reactive than the corresponding alkanes, because the larger radius of silicon compared to carbon facilitates nucleophilic attack at the silicon, the greater polarity of the Si–H bond compared to the C–H bond, and the ability of silicon to expand its octet and hence form adducts and lower the reaction's activation energy.^[23]

Silane pyrolysis gives polymeric species and finally elemental silicon and hydrogen; indeed ultrapure silicon is commercially produced by the pyrolysis of silane. While the thermal decomposition of alkanes starts by the breaking of a C–H or C–C bond and the formation of radical intermediates, polysilanes decompose by eliminating silenes :SiH₂ or :SiHR, as the activation energy of this process (~210 kJ/mol) is much less than the Si–Si and Si–H bond energies. While pure silanes do not react with pure water or dilute acids, traces of alkali catalyse immediate hydrolysis to hydrated silicon dioxide. If the reaction is carried out in methanol, controlled solvolysis results in the products SiH₂(OMe)₂, SiH(OMe)₃, and Si(OMe)₄. The Si–H bond also adds to alkenes, a reaction which proceeds slowly and speeds up with increasing substitution of the silane involved. At 450 °C, silane participates in an addition reaction with acetone, as well as a ring-opening reaction with ethylene oxide. Direct reaction of the silanes with chlorine or bromine results in explosions at room temperature, but the reaction of silane with bromine at −80 °C is controlled and yields bromosilane and dibromosilane. The monohalosilanes may be formed by reacting silane with the appropriate hydrogen halide with an Al₂X₆ catalyst, or by reacting silane with a solid silver halide in a heated flow reactor:^[23]

SiH₄ + 2 AgCl 260 °C→  SiH₃Cl + HCl + 2 Ag

Among the derivatives of silane, iodosilane (SiH₃I) and potassium silanide (KSiH₃) are very useful synthetic intermediates in the production of more complicated silicon-containing compounds: the latter is a colourless crystalline ionic solid containing K⁺ cations and SiH⁻
₃ anions in the NaCl structure, and is made by the reduction of silane by potassium metal.^[41] Additionally, the reactive hypervalent species SiH⁻
₅ is also known.^[23] With suitable organic substituents it is possible to produce stable polysilanes: they have surprisingly high electric conductivities, arising from sigma delocalisation of the electrons in the chain.^[42]

Halides

Silicon and silicon carbide readily react with all four stable halogens, forming the colourless, reactive and volatile silicon tetrahalides.^[43] Silicon tetrafluoride may also be made by fluorinating the other silicon halides, and is produced by the attack of hydrofluoric acid on glass.^[44] Heating two different tetrahalides together also produce a random mixture of mixed halides, which may also be produced by halogen exchange reactions. The melting and boiling points of these species usually rise with increasing atomic weight, though there are many exceptions: for example, the melting and boiling points drop as one passes from SiFBr₃ through SiFClBr₂ to SiFCl₂Br. The shift from the hypoelectronic elements in group 13 and earlier to the group 14 elements is illustrated by the change from an infinite ionic structure in aluminium fluoride to a lattice of simple covalent silicon tetrafluoride molecules, as dictated by the lower electronegativity of aluminium than silicon, the stoichiometry (the +4 oxidation state being too high for true ionicity), and the smaller size of the silicon atom compared to the aluminium atom.^[43] Silicon tetrachloride is manufactured on a huge scale as a precursor to the production of pure silicon, silicon dioxide, and some silicon esters.^[43] The silicon tetrahalides hydrolyse readily in water, unlike the carbon tetrahalides, again because of the larger size of the silicon atom rendering it more open to nucleophilic attack and the ability of the silicon atom to expand its octet which carbon lacks.^[44] The reaction of silicon fluoride with excess hydrofluoric acid produces the octahedral hexafluorosilicate anion SiF²⁻
₆.^[44]

Analogous to the silanes, halopolysilanes Si_nX_2n + 2 are also known. While catenation in carbon compounds is maximised in the hydrogen compounds rather than the halides, the opposite is true for silicon, so that the halopolysilanes are known up to at least Si₁₄F₃₀, Si₆Cl₁₄, and Si₄Br₁₀. A suggested explanation for this phenomenon is the compensation for the electron loss of silicon to the more electronegative halogen atoms by pi backbonding from the filled p_π orbitals on the halogen atoms to the empty d_π orbitals on silicon: this is similar to the situation of carbon monoxide in metal carbonyl complexes and explains their stability. These halopolysilanes may be produced by comproportionation of silicon tetrahalides with elemental silicon, or by condensation of lighter halopolysilanes (trimethylammonium being a useful catalyst for this reaction).^[43]

Silica

Silicon dioxide (SiO₂), also known as silica, is one of the most well-studied compounds, second only to water. Twelve different crystal modifications of silica are known, the most common being α-quartz, a major constituent of many rocks such as granite and sandstone. It is also known to occur pure as rock crystal; impure forms are known as rose quartz, smoky quartz, morion, amethyst, and citrine. Some poorly crystalline forms of quartz are also known, such as chalcedony, chrysoprase, carnelian, agate, onyx, jasper, heliotrope, and flint. Other modifications of silicon dioxide are known in some other minerals such as tridymite and cristobalite, as well as the much less common coesite and stishovite. Biologically generated forms are also known as kieselguhr and diatomaceous earth.  Vitreous silicon dioxide is known as tektites, and obsidian, and rarely as lechatelierite. Some synthetic forms are known as keatite and W-silica. Opals are composed of complicated crystalline aggregates of partially hydrated silicon dioxide.^[45]

• 

  Quartz
• 

  Sandstone
• 

  Smoky quartz
• 

  Amethyst
• 

  Citrine
• 

  Agate
• 

  Chrysoprase
• 

  Onyx
• 

  Jasper
• 

  Heliotrope
• 

  Tridymite
• 

  Cristobalite
• 

  Coesite
• 

  Lechatelierite

Most crystalline forms of silica are made of infinite arrangements of {SiO₄} tetrahedra (with Si at the centre) connected at their corners, with each oxygen atom linked to two silicon atoms. In the thermodynamically stable room-temperature form, α-quartz, these tetrahedra are linked in intertwined helical chains with two different Si–O distances (159.7 and 161.7 pm) with a Si–O–Si angle of 144°. These helices can be either left- or right-handed, so that individual α-quartz crystals are optically active. At 537 °C, this transforms quickly and reversibly into the similar β-quartz, with a change of the Si–O–Si angle to 155° but a retention of handedness. Further heating to 867 °C results in another reversible phase transition to β-tridymite, in which some Si–O bonds are broken to allow for the arrangement of the {SiO₄} tetrahedra into a more open and less dense hexagonal structure. This transition is slow and hence tridymite occurs as a metastable mineral even below this transition temperature; when cooled to about 120 °C it quickly and reversibly transforms by slight displacements of individual silicon and oxygen atoms to α-tridymite, similarly to the transition from α-quartz to β-quartz. β-tridymite slowly transforms to cubic β-cristobalite at about 1470 °C, which once again exists metastably below this transition temperature and transforms at 200–280 °C to α-cristobalite via small atomic displacements. β-cristobalite melts at 1713 °C; the freezing of silica from the melt is quite slow and vitrification, or the formation of a glass, is likely to occur instead. In vitreous silica, the {SiO₄} tetrahedra remain corner-connected, but the symmetry and periodicity of the crystalline forms are lost. Because of the slow conversions between these three forms, it is possible upon rapid heating to melt β-quartz (1550 °C) or β-tridymite (1703 °C). Silica boils at approximately 2800 °C. Other high-pressure forms of silica are known, such as coesite and stishovite: these are known in nature, formed under the shock pressure of a meteorite impact and then rapidly quenched to preserve the crystal structure. Similar melting and cooling of silica occurs following lightning strikes, forming glassy lechatelierite. W-silica is an unstable low-density form involving {SiO₄} tetrahedra sharing opposite edges instead of corners, forming parallel chains similarly to silicon disulfide (SiS₂) and silicon diselenide (SiSe₂): it quickly returns to forming amorphous silica with heat or traces of water.^[45]

Condensed polysilicic acid

Silica is rather inert chemically. It is not attacked by any acids other than hydrofluoric acid. However, it slowly dissolves in hot concentrated alkalis, and does so rather quickly in fused metal hydroxides or carbonates to give metal silicates. Among the elements, it is attacked only by fluorine at room temperature to form silicon tetrafluoride: hydrogen and carbon also react, but require temperatures over 1000 °C to do so. Silica nevertheless reacts with many metal and metalloid oxides to form a wide variety of compounds important in the glass and ceramic industries above all, but also have many other uses: for example, sodium silicate is often used in detergents due to its buffering, saponifying, and emulsifying properties.^[45]

Silicic acids

Adding water to silica drops its melting point by around 800 °C due to the breaking of the structure by replacing Si–O–Si linkages with terminating Si–OH groups. Increasing water concentration results in the formation of hydrated silica gels and colloidal silica dispersions. Many hydrates and silicic acids exist in the most dilute of aqueous solutions, but these are rather insoluble and quickly precipitate and condense and cross-link to form various polysilicic acids of variable combinations following the formula [SiO_x(OH)_4−2x]_n, similar to the behaviour of boron, aluminium, and iron, among other elements. Hence, although some simple silicic acids have been identified in dilute solutions, such as orthosilicic acid Si(OH)₄ and metasilicic acid SiO(OH)₂, none of these are likely to exist in the solid state.^[45]

Silicate minerals

About 95% of the Earth's crustal rocks are made of silica or silicate and aluminosilicate minerals, as reflected in oxygen, silicon, and aluminium being the three most common elements in the crust (in that order).^[46] Measured by mass, silicon makes up 27.7% of the Earth's crust.^[47] Pure silicon crystals are very rarely found in nature, but notable exceptions are crystals as large as to 0.3 mm across found during sampling gases from the Kudriavy volcano on the island of Iturup, one of the Kuril Islands.^[48][49]

Silicate and aluminosilicate minerals have many different structures and varying stoichiometry, but they may be classified following some general principles. Tetrahedral {SiO₄} units are common to almost all these compounds, either as discrete structures, or combined into larger units by the sharing of corner oxygen atoms. These may be divided into neso-silicates (discrete {SiO₄} units) sharing no oxygen atoms, soro-silicates (discrete {Si₂O₇} units) sharing one, cyclo-silicates (closed ring structures) and ino-silicates (continuous chain or ribbon structures) both sharing two, phyllo-silicates (continuous sheets) sharing three, and tecto-silicates (continuous three-dimensional frameworks) sharing four. The lattice of oxygen atoms that results is usually close-packed or close to it, with the charge being balanced by other cations in various different polyhedral sites according to size.^[46]

The orthosilicates M^II
₂SiO
₄ (M = Be, Mg, Mn, Fe, Zn) and ZrSiO₄ are neso-silicates. Be₂SiO₄ (phenacite) is rather unusual as both Be^II and Si^IV occupy tetrahedral four-coordinated sites; the other divalent cations instead occupy six-coordinated octahedral sites and often isomorphously replace each other as in olivine, (Mg,Fe,Mn)₂SiO₄. Zircon, ZrSiO₄, demands eight-coordination of the Zr^IV cations due to stoichiometry and because of their larger ionic radius (84 pm). Also significant are the garnets, [M^II
₃M^III
₂(SiO
₄)
₃], in which the divalent cations (e.g. Ca, Mg, Fe) are eight-coordinated and the trivalent ones are six-coordinated (e.g. Al, Cr, Fe). Regular coordination is not always present: for example, it is not found in Ca₂SiO₄, which mixes six- and eight-coordinate sites for Ca^II. Soro-silicates, involving discrete double or triple tetrahedral units, are quite rare: metasilicates involving cyclic "[(SiO₃)_n]²ⁿ⁻" units of corner-abutting tetrahedra forming a polygonal ring are also known.^[46]

Chain metasilicates, {SiO²⁻
₃}
_∞, form by the corner-sharing of an indefinite chain of linked {SiO₄} tetrahedra. Many differences arise due to the differing repeat distances of conformation across the line of tetrahedra. A repeat distance of two is most common, as in most pyroxene minerals, but repeat distances of one, three, four, five, six, seven, nine, and twelve are also known. These chains can then link across each other to form double chains and ribbons, as in the asbestos minerals, involving repeated chains of cyclic tetrahedron rings.^[46]

A typical zeolite structure

Layer silicates, such as the clay minerals and the micas, are very common, and are often formed by horizontal cross-linking of metasilicate chains or planar condensation of smaller units. An example is kaolinite [Al₂(OH)₄Si₂O₅]; in many of these minerals cation and anion replacement is common, so that for example tetrahedral Si^IV may be replaced by Al^III, octahedral Al^III by Mg^II, and OH⁻ by F⁻. Three-dimensional framework aluminosilicates are structurally very complex; they may be conceived of as starting from the SiO₂ structure, but having replaced up to one-half of the Si^IV atoms with Al^III they require more cations to be included in the structure to balance charge. Examples include feldspars (the most abundant minerals on the Earth), zeolites, and ultramarines. Many feldspars can be thought of as forming part of the ternary system NaAlSi₃O₈–KAlSi₃O₈–CaAl₂Si₂O₈. Their lattice is destroyed by high pressure prompting Al^III to undergo six-coordination rather than four-coordination, and this reaction destroying feldspars may be a reason for the Mohorovičić discontinuity, which would imply that the crust and mantle have the same chemical composition but different lattices, although this is not a universally held view. Zeolites have many polyhedral cavities in their frameworks (truncated cuboctahedra being most common, but other polyhedra are also known as zeolite cavities), allowing them to include loosely bound molecules such as water in their structure. Ultramarines alternate silicon and aluminium atoms and include a variety of other anions such as Cl⁻, SO²⁻
₄, and S²⁻
₂, but are otherwise similar to the feldspars.^[46]

Other inorganic compounds

Silicon disulfide (SiS₂) is formed by burning silicon in gaseous sulfur at 100 °C; sublimation of the resulting compound in nitrogen results in white, flexible long fibres reminiscent of asbestos with a structure similar to W-silica. This melts at 1090 °C and sublimes at 1250 °C; at high temperature and pressure this transforms to a crystal structure analogous to cristobalite. However, SiS₂ lacks the variety of structures of SiO₂, and quickly hydrolyses to silica and hydrogen sulfide. It is also ammonoloysed quickly and completely by liquid ammonia as follows to form an imide:^[50]

SiS₂ + 4 NH₃ → Si(NH)₂ + 2 NH₄SH

It reacts with the sulfides of sodium, magnesium, aluminium, and iron to form metal thiosilicates: reaction with ethanol results in ethylsilicate Si(OEt)₄ and hydrogen sulfide. Ethylsilicate is useful as its controlled hydrolysis produces adhesive or film-like forms of silica. Reacting hydrogen sulfide with silicon tetrahalides yields silicon thiohalides such as S(SiCl)₃, cyclic Cl₂Si(μ-S)₂SiCl₂, and crystalline (SiSCl₂)₄. Despite the double bond rule, stable organosilanethiones RR'Si=S have been made thanks to the stabilising mechanism of intermolecular coordination via an amine group.^[50]

Silicon nitride, Si₃N₄, can be formed by directly reacting silicon with nitrogen above 1300 °C, but a more economical means of production is by heating silica and coke in a stream of nitrogen and hydrogen gas at 1500 °C. It would make a promising ceramic if not for the difficulty of working with and sintering it: it is chemically near-totally inert, and even above 1000 °C it keeps its strength, shape, and continues to be resistant to wear and corrosion. It is very hard (9 on the Mohs hardness scale), dissociates only at 1900 °C at 1 atm, and is quite dense (density 3.185 g/cm³), because of its compact structure similar to that of phenacite (Be₂SiO₄). A similar refractory material is Si₂N₂O, formed by heating silicon and silica at 1450 °C in an argon stream containing 5% nitrogen gas, involving 4-coordinate silicon and 3-coordinate nitrogen alternating in puckered hexagonal tilings interlinked by non-linear Si–O–Si linkages to each other.^[50]

Reacting silyl halides with ammonia or alkylammonia derivatives in the gaseous phase or in ethanolic solution produces various volatile silylamides, which are silicon analogues of the amines:^[50]

3 SiH₃Cl + 4 NH₃ → N(SiH₃)₃ + 3 NH₄Cl

SiH₃Br + 2 Me₂NH → SiH₃NMe₂ + Me₂NH₂Br

4 SiH₃I + 5 N₂H₄ → (SiH₃)₂NN(SiH₃)₂ + 4 N₂H₅I

Many such compounds have been prepared, the only known restriction being that the nitrogen is always tertiary, and species containing the SiH–NH group are unstable at room temperature. The stoichiometry around the nitrogen atom in compounds such as N(SiH₃)₃is planar, which has been attributed to a p_π–d_π interaction between a lone pair on nitrogen and an empty d_π orbital on silicon. Similarly, trisilylamines are weaker as ligands than their carbon analogues, the tertiary amines, although substitution of some SiH₃ groups by CH₃ groups mitigates this weakness. Thus, for example, N(SiH₃)₃ does not form an adduct with BH₃ at all, while MeN(SiH₃)₂ and Me₂NSiH₃ form adducts at low temperatures that decompose upon warming. Some silicon analogues of imines, with a Si=N double bond, are known: the first found was Bu^t₂Si=N–SiBu^t₃, which was discovered in 1986.^[50]

Silicon carbide

Silicon carbide (SiC) was first made by Edward Goodrich Acheson in 1891, who named it carborundum to reference its intermediate hardness and abrasive power between diamond (an allotrope of carbon) and corundum (aluminium oxide). He soon founded a company to manufacture it, and today about one million tonnes are produced each year.^[51] Silicon carbide exists in about 250 crystalline forms.^[52] The polymorphism of SiC is characterized by a large family of similar crystalline structures called polytypes. They are variations of the same chemical compound that are identical in two dimensions and differ in the third. Thus, they can be viewed as layers stacked in a certain sequence.^[53] It is made industrially by reduction of quartz sand with excess coke or anthracite at 2000–2500 °C in an electric furnace:^[51]

SiO₂ + 2 C → Si + 2 CO

Si + C → SiC

It is the most thermally stable binary silicon compound, only decomposing through loss of silicon starting from around 2700 °C. It is resistant to most aqueous acids, phosphoric acid being an exception. It forms a protective layer of silicon dioxide on the surface and hence only oxidises appreciably in air above 1000 °C; removal of this layer by molten hydroxides or carbonates leads to quick oxidation. Silicon carbide is rapidly attacked by chlorine gas, which forms SiCl₄ and carbon at 100 °C and SiCl₄ and CCl₄ at 1000 °C. It is mostly used as an abrasive and a refractory materia, as it is chemically stable and very strong, and it fractures to form a very sharp cutting edge. It is also useful as an intrinsic semiconductor, as well as an extrinsic semiconductor upon being doped.^[51] In its diamond-like behaviour it serves as an illustration of the chemical similarity between carbon and silicon.^[54]

Organosilicon compounds

A hydrosilylation reaction, in which Si–H is added to an unsaturated substrate

Because the Si–C bond is close in strength to the C–C bond, organosilicon compounds tend to be markedly thermally and chemically stable. For example, tetraphenylsilane (SiPh₄) may be distilled in air even at its boiling point of 428 °C, and so can its substituted derivatives Ph₃SiCl and Ph₂SiCl₂, which boil at 378 °C and 305 °C respectively. Furthermore, since carbon and silicon are chemical congeners, organosilicon chemistry shows some significant similarities with carbon chemistry, for example in the propensity of such compounds for catenation and forming multiple bonds.^[54] However, significant differences also arise: since silicon is more electropositive than carbon, bonds to more electronegative elements are generally stronger with silicon than with carbon, and vice versa. Thus the Si–F bond is significantly stronger than even the C–F bond and is one of the strongest single bonds, while the Si–H bond is much weaker than the C–H bond and is readily broken. Furthermore, the ability of silicon to expand its octet is not shared by carbon, and hence some organosilicon reactions have no organic analogues. For example, nucleophilic attack on silicon does not proceed by the S_N2 or S_N1 processes, but instead goes through a negatively charged true pentacoordinate intermediate and appears like a substitution at a hindered tertiary atom. This works for silicon, unlike for carbon, because the long Si–C bonds reduce the steric hindrance and the d-orbital of silicon is geometrically unconstrained for nucleophilic attack, unlike for example a C–O σ* antibonding orbital. Nevertheless, despite these differences, the mechanism is still often called "S_N2 at silicon" for simplicity.^[55]

One of the most useful silicon-containing groups is trimethylsilyl, Me₃Si–. The Si–C bond connecting it to the rest of the molecule is reasonably strong, allowing it to remain while the rest of the molecule undergoes reactions, but is not so strong that it cannot be removed specifically when needed, for example by the fluoride ion, which is a very weak nucleophile for carbon compounds but a very strong one for organosilicon compounds. It may be compared to acidic protons; while trisilylmethyl is removed by hard nucleophiles instead of bases, both removals usually promote elimination. As a general rule, while saturated carbon is best attacked by nucleophiles that are neutral compounds, those based on nonmetals far down on the periodic table (e.g. sulfur, selenium, or iodine), or even both, silicon is best attacked by charged nucleophiles, particularly those involving such highly electronegative nonmetals as oxygen, fluorine, or chlorine. For example, enolates react at the carbon in haloalkanes, but at the oxygen in silyl chlorides; and when trimethylsilyl is removed from an organic molecule using hydroxide as a nucleophile, the product of the reaction is not the silanol as one would expect from using carbon chemistry as an analogy, because the siloxide is strongly nucleophilic and attacks the original molecule to yield the silyl ether hexamethyldisiloxane, (Me₃Si)₂O. Conversely, while the S_N2 reaction is mostly unaffected by the presence of a partial positive charge (δ+) at the carbon, the analogous "S_N2" reaction at silicon is so affected. Thus, for example, the silyl triflates are so electrophilic that they react 10⁸ to 10⁹ times faster than silyl chlorides with oxygen-containing nucleophiles. Trimethylsilyl triflate is in particular a very good Lewis acid and is used to convert carbonyl compounds to acetals and silyl enol ethers, reacting them together analogously to the aldol reaction.^[55]

Si–C bonds are commonly formed in three ways. In the laboratory, preparation is often carried out in small quantities by reacting tetrachlorosilane with organolithium, Grignard, or organoaluminium reagents, or by catalytic addition of Si–H across C=C double bonds. The second route has the drawback of not being applicable to the most important silanes, the methyl and phenyl silanes. Organosilanes are made industrially by directly reacting alkyl or aryl halides with silicon with 10% by weight metallic copper as a catalyst. Standard organic reactions suffice to produce many derivatives; the resulting organosilanes are often significantly more reactive than their carbon congeners, readily undergoing hydrolysis, ammonolysis, alcoholysis, and condensation to form cyclic oligomers or linear polymers.^[54]

Silicone polymers

Structure of polydimethylsiloxane, the principal component of silicones

The word "silicone" was first used by Frederick Kipping in 1901. He invented the word to illustrate the similarity of chemical formulae between Ph₂SiO and benzophenone, Ph₂CO, although he also stressed the lack of chemical resemblance due to the polymeric structure of Ph₂SiO, which is not shared by Ph₂CO.^[54]

Silicones may be considered analogous to mineral silicates, in which the methyl groups of the silicones correspond to the isoelectronic O⁻ of the silicates.^[54] They are quite stable to extreme temperatures, oxidation, and water, and have useful dielectric, antistick, and antifoam properties. Furthermore, they are resistant over long periods of time to ultraviolet radiation and weathering, and are inert physiologically. They are fairly unreactive, but do react with concentrated solutions bearing the hydroxide ion and fluorinating agents, and occasionally can be even used as mild reagents for selective syntheses. For example, (Me₃Si)₂O is valuable for the preparation of derivatives of molybdenum and tungsten oxyhalides, converting a tungsten hexachloride suspension in dichloroethane solution quantitatively to WOCl₄ in under an hour at room temperature, and then to yellow WO₂Cl₂ at 100 °C in light petroleum at a yield of 95% overnight.^[54]

Occurrence

Olivine

In the universe, silicon is the seventh most abundant element, coming after hydrogen, helium, carbon, nitrogen, oxygen, and neon. These abundances are not replicated well on Earth due to substantial separation of the elements taking place during the formation of the Solar System. Silicon makes up 27.2% of the Earth's crust by weight, second only to oxygen at 45.5%, with which it is always associated in nature. Further fractionation took place in the formation of the Earth by planetary differentiation: Earth's core, which makes up 31.5% of the mass of the Earth, has approximate composition Fe₂₅Ni₂Co_0.1S₃; the mantle makes up 68.1% of the Earth's mass and is composed mostly of denser oxides and silicates, an example being olivine, (Mg,Fe)₂SiO₄; while the lighter siliceous minerals such as aluminosilicates rise to the surface and form the crust, making up 0.4% of the Earth's mass.^[56]

The crystallisation of igneous rocks from magma depends on a number of factors; among them are the chemical composition of the magma, the cooling rate, and some properties of the individual minerals to be formed, such as lattice energy, melting point, and complexity of their crystal structure. As magma is cooled, olivine appears first, followed by pyroxene, amphibole, biotite mica, orthoclase feldspar, muscovite mica, quartz, zeolites, and finally hydrothermal minerals. This sequence shows a trend towards increasingly complex silicate units with cooling, and the introduction of hydroxide and fluoride anions in addition to oxides. Many metals can substitute for silicon. After these igneous rocks undergo weathering, transport, and deposition, sedimentary rocks like clay, shale, and sandstone are formed. Metamorphism also can occur at high temperatures and pressures, creating an even vaster variety of minerals.^[56]

Production

Silicon of 96–99% purity is made by reducing quartzite or sand with highly pure coke. The reduction is carried out in an electric arc furnace, with an excess of SiO₂ used to stop silicon carbide (SiC) from accumulating:^[30]

SiO₂ + 2 C → Si + 2 CO

2 SiC + SiO₂ → 3 Si + 2 CO

Ferrosilicon alloy

This reaction, known as carbothermal reduction of silicon dioxide, is usually conducted in the presence of scrap iron with low amounts of phosphorus and sulfur, produing ferrosilicon.^[30] Ferrosilicon, an iron-silicon alloy that contains varying ratios of elemental silicon and iron, accounts for about 80% of the world's production of elemental silicon, with China, the leading supplier of elemental silicon, providing 4.6 million tonnes (or 2/3 of the world output) of silicon, most of which is in the form of ferrosilicon. It is followed by Russia (610,000 t), Norway (330,000 t), Brazil (240,000 t) and the United States (170,000 t).^[57] Ferrosilicon is primarily used by the iron and steel industry (see below) with primary use as alloying addition in iron or steel and for de-oxidation of steel in integrated steel plants.^[30] Another sometimes used reaction is aluminothermal reduction of silicon dioxide, as follows:^[58]

3 SiO₂ + 4 Al → 3 Si + 2 Al₂O₃

Leaching powdered 96–97% pure silicon with water results in ~98.5% pure silicon, which is used in the chemical industry. However, even greater purity is needed for semiconductor applications, and this is produced from the reduction of tetrachlorosilane or trichlorosilane. The former is made by chlorinating scrap silicon and the latter is a byproduct of silicone production. These compounds are volatile and hence can be purified by repeated fractional distillation, followed by reduction to elemental silicon with very pure zinc metal as the reducing agent. The spongy pieces of silicon thus produced are melted and then grown to form cylindrical single crystals, before being purified by zone refining. Other routes use the thermal decomposition of silane or tetraiodosilane. Another process used is the reduction of sodium hexafluorosilicate, a common waste product of the phosphate fertiliser industry, by metallic sodium: this is highly exothermic and hence requires no outside fuel source. Hyperfine silicon is made at a higher purity than almost every other material: transistor production requires impurity levels in silicon crystals less than 1 part per 10¹⁰, and in special cases impurity levels below 1 part per 10¹² are needed and attained.^[30]

Applications

Compounds

Most silicon is used industrially without being purified, and indeed often with comparatively little processing from its natural form. Over 90% of the Earth's crust is composed of silicate minerals, which are compounds of silicon and oxygen, often with metallic ions when negatively charged silicate anions require cations to balance the charge. Many of these have direct commercial uses, such as clays, silica sand and most kinds of building stone. Thus, the vast majority of uses for silicon are as structural compounds, either as the silicate minerals or silica (crude silicon dioxide). Silicates are used in making Portland cement (made mostly of calcium silicates) which is used in building mortar and modern stucco, but more importantly, combined with silica sand, and gravel (usually containing silicate minerals like granite), to make the concrete that is the basis of most of the very largest industrial building projects of the modern world.^[59]

Silica is used to make fire brick, a type of ceramic. Silicate minerals are also in whiteware ceramics, an important class of products usually containing various types of fired clay minerals (natural aluminium phyllosilicates). An example is porcelain which is based on the silicate mineral kaolinite. Traditional glass (silica-based soda-lime glass) also functions in many of the same ways, and is also used for windows and containers. In addition, specialty silica based glass fibers are used for optical fiber, as well as to produce fiberglass for structural support and glass wool for thermal insulation.

Silicones are often used in waterproofing treatments, molding compounds, mold-release agents, mechanical seals, high temperature greases and waxes, and caulking compounds. Silicone is also sometimes used in breast implants, contact lenses, explosives and pyrotechnics.^[60] Silly Putty was originally made by adding boric acid to silicone oil.^[61] Other silicon compounds function as high-technology abrasives and new high-strength ceramics based upon silicon carbide. Silicon is a component of some superalloys.

Alloys

Elemental silicon is added to molten cast iron as ferrosilicon or silicocalcium alloys to improve performance in casting thin sections and to prevent the formation of cementite where exposed to outside air. The presence of elemental silicon in molten iron acts as a sink for oxygen, so that the steel carbon content, which must be kept within narrow limits for each type of steel, can be more closely controlled. Ferrosilicon production and use is a monitor of the steel industry, and although this form of elemental silicon is grossly impure, it accounts for 80% of the world's use of free silicon. Silicon is an important constituent of electrical steel, modifying its resistivity and ferromagnetic properties.

The properties of silicon can be used to modify alloys with metals other than iron. "Metallurgical grade" silicon is silicon of 95–99% purity. About 55% of the world consumption of metallurgical purity silicon goes for production of aluminium-silicon alloys (silumin alloys) for aluminium part casts, mainly for use in the automotive industry. Silicon's importance in aluminium casting is that a significantly high amount (12%) of silicon in aluminium forms a eutectic mixture which solidifies with very little thermal contraction. This greatly reduces tearing and cracks formed from stress as casting alloys cool to solidity. Silicon also significantly improves the hardness and thus wear-resistance of aluminium.^[62][63]

Electronics

Silicon wafer with mirror finish

Most elemental silicon produced remains as a ferrosilicon alloy, and only about 20% is refined to metallurgical grade purity (a total of 1.3–1.5 million metric tons/year). An estimated 15% of the world production of metallurgical grade silicon is further refined to semiconductor purity.^[63] This typically is the "nine-9" or 99.9999999% purity ^[64] nearly defect-free single crystalline material.^[65]

Monocrystalline silicon of such purity is usually produced by the Czochralski process, is used to produce silicon wafers used in the semiconductor industry, in electronics, and in some high-cost and high-efficiency photovoltaic applications.^[66] Pure silicon is an intrinsic semiconductor, which means that unlike metals, it conducts electron holes and electrons released from atoms by heat; silicon's electrical conductivity increases with higher temperatures. Pure silicon has too low a conductivity (i.e., too high a resistivity) to be used as a circuit element in electronics. In practice, pure silicon is doped with small concentrations of certain other elements, which greatly increase its conductivity and adjust its electrical response by controlling the number and charge (positive or negative) of activated carriers. Such control is necessary for transistors, solar cells, semiconductor detectors, and other semiconductor devices used in the computer industry and other technical applications.^[67] In silicon photonics, silicon can be used as a continuous wave Raman laser medium to produce coherent light.^[68]

In common integrated circuits, a wafer of monocrystalline silicon serves as a mechanical support for the circuits, which are created by doping and insulated from each other by thin layers of silicon oxide, an insulator that is easily produced by exposing the element to oxygen under the proper conditions. Silicon has become the most popular material for both high power semiconductors and integrated circuits because it can withstand the highest temperatures and greatest electrical activity without suffering avalanche breakdown (an electron avalanche is created when heat produces free electrons and holes, which in turn pass more current, which produces more heat). In addition, the insulating oxide of silicon is not soluble in water, which gives it an advantage over germanium (an element with similar properties which can also be used in semiconductor devices) in certain fabrication techniques.^[69]

Monocrystalline silicon is expensive to produce, and is usually justified only in production of integrated circuits, where tiny crystal imperfections can interfere with tiny circuit paths. For other uses, other types of pure silicon may be employed. These include hydrogenated amorphous silicon and upgraded metallurgical-grade silicon (UMG-Si) used in the production of low-cost, large-area electronics in applications such as liquid crystal displays and of large-area, low-cost, thin-film solar cells. Such semiconductor grades of silicon are either slightly less pure or polycrystalline rather than monocrystalline, and are produced in comparable quatities as the monocrystalline silicon: 75,000 to 150,000 metric tons per year. The market for the lesser grade is growing more quickly than for monocrystalline silicon. By 2013, polycrystalline silicon production, used mostly in solar cells, was projected to reach 200,000 metric tons per year, while monocrystalline semiconductor grade silicon was expected to remain less than 50,000 tons/year.^[63]

Biological role

A diatom, enclosed in a silica cell wall

Although silicon is readily available in the form of silicates, very few organisms use it directly. Diatoms, radiolaria and siliceous sponges use biogenic silica as a structural material for skeletons. In more advanced plants, the silica phytoliths (opal phytoliths) are rigid microscopic bodies occurring in the cell; some plants, for example rice, need silicon for their growth.^[70][71][72] There is some evidence that silicon is important to nail, hair, bone and skin health in humans,^[73] for example in studies that show that premenopausal women with higher dietary silicon intake have higher bone density, and that silicon supplementation can increase bone volume and density in patients with osteoporosis.^[74] Silicon is needed for synthesis of elastin and collagen, of which the aorta contains the greatest quantity in the human body^[75] and has been considered an essential element;^[76] nevertheless, it is difficult to prove its essentiality, because silicon is very common and hence deficiency symptoms are difficult to reproduce.^[77]

Silicon is currently under consideration for elevation to the status of a "plant beneficial substance by the Association of American Plant Food Control Officials (AAPFCO)."^[78][79] Silicon has been shown to improve plant cell wall strength and structural integrity in some plants.^[80]

Safety

People can be exposed to elemental silicon in the workplace by breathing it in, swallowing it, skin contact, and eye contact. In the latter two cases, silicon poses a slight hazard as an irritant; it is hazardous if inhaled.^[81] The Occupational Safety and Health Administration (OSHA) has set the legal limit (Permissible exposure limit) for silicon exposure in the workplace as 15 mg/m³ total exposure and 5 mg/m³ respiratory exposure over an 8-hour workday. The National Institute for Occupational Safety and Health (NIOSH) has set a Recommended exposure limit (REL) of 10 mg/m³ total exposure and 5 mg/m³ respiratory exposure over an 8-hour workday.^[82] Inhalation of crystalline silica dust may lead to silicosis, an occupational lung disease marked by inflammation and scarring in the form of nodular lesions in the upper lobes of the lungs.^[83]

Nanonetwork

From Wikipedia, the free encyclopedia

A nanonetwork or nanoscale network is a set of interconnected nanomachines (devices a few hundred nanometers or a few micrometers at most in size), which are able to perform only very simple tasks such as computing, data storing, sensing and actuation.^[1][2] Nanonetworks are expected to expand the capabilities of single nanomachines both in terms of complexity and range of operation by allowing them to coordinate, share and fuse information. Nanonetworks enable new applications of nanotechnology in the biomedical field, environmental research, military technology and industrial and consumer goods applications. Nanoscale communication is defined in IEEE P1906.1.

Communication approaches

Classical communication paradigms need to be revised for the nanoscale. The two main alternatives for communication in the nanoscale are based either on electromagnetic communication or on molecular communication.

Electromagnetic

This is defined as the transmission and reception of electromagnetic radiation from components based on novel nanomaterials.^[3] Recent advancements in carbon and molecular electronics have opened the door to a new generation of electronic nanoscale components such as nanobatteries,^[4] nanoscale energy harvesting systems,^[5] nano-memories,^[6] logical circuitry in the nanoscale and even nano-antennas.^[7][8] From a communication perspective, the unique properties observed in nanomaterials will decide on the specific bandwidths for emission of electromagnetic radiation, the time lag of the emission, or the magnitude of the emitted power for a given input energy, amongst others.

For the time being, two main alternatives for electromagnetic communication in the nanoscale have been envisioned. First, it has been experimentally demonstrated that is possible to receive and demodulate an electromagnetic wave by means of a nanoradio, i.e., an electromechanically resonating carbon nanotube which is able to decode an amplitude or frequency modulated wave.^[9] Second, graphene-based nano-antennas have been analyzed as potential electromagnetic radiators in the Terahertz band^[10]

Molecular

Molecular communication is defined as the transmission and reception of information by means of molecules^[11]. The different molecular communication techniques can be classified according to the type of molecule propagation in walkaway-based, flow-based or diffusion-based communication.

In walkway-based molecular communication, the molecules propagate through pre-defined pathways by using carrier substances, such as molecular motors.^[12] This type of molecular communication can also be achieved by using E. coli bacteria as chemotaxis.^[13]

In flow-based molecular communication, the molecules propagate through diffusion in a fluidic medium whose flow and turbulence are guided and predictable. The hormonal communication through blood streams inside the human body is an example of this type of propagation. The flow-based propagation can also be realized by using carrier entities whose motion can be constrained on the average along specific paths, despite showing a random component. A good example of this case is given by pheromonal long range molecular communications.^[14]

In diffusion-based molecular communication, the molecules propagate through spontaneous diffusion in a fluidic medium. In this case, the molecules can be subject solely to the laws of diffusion or can also be affected by non-predictable turbulence present in the fluidic medium. Pheromonal communication, when pheromones are released into a fluidic medium, such as air or water, is an example of diffusion-based architecture. Other examples of this kind of transport include calcium signaling among cells ^[15], as well as quorum sensing among bacteria.^[16]

Based on the macroscopic theory^[17] of ideal (free) diffusion the impulse response of a unicast molecular communication channel was reported in a paper^[18] that identified that the impulse response of the ideal diffusion based molecular communication channel experiences temporal spreading. Such temporal spreading has a deep impact in the performance of the system e.g. in creating the intersymbol interference (ISI) at the receiving nanomachine.^[19] In order to detect the concentration-encoded molecular signal two detection methods named sampling-based detection (SD) and energy-based detection (ED) have been proposed.^[20] While the SD approach is based on the concentration amplitude of only one sample taken at a suitable time instant during the symbol duration, the ED approach is based on the total accumulated number of molecules received during the entire symbol duration. In order to reduce the impact of ISI a controlled pulse-width based molecular communication scheme has been analysed.^[21] The work presented in ^[22] showed that it is possible to realize multilevel amplitude modulation based on ideal diffusion. A comprehensive study of pulse-based binary^[23] and sinus-based,^{[24][25][26][27]} concentration-encoded molecular communication system have also been investigated.

Cell signaling

From Wikipedia, the free encyclopedia

Cell signaling (cell signalling in British English) is part of any communication process that governs basic activities of cells and coordinates all cell actions. The ability of cells to perceive and correctly respond to their microenvironment is the basis of development, tissue repair, and immunity, as well as normal tissue homeostasis. Errors in signaling interactions and cellular information processing are responsible for diseases such as cancer, autoimmunity, and diabetes.^[1][2][3] By understanding cell signaling, diseases may be treated more effectively and, theoretically, artificial tissues may be created.^[4]

Traditional work in biology has focused on studying individual parts of cell signaling pathways. Systems biology research helps us to understand the underlying structure of cell signaling networks and how changes in these networks may affect the transmission and flow of information (signal transduction). Such networks are complex systems in their organization and may exhibit a number of emergent properties including bistability and ultrasensitivity. Analysis of cell signaling networks requires a combination of experimental and theoretical approaches including the development and analysis of simulations and modeling.^{[5][citation needed][6]} Long-range allostery is often a significant component of cell signaling events.^[7]

Signaling between cells of one organism and multiple organisms

Figure 1. Example of signaling between bacteria. Salmonella enteritidis uses N-Acyl homoserine lactone for Quorum sensing (see: Inter-Bacterial Communication)

Cell signaling has been most extensively studied in the context of human diseases and signaling between cells of a single organism. However, cell signaling may also occur between the cells of two different organisms. In many mammals, early embryo cells exchange signals with cells of the uterus.^[8] In the human gastrointestinal tract, bacteria exchange signals with each other and with human epithelial and immune system cells.^[9] For the yeast Saccharomyces cerevisiae during mating, some cells send a peptide signal (mating factor pheromones) into their environment. The mating factor peptide may bind to a cell surface receptor on other yeast cells and induce them to prepare for mating.^[10]

Classification

Cell signaling can be classified to be mechanical and biochemical based on the type of the signal. Mechanical signals are the forces exerted on the cell and the forces produced by the cell. These forces can both be sensed and responded by the cells.^[11] Biochemical signals are the biochemical molecules such as proteins, lipids, ions and gases. These signals can be categorized based on the distance between signaling and responder cells. Signaling within, between, and amongst cells is subdivided into the following classifications:

• Intracrine signals are produced by the target cell that stay within the target cell.
• Autocrine signals are produced by the target cell, are secreted, and affect the target cell itself via receptors. Sometimes autocrine cells can target cells close by if they are the same type of cell as the emitting cell. An example of this are immune cells.
• Juxtacrine signals target adjacent (touching) cells. These signals are transmitted along cell membranes via protein or lipid components integral to the membrane and are capable of affecting either the emitting cell or cells immediately adjacent.
• Paracrine signals target cells in the vicinity of the emitting cell. Neurotransmitters represent an example.
• Endocrine signals target distant cells. Endocrine cells produce hormones that travel through the blood to reach all parts of the body.

Figure 2. Notch-mediated juxtacrine signal between adjacent cells.

Cells communicate with each other via direct contact (juxtacrine signaling), over short distances (paracrine signaling), or over large distances and/or scales (endocrine signaling).

Some cell–cell communication requires direct cell–cell contact. Some cells can form gap junctions that connect their cytoplasm to the cytoplasm of adjacent cells. In cardiac muscle, gap junctions between adjacent cells allows for action potential propagation from the cardiac pacemaker region of the heart to spread and coordinately cause contraction of the heart.

The notch signaling mechanism is an example of juxtacrine signaling (also known as contact-dependent signaling) in which two adjacent cells must make physical contact in order to communicate. This requirement for direct contact allows for very precise control of cell differentiation during embryonic development. In the worm Caenorhabditis elegans, two cells of the developing gonad each have an equal chance of terminally differentiating or becoming a uterine precursor cell that continues to divide. The choice of which cell continues to divide is controlled by competition of cell surface signals. One cell will happen to produce more of a cell surface protein that activates the Notch receptor on the adjacent cell. This activates a feedback loop or system that reduces Notch expression in the cell that will differentiate and that increases Notch on the surface of the cell that continues as a stem cell.^[12]

Many cell signals are carried by molecules that are released by one cell and move to make contact with another cell. Endocrine signals are called hormones. Hormones are produced by endocrine cells and they travel through the blood to reach all parts of the body. Specificity of signaling can be controlled if only some cells can respond to a particular hormone. Paracrine signals such as retinoic acid target only cells in the vicinity of the emitting cell.^[13] Neurotransmitters represent another example of a paracrine signal. Some signaling molecules can function as both a hormone and a neurotransmitter. For example, epinephrine and norepinephrine can function as hormones when released from the adrenal gland and are transported to the heart by way of the blood stream. Norepinephrine can also be produced by neurons to function as a neurotransmitter within the brain.^[14] Estrogen can be released by the ovary and function as a hormone or act locally via paracrine or autocrine signaling.^[15] Active species of oxygen and nitric oxide can also act as cellular messengers. This process is dubbed redox signaling.

Cell signaling in multicellular organisms

In a multicellular organism, signaling between cells occurs either through release into the extracellular space, divided in paracrine signaling (over short distances) and endocrine signaling (over long distances), or by direct contact, known as juxtacrine signaling.^[16] Autocrine signaling is a special case of paracrine signaling where the secreting cell has the ability to respond to the secreted signaling molecule.^[17] Synaptic signaling is a special case of paracrine signaling (for chemical synapses) or juxtacrine signaling (for electrical synapses) between neurons and target cells. Signaling molecules interact with a target cell as a ligand to cell surface receptors, and/or by entering into the cell through its membrane or endocytosis for intracrine signaling. This generally results in the activation of second messengers, leading to various physiological effects.

A particular molecule is generally used in diverse modes of signaling, and therefore a classification by mode of signaling is not possible. At least three important classes of signaling molecules are widely recognized, although non-exhaustive and with imprecise boundaries, as such membership is non-exclusive and depends on the context:

• Hormones are the major signaling molecules of the endocrine system, though they often regulate each other's secretion via local signaling (e.g. islet of Langerhans cells), and most are also expressed in tissues for local purposes (e.g. angiotensin) or failing that, structurally related molecules are (e.g. PTHrP).
• Neurotransmitters are signaling molecules of the nervous system, also including neuropeptides and neuromodulators. Neurotransmitters like the catecholamines are also secreted by the endocrine system into the systemic circulation.
• Cytokines are signaling molecules of the immune system, with a primary paracrine or juxtacrine role, though they can during significant immune responses have a strong presence in the circulation, with systemic effect (altering iron metabolism or body temperature). Growth factors can be considered as cytokines or a different class.

Signaling molecules can belong to several chemical classes: lipids, phospholipids, amino acids, monoamines, proteins, glycoproteins, or gases. Signaling molecules binding surface receptors are generally large and hydrophilic (e.g. TRH, Vasopressin, Acetylcholine), while those entering the cell are generally small and hydrophobic (e.g. glucocorticoids, thyroid hormones, cholecalciferol, retinoic acid), but important exceptions to both are numerous, and a same molecule can act both via surface receptor or in an intracrine manner to different effects.^[17] In intracrine signaling, once inside the cell, a signaling molecule can bind to intracellular receptors, other elements, or stimulate enzyme activity (e.g. gasses). The intracrine action of peptide hormones remains a subject of debate.^[18]

Hydrogen sulfide is produced in small amounts by some cells of the human body and has a number of biological signaling functions. Only two other such gases are currently known to act as signaling molecules in the human body: nitric oxide and carbon monoxide.^[19]

Receptors for cell motility and differentiation

Cells receive information from their neighbors through a class of proteins known as receptors. Notch is a cell surface protein that functions as a receptor. Animals have a small set of genes that code for signaling proteins that interact specifically with Notch receptors and stimulate a response in cells that express Notch on their surface. Molecules that activate (or, in some cases, inhibit) receptors can be classified as hormones, neurotransmitters, cytokines, and growth factors, in general called receptor ligands. Ligand receptor interactions such as that of the Notch receptor interaction, are known to be the main interactions responsible for cell signaling mechanisms and communication.^[20]

As shown in Figure 2 (above; left), notch acts as a receptor for ligands that are expressed on adjacent cells. While some receptors are cell surface proteins, others are found inside cells. For example, estrogen is a hydrophobic molecule that can pass through the lipid bilayer of the membranes. As part of the endocrine system, intracellular estrogen receptors from a variety of cell types can be activated by estrogen produced in the ovaries.

A number of transmembrane receptors^[21][22] for small molecules and peptide hormones,^[23] as well as intracellular receptors for steroid hormones exist, giving cells the ability to respond to a great number of hormonal and pharmacological stimuli. In diseases, often, proteins that interact with receptors are aberrantly activated, resulting in constitutively activated downstream signals.^[24]

For several types of intercellular signaling molecules that are unable to permeate the hydrophobic cell membrane due to their hydrophilic nature, the target receptor is expressed on the membrane. When such a signaling molecule activates its receptor, the signal is carried into the cell usually by means of a second messenger such as cAMP.^[25][26]

Signaling pathways

 

Overview of signal transduction pathways

 

Figure 3. Key components of a signal transduction pathway (MAPK/ERK pathway shown)

In some cases, receptor activation caused by ligand binding to a receptor is directly coupled to the cell's response to the ligand. For example, the neurotransmitter GABA can activate a cell surface receptor that is part of an ion channel. GABA binding to a GABA_A receptor on a neuron opens a chloride-selective ion channel that is part of the receptor. GABA_A receptor activation allows negatively charged chloride ions to move into the neuron, which inhibits the ability of the neuron to produce action potentials. However, for many cell surface receptors, ligand-receptor interactions are not directly linked to the cell's response. The activated receptor must first interact with other proteins inside the cell before the ultimate physiological effect of the ligand on the cell's behavior is produced. Often, the behavior of a chain of several interacting cell proteins is altered following receptor activation. The entire set of cell changes induced by receptor activation is called a signal transduction mechanism or pathway.^[27]

In the case of Notch-mediated signaling, the signal transduction mechanism can be relatively simple. As shown in Figure 2, activation of Notch can cause the Notch protein to be altered by a protease. Part of the Notch protein is released from the cell surface membrane and takes part in gene regulation. Cell signaling research involves studying the spatial and temporal dynamics of both receptors and the components of signaling pathways that are activated by receptors in various cell types.^{[citation needed]}

A more complex signal transduction pathway is shown in Figure 3. This pathway involves changes of protein–protein interactions inside the cell, induced by an external signal. Many growth factors bind to receptors at the cell surface and stimulate cells to progress through the cell cycle and divide. Several of these receptors are kinases that start to phosphorylate themselves and other proteins when binding to a ligand. This phosphorylation can generate a binding site for a different protein and thus induce protein–protein interaction. In Figure 3, the ligand (called epidermal growth factor (EGF)) binds to the receptor (called EGFR). This activates the receptor to phosphorylate itself. The phosphorylated receptor binds to an adaptor protein (GRB2), which couples the signal to further downstream signaling processes. For example, one of the signal transduction pathways that are activated is called the mitogen-activated protein kinase (MAPK) pathway. The signal transduction component labeled as "MAPK" in the pathway was originally called "ERK," so the pathway is called the MAPK/ERK pathway. The MAPK protein is an enzyme, a protein kinase that can attach phosphate to target proteins such as the transcription factor MYC and, thus, alter gene transcription and, ultimately, cell cycle progression. Many cellular proteins are activated downstream of the growth factor receptors (such as EGFR) that initiate this signal transduction pathway.^{[citation needed]}

Some signaling transduction pathways respond differently, depending on the amount of signaling received by the cell. For instance, the hedgehog protein activates different genes, depending on the amount of hedgehog protein present.^{[citation needed]}

Complex multi-component signal transduction pathways provide opportunities for feedback, signal amplification, and interactions inside one cell between multiple signals and signaling pathways.^{[citation needed]}

Intraspecies and interspecies signaling

Molecular signaling can occur between different organisms, whether unicellular or multicellular. The emitting organism produces the signaling molecule, secretes it into the environment, where it diffuses, and it is sensed or internalized by the receiving organism. In some cases of interspecies signaling, the emitting organism can actually be a host of the receiving organism, or vice versa.

Intraspecies signaling occurs especially in bacteria, yeast, social insects, but also many vertebrates. The signaling molecules used by multicellular organisms are often called pheromones. They can have such purposes as alerting against danger, indicating food supply, or assisting in reproduction.^[28] In unicellular organisms such as bacteria, signaling can be used to 'activate' peers from a dormant state, enhance virulence, defend against bacteriophages, etc.^[29] In quorum sensing, which is also found in social insects, the multiplicity of individual signals has the potentiality to create a positive feedback loop, generating coordinated response. In this context, the signaling molecules are called autoinducers.^[30][31][32] This signaling mechanism may have been involved in evolution from unicellular to multicellular organisms.^[30][33] Bacteria also use contact-dependent signaling, notably to limit their growth.^[34]

Molecular signaling can also occur between individuals of different species. This has been particularly studied in bacteria.^[35][36][37] Different bacterial species can coordinate to colonize a host and participate in common quorum sensing.^[38] Therapeutic strategies to disrupt this phenomenon are being investigated.^[39][40] Interactions mediated through signaling molecules are also thought to occur between the gut flora and their host, as part of their commensal or symbiotic relationship.^[40][41] Gram negative microbes deploy bacterial outer membrane vesicles for intra- and inter-species signaling in natural environments and at the host-pathogen interface.

Additionally, interspecies signaling occurs between multicellular organisms. In Vespa mandarinia, individuals release a scent that directs the colony to a food source.^[42]