Silicon chapter
Silicon has four valence electrons and thus it looks like a quadrant. It is no coincidence that carbon, which also looks like a quadrant, and silicon are both considered the miracle elements because of how important they are.
QMRSilicon sulfide is the inorganic compound with the formula SiS2. Like the silicon dioxide, this material is polymeric, but it adopts a 1-dimensional structure quite different from the usual forms of SiO2.
Synthesis, structure, and properties[edit]
The material is formed by heating silicon and sulfur or by the exchange reaction between SiO2 and Al2S3. The material consists of chains of edge-shared tetrahedra, Si(μ-S)2Si(μS)2, etc.[2]
Like other silicon sulfur-compounds (e.g., bis(trimethylsilyl)sulfide) SiS2 hydrolyzes readily to release H2S. In liquid ammonia it is reported to form the imide Si(NH)2 and NH4SH,[3] but a recent report has identified crystalline (NH4)2[SiS3(NH3)]·2NH3 as a product which contains the tetrahedral thiosilicate anion, SiS3(NH3).[4]
Reaction with ethanol gives the alkoxide, tetraethyl orthosilicate and H2S.[3] Reaction with sodium sulfide, magnesium sulfide and aluminum sulfide give thiosilicates.[3]
SiS2 is claimed to occur in certain interstellar objects.[
QMRTetraethyl orthosilicate is the chemical compound with the formula Si(OC2H5)4. Often abbreviated TEOS, it is a colorless liquid that degrades in water. TEOS is the ethyl ester of orthosilicic acid, Si(OH)4. It is a prototypical alkoxide.
TEOS is a tetrahedral molecule. Like its many analogues, TEOS is prepared by alcoholysis of silicon tetrachloride:
SiCl4 + 4 EtOH → Si(OEt)4 + 4 HCl
where Et = C2H5
Applications[edit]
TEOS is mainly used as a crosslinking agent in silicone polymers and as a precursor to silicon dioxide in the semiconductor industry.[3] TEOS is also used as the silica source for synthesis of zeolites.[4] Other applications include coatings for carpets and other objects. TEOS is used in the production of aerogel. These applications exploit the reactivity of the Si-OR bonds.[5]
Other reactions[edit]
TEOS has the remarkable property of easily converting into silicon dioxide. This reaction occurs upon the addition of water:
Si(OC2H5)4 + 2 H2O → SiO2 + 4 C2H5OH
This hydrolysis reaction is an example of a sol-gel process. The side product is ethanol. The reaction proceeds via a series of condensation reactions that convert the TEOS molecule into a mineral-like solid via the formation of Si-O-Si linkages. Rates of this conversion are sensitive to the presence of acids and bases, both of which serve as catalysts. The Stöber process allows the formation of monodisperse silica particles.
At elevated temperatures (>600 °C), TEOS converts to silicon dioxide:
Si(OC2H5)4 → SiO2 + 2 (C2H5)2O
The volatile coproduct is diethyl ether.
Safety[edit]
Although nontoxic to ingestion, TEOS vapor is highly damaging to eyes since it deposits silica.[6]
QMRTetramethyl orthosilicate is the chemical compound with the formula Si(OCH3)4. This molecule consists of four methyl groups attached to the hypothetical anion SiO44−. The basic properties are similar to the more popular tetraethyl orthosilicate, which is usually preferred because the product of hydrolysis, ethanol, is less toxic than methanol.
Tetramethyl orthosilicate hydrolyzes to SiO2:
Si(OCH3)4 + 2 H2O → SiO2 + 4 CH3OH
In organic synthesis, Si(OCH3)4 has been used to convert ketones and aldehydes to the corresponding ketals and acetals, respectively.[2]
Safety[edit]
The hydrolysis of Si(OCH3)4 produces insoluble SiO2 and CH3OH (methanol). Even at low concentrations inhalation causes lung lesions, and at slightly higher concentrations eye contact with the vapor causes blindness. Worse, at low concentrations (200ppm/15min) the damage is often insidious, with onset of symptoms hours after exposure.[3] The mode of action is the precipitation of silica in the eyes and/or lungs. Contrary to common information, including several erroneous MSDS sheets, the methanol produced is relatively innocuous. The mechanisms of methanol toxicity are well established, methanol causes blindness via conversion to toxic formaldehyde in the liver; methanol splashes to the eye cause only moderate and reversible eye irritation.[4]
Tetraethyl orthosilicate is safer for two reasons. 1) It is considerably less volatile. 2) It is hydrolyzed more slowly by the water in the lungs and on the eyes (due to steric bulk).
QMRAn acetal is a functional group with the following connectivity R2C(OR')2, where both R' groups are organic fragments. The central carbon atom has four bonds to it, and is therefore saturated and has tetrahedral geometry.
Electronic properties[edit]
The diameter of carbon peapods range from ca. 1 to 50 nanometers. Various combinations of fullerene C60 sizes and nanotube structures can lead to various electric conductivity property of carbon peapods due to orientation of rotations. For example, the C60 @ (10,10) is a good superconductor and the C60 @ (17,0) peapod is a semiconductor. The calculated band gap of C60 @ (17,0) equals 0.1 eV.[13] Research into their potential as semiconductors is still ongoing. Although both the doped fullerides and ropes of SWNTs are superconductors, unfortunately, the critical temperatures for the superconducting phase transition in these materials are low. There are hopes that carbon nano-peapods could be superconducting at room temperature.[14]
With chemical doping, the electronic characteristics of peapods can be further adjusted. When carbon peapod is doped with alkali metal atoms like potassium, the dopants will react with the C60 molecules inside the SWNT. It forms a negatively charged C606− covalently bound, one-dimensional polymer chain with metallic conductivity. Overall, the doping of SWNTs and peapods by alkali metal atoms actively enhances the conductivity of the molecule since the charge is relocated from the metal ions to the nanotubes.[15] Doping carbon nanotubes with oxidized metal is another way to adjust conductivity. It creates a very interesting high temperature superconducting state as the Fermi level is significantly reduced. A good application would be the introduction of silicon dioxide to carbon nanotubes. It constructs memory effect as some research group has invented ways to create memory devices based on carbon peapods grown on Si/SiO2 surfaces.[
QMRSilicones are polymers that include any inert, synthetic compound made up of repeating units of siloxane, which is a chain of alternating silicon atoms and oxygen atoms, frequently combined with carbon and/or hydrogen. They are typically heat-resistant and rubber-like, and are used in sealants, adhesives, lubricants, medicine, cooking utensils, and thermal and electrical insulation. Some common forms include silicone oil, silicone grease, silicone rubber, silicone resin, and silicone caulk.[1]
More precisely called polymerized siloxanes or polysiloxanes, silicones consist of an inorganic silicon-oxygen backbone chain (⋯-Si-O-Si-O-Si-O-⋯) with organic side groups attached to the silicon atoms. These silicon atoms are tetravalent. So, silicones are polymers constructed from inorganic-organic monomers. Silicones have in general the chemical formula [R2SiO]n, where R is an organic group such as methyl, ethyl, or phenyl.
In some cases, organic side groups can be used to link two or more of these -Si-O- backbones together. By varying the -Si-O- chain lengths, side groups, and crosslinking, silicones can be synthesized with a wide variety of properties and compositions. They can vary in consistency from liquid to gel to rubber to hard plastic. The most common siloxane is linear polydimethylsiloxane (PDMS), a silicone oil. The second largest group of silicone materials is based on silicone resins, which are formed by branched and cage-like oligosiloxanes.
F. S. Kipping and Matt Saunders coined the word silicone in 1901 to describe polydiphenylsiloxane by analogy of its formula, Ph2SiO (Ph stands for phenyl, C6H5), with the formula of the ketone benzophenone, Ph2CO (his term was originally silicoketone). Kipping was well aware that polydiphenylsiloxane is polymeric whereas benzophenone is monomeric and noted that Ph2SiO and Ph2CO had very different chemistry.[2][3] The discovery of the structural differences between Kippings' molecules and the ketones means that silicone is no longer the correct term (though it remains in common usage) and that the term siloxanes is correct according to the nomenclature of modern chemistry.[4]
Silicone is sometimes mistakenly referred to as silicon. The chemical element silicon is a crystalline metalloid widely used in computers and other electronic equipment. Although silicones contain silicon atoms, they also include carbon, hydrogen, oxygen, and perhaps other kinds of atoms as well, and have very different physical and chemical properties to elemental silicon.
A true silicone group with a double bond between oxygen and silicon does not commonly exist in nature; chemists find that the silicon atom usually forms single bonds with each of two oxygen atoms, rather than a double bond to a single atom. Polysiloxanes are among the many substances commonly known as "silicones".
Molecules containing silicon-oxygen double bonds do exist and are called silanones but they are very reactive. Despite this, silanones are important as intermediates in gas-phase processes such as chemical vapor deposition in microelectronics production, and in the formation of ceramics by combustion.[5]
Combustion[edit]
When silicone is burned in air or oxygen, it forms solid silica (silicon dioxide) as a white powder, char, and various gases. The readily dispersed powder is sometimes called silica fume.
Properties[edit]
Silicones exhibit many useful characteristics, including:[1]
Low thermal conductivity
Low chemical reactivity
Low toxicity
Thermal stability (constancy of properties over a wide temperature range of −100 to 250 °C).
The ability to repel water and form watertight seals.
Does not stick to many substrates, but adheres very well to others, e.g. glass.
Does not support microbiological growth.
Resistance to oxygen, ozone, and ultraviolet (UV) light. This property has led to widespread use of silicones in the construction industry (e.g. coatings, fire protection, glazing seals) and the automotive industry (external gaskets, external trim).
Electrical insulation properties. Because silicone can be formulated to be electrically insulative or conductive, it is suitable for a wide range of electrical applications.
High gas permeability: at room temperature (25 °C), the permeability of silicone rubber for such gases as oxygen is approximately 400 times[citation needed] that of butyl rubber, making silicone useful for medical applications in which increased aeration is desired. Consequently, silicone rubbers cannot be used where gas-tight seals are necessary.
Uses
Uses[edit]
Silicones are used in many products. Ullmann's Encyclopedia of Industrial Chemistry lists the following major categories of application: Electrical (e.g., insulation), electronics (e.g., coatings), household (e.g., sealants for cooking apparatus), automobile (e.g., gaskets), aeroplane (e.g., seals), office machines (e.g., keyboard pads), medicine/dentistry (e.g., teeth impression molds), textiles/paper (e.g., coatings). For these applications, an estimated 400,000 tons of silicones were produced in 1991. Specific examples, both large and small are presented below.[1]
Automotive[edit]
In the automotive field, silicone grease is typically used as a lubricant for brake components since it is stable at high temperatures, is not water-soluble, and is far less likely than other lubricants to foul. It is also used as DOT 5 brake fluid.
Automotive spark plug wires are insulated by multiple layers of silicone to prevent sparks from jumping to adjacent wires, causing misfires. Silicone tubing is sometimes used in automotive intake systems (especially for engines with forced induction).
Sheet silicone is used to manufacture gaskets used in automotive engines, transmissions, and other applications.
Automotive body manufacturing plants and paint shops avoid silicones, as they may cause "fish eyes", small, circular craters in the finish.
Additionally, silicone compounds such as silicone rubber are used as coatings and sealants for airbags; the high strength of silicone rubber makes it an optimal adhesive/sealant for high impact airbags. Recent technological advancements allow convenient use of silicone in combination with thermoplastics to provide improvements in scratch and mar resistance and lowered coefficient of friction.
Silicone films can be applied to such silica-based substrates as glass to form a covalently bonded hydrophobic coating.
Many fabrics can be coated or impregnated with silicone to form a strong, waterproof composite such as silnylon.
Cookware[edit]
Soup ladle and pasta ladle made of silicone.
A silicone food steamer to be placed inside a pot of boiling water.
Ice cube trays made of silicone.
As a low-taint, non-toxic material, silicone can be used where contact with food is required. Silicone is becoming an important product in the cookware industry, particularly bakeware and kitchen utensils.
Silicone is used as an insulator in heat-resistant potholders and similar items; however, it is more conductive of heat than similar less dense fiber-based products. Silicone oven mitts are able to withstand temperatures up to 260 °C (500 °F), allowing reaching into boiling water.
Molds for chocolate, ice, cookies, muffins and various other foods.
Non-stick bakeware and reusable mats used on baking sheets.
Other products such as steamers, egg boilers or poachers, cookware lids, pot holders, trivets, and kitchen mats.
Defoaming[edit]
Silicones are used as active compound in defoamers due to their low water solubility and good spreading properties.
Dry cleaning[edit]
Liquid silicone can be used as a dry cleaning solvent, providing an alternative to the traditional chlorine-containing perchloroethylene (perc) solvent. Use of silicones in dry cleaning reduces the environmental impact of a typically high-polluting industry.
Electronics[edit]
Electronic components are sometimes encased in silicone to increase stability against mechanical and electrical shock, radiation and vibration, a process called "potting".
Silicones are used where durability and high performance are demanded of components under hard conditions, such as in space (satellite technology). They are selected over polyurethane or epoxy encapsulation when a wide operating temperature range is required (−65 to 315 °C). Silicones also have the advantage of little exothermic heat rise during cure, low toxicity, good electrical properties and high purity.
The use of silicones in electronics is not without problems, however. Silicones are relatively expensive and can be attacked by solvents. Silicone easily migrates as either a liquid or vapor onto other components.
Silicone contamination of electrical switch contacts can lead to failures by causing an increase in contact resistance, often late in the life of the contact, well after any testing is completed.[6][7] Use of silicone-based spray products in electronic devices during maintenance or repairs can cause later failures.
Firestops[edit]
Wikimedia Commons has media related to Silicone foam.
Silicone foam has been used in North American buildings in an attempt to firestop openings within fire-resistance-rated wall and floor assemblies to prevent the spread of flames and smoke from one room to another. When properly installed, silicone-foam firestops can be fabricated for building code compliance. Advantages include flexibility and high dielectric strength. Disadvantages include combustibility (hard to extinguish) and significant smoke development.
Silicone-foam firestops have been the subject of controversy and press attention due to smoke development from pyrolysis of combustible components within the foam, hydrogen gas escape, shrinkage, and cracking. These problems have led to reportable events among licensees (operators of nuclear power plants) of the Nuclear Regulatory Commission (NRC).
Silicone firestops are also used in aircraft.
Lubricants[edit]
Silicone greases are used for many purposes, such as bicycle chains, airsoft gun parts, and a wide range of other mechanisms. Typically, a dry-set lubricant is delivered with a solvent carrier to penetrate the mechanism. The solvent then evaporates, leaving a clear film that lubricates but does not attract dirt and grit as much as an oil-based or other traditional "wet" lubricant.
Silicone personal lubricants are also available for use in medical procedures or sexual activity. See below.
Medicine[edit]
Silicone is used in microfluidics, seals, gaskets, shrouds, and other applications requiring high biocompatibility. Additionally, the gel form is used in bandages and dressings, breast implants, testicle implants, pectoral implants, contact lenses, and a variety of other medical uses.
Scar treatment sheets are often made of medical grade silicone due to its durability and biocompatibility. Polydimethylsiloxane is often used for this purpose, since its specific crosslinking results in a flexible and soft silicone with high durability and tack.
Polydimethylsiloxane (PDMS) has been used as the hydrophobic block of amphiphilic synthetic block copolymers used to form the vesicle membrane of polymersomes
Moldmaking[edit]
Two-part silicone systems are used to create rubber molds used to cast resins, foams, rubber, and low-temperature alloys. A silicone mold generally requires little or no mold-release or surface preparation, as most materials do not adhere to silicone. For experimental uses, ordinary one-part silicone can be used to make molds or to mold into shapes. If needed, common vegetable cooking oils or petroleum jelly can be used on mating surfaces as a mold-release agent.[8]
Silicone Mold
Cooking molds used as bakeware do not require coating with cooking oil, allowing the baked food to be more easily removed from the mold after cooking.
Ophthalmology[edit]
Silicone has many applications like silicone oil used to replace vitreous following vitrectomy, silicone intraocular lenses following cataract extraction, silicone tubes to keep nasolacrimal passage open following dacrycystorhinostomy, canalicular stents for canalicular stenosis, punctal plugs for punctal occlusion in dry eyes, silicone rubber and bands as an external temponade in tractional retinal detachment, and anteriorly located break in rhegmatogenous retinal detachment.
Personal care[edit]
Silicones are ingredients in many hair conditioners, shampoos, and hair gel products. Some silicones, notably the amine functionalized amodimethicones, are excellent conditioners, providing improved compatibility, feel, and softness, and lessening frizz. The phenyltrimethicones, in another silicone family, are used in reflection-enhancing and color-correcting hair products, where they increase shine and glossiness (and possibly effect subtle color changes). Phenyltrimethicones, unlike the conditioning amodimethicones, have refractive indices (typically 1.46) close to that of human hair (1.54). However, if included in the same formulation, amodimethicone and phenyltrimethicone interact and dilute each other, making it difficult to achieve both high shine and excellent conditioning in the same product.[9]
Silicone rubber is commonly used in baby bottle nipples (teats) for its cleanliness, aesthetic appearance, and low extractable content.
Silicones are used in shaving products and personal lubricants.[10]
Plumbing and building construction[edit]
The strength and reliability of silicone rubber is widely acknowledged in the construction industry. One-part silicone sealants and caulks are in common use to seal gaps, joints and crevices in buildings. One-part silicones cure by absorbing atmospheric moisture, which simplifies installation. In plumbing, silicone grease is typically applied to O-rings in brass taps and valves, preventing lime from sticking to the metal.
Toys and hobbies[edit]
Silly Putty and similar materials are composed of silicones dimethyl siloxane, polydimethylsiloxane, and decamethyl cyclopentasiloxane, with other ingredients. This substance is noted for its unusual characteristic that it bounces, but breaks when given a sharp blow; it can also flow like a liquid and will form a puddle given enough time.
Silicone "rubber bands" are a long-lasting popular replacement refill for real rubber bands in the new (2013) fad "rubber band loom" toys at two to four times the price (in 2014). However, they only come in one size, small. Silicone bands also come in bracelet sizes that can be custom embossed with a name or message. Large silicone bands are also sold as utility tie-downs and such.
Formerol is a silicone rubber (marketed as Sugru) used as an arts-and-crafts material as its plasticity allows it to be moulded by hand like modeling clay. It hardens at room temperature and it is adhesive to various substances including glass and aluminum.[11]
In making aquariums, manufacturers now commonly use 100% silicone sealant to join glass plates. Glass joints made with silicone sealant can withstand great pressure, making obsolete the original aquarium construction method of angle-iron and putty. This same silicone is used to make hinges in aquarium lids or for minor repairs. However, not all commercial silicones are safe for aquarium manufacture, nor is silicone used for the manufacture of acrylic aquariums as silicones do not have long-term adhesion to plastics.[12]
Sex toys and lubricants[edit]
Silicone is a material of choice for soft sex toys, due to its durability, cleanability, non-degradation by petroleum-based lubricants, and lack of phthalates, chemicals suspected of having carcinogenic and mutagenic effects on the skin and mucous membranes.[13][14][15]
Production and marketing[edit]
The global demand for silicones approached US$12.5 billion in 2008, approximately 4% up from the previous year. It continues similar growth in the following years to reach $13.5 billion by 2010. The annual growth is expected to be boosted by broader applications, introduction of novel products and increasing awareness of using more environmentally friendly materials.[16]
QMR Silicon has four valence electrons and thus looks like a quadrant
Silicone resins are a type of silicone material which is formed by branched, cage-like oligosiloxanes with the general formula of RnSiXmOy, where R is a non reactive substituent, usually Methyl (Me) or Phenyl (Ph), and X is a functional group Hydrogen (H), Hydroxyl group (OH), Chlorine (Cl) or Alkoxy group (OR). These groups are further condensed in many applications, to give highly crosslinked, insoluble polysiloxane networks.[1]
When R is methyl, the four possible functional siloxane monomeric units are described as follows:[2]
"M" stands for Me3SiO,
"D" for Me2SiO2,
"T" for MeSiO3 and
"Q" for SiO4.
Note that a network of only Q groups becomes fused quartz.
The most abundant silicone resins are built of D and T units (DT resins) or from M and Q units (MQ resins), however many other combinations (MDT, MTQ, QDT) are also used in industry.
Silicone resins represent a broad range of products. Materials of molecular weight in the range of 1000–10,000 are very useful in pressure-sensitive adhesives, silicone rubbers, coatings and additives.[3][4]
Silicone resins are prepared by hydrolytic condensation of various silicone precursors. In early processes of preparation of silicone resins sodium silicate and various chlorosilanes were used as starting materials. Although the starting materials were the least expensive ones (something typical for industry), structural control of the product was very difficult. More recently, a less reactive tetraethoxysilane - (TEOS) or ethyl polysilicate and various disiloxanes are used as starting materials.[1]
Microbial deterioration[edit]
The algae Stichococcus bacillaris, and certain fungal species have been seen to colonize silcone resins used at archaeological sites. [5]
QMRSodium silicate is the common name for compounds with the formula Na2(SiO2)nO. A well-known member of this series is sodium metasilicate, Na2SiO3. Also known as waterglass or liquid glass, these materials are available in aqueous solution and in solid form. The pure compositions are colourless or white, but commercial samples are often greenish or blue owing to the presence of iron-containing impurities.
They are used in cements, passive fire protection, textile and lumber processing, refractories, and automobiles. Sodium carbonate and silicon dioxide react when molten to form sodium silicate and carbon dioxide:[1]
Na2CO3 + SiO2 → Na2SiO3 + CO2
Uses[edit]
In 1990, 4M tons of alkali metal silicates were produced. The main applications were in detergents, paper, water treatment, and construction materials.[3]
Adhesive[edit]
The largest application of sodium silicate solutions is a cement for producing cardboard.[3] When used as a paper cement, the tendency is for the sodium silicate joint eventually to crack within a few years, at which point it no longer holds the paper surfaces cemented together.
Drilling fluids[edit]
Sodium silicate is frequently used in drilling fluids to stabilize borehole wells and to avoid the collapse of bore walls. It is particularly useful when drill holes pass through argillaceous formations containing swelling clay minerals such as smectite or montmorillonite.
Concrete and general masonry treatment[edit]
Concrete treated with a sodium silicate solution helps to significantly reduce porosity in most masonry products such as concrete, stucco, and plasters. A chemical reaction occurs with the excess Ca(OH)2 (portlandite) present in the concrete that permanently binds the silicates with the surface, making them far more durable and water repellent. This treatment generally is applied only after the initial cure has taken place (7 days or so depending on conditions). These coatings are known as silicate mineral paint.
Detergent auxiliaries[edit]
It is used in detergent auxiliaries such as complex sodium disilicate and modified sodium disilicate. The detergent granules gain their ruggedness from a coating of silicates.[3]
Water treatment[edit]
Water glass is used as coagulant/deflocculant agent in wastewater treatment plants. Waterglass binds to colloidal molecules, creating larger aggregates that sink to the bottom of the water column. The microscopic negatively charged particles suspended in water interact with sodium silicate. Their electrical double layer collapses due to the increase of ionic strength caused by the addition of sodium silicate (doubly negatively charged anion accompanied by two sodium cations) and they subsequently aggregate. This process is called coagulation/deflocculation.[3]
Refractory use[edit]
Water glass is a useful binder of solids, such as vermiculite and perlite. When blended with the aforementioned lightweight aggregates, water glass can be used to make hard, high-temperature insulation boards used for refractories, passive fire protection and high temperature insulations, such as moulded pipe insulation applications. When mixed with finely divided mineral powders, such as vermiculite dust (which is common scrap from the exfoliation process), one can produce high temperature adhesives. The intumescence disappears in the presence of finely divided mineral dust, whereby the waterglass becomes a mere matrix. Waterglass is inexpensive and abundantly available, which makes its use popular in many refractory applications.
Dye auxiliary[edit]
Sodium silicate solution is used as a fixative for hand dyeing with reactive dyes that require a high pH to react with the textile fiber. After the dye is applied to a cellulose-based fabric, such as cotton or rayon, or onto silk, it is allowed to dry, after which the sodium silicate is painted on to the dyed fabric, covered with plastic to retain moisture, and left to react for an hour at room temperature.[10]
Niche and hobby uses[edit]
Passive fire protection[edit]
Expantrol proprietary sodium silicate suspended in about a 6.5-mm-thick layer of red rubber, type 3M FS195, inserted into a metal pipe, then heated, to demonstrate hard char intumescence, strong enough to shut a melting plastic pipe
Palusol-based intumescent plastic pipe device used for commercial firestopping
Sodium silicates are inherently intumescent. They come in prill (solid beads) form, as well as the liquid, water glass. The solid sheet form (Palusol) must be waterproofed to ensure long-term passive fire protection (PFP).
Standard, solid, bead-form sodium silicates have been used as aggregate within silicone rubber to manufacture plastic pipe firestop devices. The silicone rubber was insufficient waterproofing to preserve the intumescing function and the products had to be recalled, which is problematic for firestops concealed behind drywall in buildings.
Pastes for caulking purposes are similarly unstable. This, too, has resulted in recalls and even litigation. Only 3M's "Expantrol" version, which has an external heat treatment that helps to seal the outer surface, as part of its process standard, has achieved sufficient longevity to qualify for DIBt approvals in the US for use in firestopping.
Not unlike other intumescents, sodium silicate, both in bead form and in liquid form, are inherently endothermic, due to liquid water in the water glass and hydrates in the prill form. The absence in the US of mandatory aging tests, whereby PFP systems are made to undergo system performance tests after the aging and humidity exposures, are at the root of the continued availability, in North America, of PFP products that can become inoperable within weeks of installation. Indiscriminate use of sodium silicates without proper waterproofing measures are contributors to the problems and risk. When sodium silicates are adequately protected, they function extremely well and reliably for long periods. Evidence of this can be seen in the many DIBt approvals for plastic pipe firestop devices using Palusol, which use waterproofed sodium silicate sheets.
Food preservation[edit]
World War I poster suggesting the use of waterglass to preserve eggs (lower right).
Sodium silicate was also used as an egg preservation agent through the early 20th century with large success. When fresh eggs are immersed in it, bacteria which cause the eggs to spoil are kept out and water is kept in. Eggs can be kept fresh using this method for up to five months. When boiling eggs preserved this way, it is well advised to pin-prick the egg to allow steam to escape because the shell is no longer porous.[11]
Metal repair[edit]
Sodium silicate is used, along with magnesium silicate, in muffler repair and fitting paste. When dissolved in water, both sodium silicate and magnesium silicate form a thick paste that is easy to apply. When the exhaust system of an internal combustion engine heats up to its operating temperature, the heat drives out all of the excess water from the paste. The silicate compounds that are left over have glass-like properties, making a temporary, brittle repair.
Automotive repair[edit]
Sodium silicate is also used currently as an exhaust system joint and crack sealer for repairing mufflers, resonators, tailpipes, and other exhaust components, with and without fiberglass reinforcing tapes. In this application, the sodium silicate (60-70%) is typically mixed with kaolin (40-30%), an aluminium silicate mineral, to make the sodium silicate "glued" joint opaque. The sodium silicate, however, is the high-temperature adhesive; the kaolin serves simply as a compatible high-temperature coloring agent. Some of these repair compounds also contain glass fibres to enhance their gap-filling abilities and reduce brittleness.
Sodium silicate can be used to fill gaps within the head gasket. Commonly used on aluminum alloy cylinder heads, which are sensitive to thermally induced surface deflection. This can be caused by many things including head-bolt stretching, deficient coolant delivery, high cylinder head pressure, overheating, etc.
"Liquid glass" (sodium silicate) is added to the system through the radiator, and allowed to circulate. Sodium silicate is suspended in the coolant until it reaches the cylinder head. At 100–105°C, sodium silicate loses water molecules to form a glass seal with a remelt temperature above 810°C.
A sodium silicate repair can last two years or longer. The repair occurs rapidly, and symptoms disappear instantly. This repair only works when the sodium silicate reaches its "conversion" temperature at 100–105°C. Contamination of engine oil is a serious possibility in situations in which a coolant-to-oil leak is present. Sodium silicate (glass particulate) contamination of lubricants is detrimental to their function.
Sodium silicate solution is used to inexpensively, quickly, and permanently disable automobile engines. Running an engine with about 2 liters of a sodium silicate solution instead of motor oil causes the solution to precipitate, catastrophically damaging the engine's bearings and pistons within a few minutes.[12] In the United States, this procedure was used to comply with requirements of the Car Allowance Rebate System (CARS) program.[12][13
Homebrewing[edit]
Sodium silicate flocculant properties are also used to clarify wine and beer by precipitating colloidal particles. As a clearing agent, though, sodium silicate (water glass) is sometimes confused with isinglass which is prepared from collagen extracted from the dried swim bladders of sturgeon and other fishes. Eggs preserved in a bucket of waterglass gel, and their shells, are sometimes also used (baked and crushed) to clear wine.[14]
Aquaculture[edit]
Sodium silicate gel is also used as a substrate for algal growth in aquaculture hatcheries.
Safe construction[edit]
A mixture of sodium silicate and sawdust has been used in between the double skin of certain safes. This not only makes them more fire resistant, but also makes cutting them open with an oxyacetylene torch extremely difficult due to the smoke emitted.
Crystal gardens[edit]
When crystals of a number of metallic salts are dropped into a solution of water glass, simple or branching stalagmites of coloured metal silicates are formed. This phenomenon has been used by manufacturers of toys and chemistry sets to provide instructive enjoyment to many generations of children from the early 20th century until the present. An early mention of crystals of metallic salts forming a "chemical garden" in sodium silicate is found in the 1946 Modern Mechanix magazine.[15] Metal salts used included the sulfates and/or chlorides of copper, cobalt, iron, nickel, and manganese.
Pottery[edit]
Sodium silicate is used as a deflocculant in casting slips helping reduce viscosity and the need for large amounts of water to liquidize the clay body. It is also used to create a crackle effect in pottery, usually wheel-thrown. A vase or bottle is thrown on the wheel, fairly narrow and with thick walls. Sodium silicate is brushed on a section of the piece. After 5 minutes, the wall of the piece is stretched outward with a rib or hand. The result is a wrinkled or cracked look.
It is also the main agent in "magic water", which is used when joining clay pieces, especially if the moisture level of the two differs.[16]
Sealing of leaking water-containing structures[edit]
Sodium silicate with additives was injected into the ground to harden it and thereby to prevent further leakage of highly radioactive water from the Fukushima Daiichi nuclear power plant in Japan in April, 2011.[17] The residual heat carried by the water used for cooling the damaged reactors accelerated the setting of the injected mixture.
On June 3, 1958, the USS Nautilus, the world's first nuclear submarine, visited Everett and Seattle. In Seattle, crewmen dressed in civilian clothing were sent in to secretly buy 140 quarts of an automotive product containing sodium silicate (originally identified as Stop Leak) to repair a leaking condenser system. The Nautilus was en route to the North Pole on a top secret mission to cross the North Pole submerged.[18]
Cartridges[edit]
A historical use of the adhesive properties of sodium silicates is the production of paper cartridges for black powder revolvers produced by Colt's Manufacturing Company during the period from 1851 until 1873, especially during the American Civil War. Sodium silicate was used to seal combustible nitrated paper together to form a conical paper cartridge to hold the black powder, as well as to cement the lead ball or conical bullet into the open end of the paper cartridge. Such sodium silicate cemented paper cartridges were inserted into the cylinders of revolvers, thereby speeding the reloading of cap-and-ball black powder revolvers. This use largely ended with the introduction of Colt revolvers employing brass-cased cartridges starting in 1873.[19][20] Similarly, sodium silicate was also used to cement the top wad into brass shotgun shells, thereby eliminating any need for a crimp at the top of the brass shotgun shell to hold a shotgun shell together. Reloading brass shotgun shells was widely practiced by self-reliant American farmers during the 1870s, using the same waterglass material that was also used to preserve eggs. The cementing of the top wad on a shotgun shell consisted of applying from three to five drops of waterglass on the top wad to secure it to the brass hull. Brass hulls for shotgun shells were superseded by paper hulls starting around 1877. The newer paper-hulled shotgun shells used a roll crimp in place of a waterglass-cemented joint to hold the top wad in the shell. However, whereas brass shotshells with top wads cemented with waterglass could be reloaded nearly indefinitely (given powder, wad, and shot, of course), the paper hulls that replaced the brass hulls could be reloaded only a few times.
QMRThe silicate minerals are rock-forming minerals, constituting approximately 90 percent of the crust of the Earth. They are classified based on the structure of their silicate group which contain different ratios of silicon and oxygen. They make up the largest and most important class of rock-forming minerals.
QMRCoesite is a form (polymorph) of silicon dioxide SiO2 that is formed when very high pressure (2–3 gigapascals), and moderately high temperature (700 °C or 1,300 °F), are applied to quartz. Coesite was first synthesized by Loring Coes, Jr., a chemist at the Norton Company, in 1953.[2][3]
QMRQuartz-porphyry, in layman's terms, is a type of volcanic (igneous) rock containing large porphyritic crystals of quartz.[1][2] These rocks are classified as hemi-crystalline acid rocks.
The quartz crystals exist in a fine-grained matrix, usually of micro-crystalline or felsitic structure. In specimens, the quartz appears as small rounded, clear, greyish, vitreous blebs, which are crystals, double hexagonal pyramids, with their edges and corners rounded by resorption or corrosion.
Under the microscope they are often seen to contain rounded enclosures of the ground-mass or fluid cavities, which are frequently negative crystals with regular outlines resembling those of perfect quartz crystals. Many of the latter contain liquid carbonic acid and a bubble of gas that may exhibit vibratile motion under high magnifying powers.
Variants[edit]
In addition to quartz there are usually phenocrysts of feldspar, mostly orthoclase, though a varying amount of plagioclase is often present. The feldspars are usually full and cloudy from the formation of secondary kaolin and muscovite throughout their substance. Their crystals are larger than those of quartz and sometimes attain a length of two inches.
Not uncommonly scales of biotite are visible in the specimens, being hexagonal plates, which may be weathered into a mixture of chlorite and epidote.
Apatite, magnetite, and zircon, all in small but frequently perfect crystals, are almost universal minerals of the quartz-porphyries. The ground-mass is finely crystalline and to the unaided eye has usually a dull aspect resembling common earthenware; it is grey, green, reddish or white. Often it is streaked or banded by flow during cooling, but as a rule these rocks are not vesicular.
Two main types may be recognized by means of the microscope; the felsitic and the microcrystalline. In the former the ingredients are so fine-grained that in the thinnest slices they cannot be determined by means of the microscope. Some of these rocks show perlitic or spherulitic structure, and such rocks were probably originally glassy (obsidians or pitchstones), but by lapse of time and processes of alteration have slowly passed into very finely crystal-line state. This change is called devitrification; it is common in glasses, as these are essentially unstable. A large number of the finer quartz-porphyries are also in some degree silicified of impregnated by quartz, chalcedony and opal, derived from the silica set free by decomposition (kaolinization) of the original feldspar. This re-deposited silica forms veins and patches of indefinite shape or may bodily replace a considerable area of the rock by metasomatic substitution. The opal is amorphous, the chalcedony finely crystalline and often arranged in spherulitic growths that yield an excellent black cross in polarized light. The microcrystalline ground-masses are those that can be resolved into their component minerals in thin slices by use of the microscope. They prove to consist essentially of quartz and feldspars, which are often in grains of quite irregular shape (microgranitic).
In other cases these two minerals are in graphic intergrowth, often forming radiate growths of spherulites consisting of fibers of extreme tenuity; this type is known as granophyric. There is another group in which the matrix contains small rounded or shapeless patches of quartz in which many rectangular feldspars are embedded; this structure is called micropoikilitic, and though often primary is sometimes developed by secondary changes that involve the deposit of new quartz in the ground-mass. As a whole those quartz-porphyries that have microcrystalline ground-masses are rocks of intrusive origin. Elvan is a name given locally to the quartz-porphyries that occur as dikes in Cornwall. In many of them the matrix contains scales of colorless muscovite or minute needles of blue tourmaline. Fluorite and kaolin appear also in these rocks, and the whole of these minerals are due to pneumatolytic action by vapors permeating the porphyry after it had consolidated but probably before it had entirely cooled. Many ancient rhyolitic quartz-porphyries show on their weathered surfaces numerous globular projections. They may be several inches in diameter, and vary from this size down to a minute fraction of an inch. When struck with a hammer they may detach readily from the matrix as if their margins were defined by a fissure. If they are broken across their inner portions are often seen to be filled with secondary quartz, chalcedony or agate: some of them have a central cavity, often with deposits of quartz crystals; they also frequently exhibit a succession of rounded cracks or dark lines occupied by secondary products. Rocks having these structures are common in north Wales and Cumberland; they occur also in Jersey, the Vosges and Hungary. It has been proposed to call them pyromerides.
Much discussion has taken place regarding the origin of these spheroids, but it is generally admitted that most of them were originally spherulites, and that they have suffered extensive changes through decomposition and silicification. Many of the older quartz-porphyries that occur in Paleozoic and Pre-Cambrian rocks have been affected by earth movements, and have experienced crushing and shearing. In this way they become schistose, and from their feldspar minute plates of sericitic white mica are developed, giving the rock in some cases very much of the appearance of mica-schists. If there have been no phenocrysts in the original rock, very perfect mica-schists may be produced, which can hardly be distinguished from sedimentary schists, though chemically somewhat different on account of the larger amounts of alkalis igneous rocks contain. When phenocrysts were present they often remain, though rounded and dragged apart while the matrix flows around them. The glassy or felsitic enclosures in the quartz are then very suggestive of an igneous origin for the rock. Such porphyry-schists have been called porphyroids or porphyroid-schists, and in the United States the name aporhyolite has been used for them. They are well known in some parts of the Alps, Westphalia, Charnwood (England), and Pennsylvania. The halleflintas of Sweden are also in part acid igneous rocks with a well-banded schistose or granulitic texture. The quartz-porphyries are distinguished from the rhyolites by being either intrusive rocks or Palaeozoic lavas. All Tertiary acid lavas are included under rhyolites. The intrusive quartz-porphyries are equally well described as granite-porphyries.
QMRZircon (/ˈzɜrkɒn/[5][6] or /ˈzɜrkən/;[7] including hyacinth or yellow zircon) is a mineral belonging to the group of nesosilicates. Its chemical name is zirconium silicate and its corresponding chemical formula is ZrSiO4. A common empirical formula showing some of the range of substitution in zircon is (Zr1–y, REEy)(SiO4)1–x(OH)4x–y. Zircon forms in silicate melts with large proportions of high field strength incompatible elements. For example, hafnium is almost always present in quantities ranging from 1 to 4%. The crystal structure of zircon is tetragonal crystal system. The natural color of zircon varies between colorless, yellow-golden, red, brown, blue, and green. Colorless specimens that show gem quality are a popular substitute for diamond and are also known as "Matura diamond".
Zircon is mainly consumed as an opacifier, and has been known to be used in the decorative ceramics industry.[11] It is also the principal precursor not only to metallic zirconium, although this application is small, but also to all compounds of zirconium including zirconium dioxide (ZrO2), one of the most refractory materials known.
Zircon has played an important role during the evolution of radiometric dating. Zircons contain trace amounts of uranium and thorium (from 10 ppm up to 1 wt%) and can be dated using several modern analytical techniques. Because zircons can survive geologic processes like erosion, transport, even high-grade metamorphism, they contain a rich and varied record of geological processes. Currently, zircons are typically dated by uranium-lead (U-Pb), fission-track, and U+Th/He techniques.
Zircons from Jack Hills in the Narryer Gneiss Terrane, Yilgarn Craton, Western Australia, have yielded U-Pb ages up to 4.404 billion years,[13] interpreted to be the age of crystallization, making them the oldest minerals so far dated on Earth. In addition, the oxygen isotopic compositions of some of these zircons have been interpreted to indicate that more than 4.4 billion years ago there was already water on the surface of the Earth.[13][14] This interpretation is supported by additional trace element data,[15][16] but is also the subject of debate.[17][18] In 2015, "remains of biotic life" were found in 4.1 billion-year-old rocks in the Jack Hills of Western Australia.[19][20] According to one of the researchers, "If life arose relatively quickly on Earth ... then it could be common in the universe."[19]
QMRIn the majority of silicates, the Si atom shows tetrahedral coordination, with 4 oxygen atoms surrounding a central Si atom. The most common example is seen in the quartz crystalline form of silica SiO2. In each of the most thermodynamically stable crystalline forms of silica, on average, all 4 of the vertices (or oxygen atoms) of the SiO4 tetrahedra are shared with others, yielding the net chemical formula: SiO2.
Relation between refractive index and density for some SiO2 forms.[14]
For example, in the unit cell of α-quartz, the central tetrahedron shares all 4 of its corner O atoms, the 2 face-centered tetrahedra share 2 of their corner O atoms, and the 4 edge-centered tetrahedra share just one of their O atoms with other SiO4 tetrahedra. This leaves a net average of 12 out of 24 total vertices for that portion of the 7 SiO4 tetrahedra that are considered to be a part of the unit cell for silica (see 3-D Unit Cell).
SiO2 has a number of distinct crystalline forms (polymorphs) in addition to amorphous forms. With the exception of stishovite and fibrous silica, all of the crystalline forms involve tetrahedral SiO4 units linked together by shared vertices in different arrangements. Silicon–oxygen bond lengths vary between the different crystal forms, for example in α-quartz the bond length is 161 pm, whereas in α-tridymite it is in the range 154–171 pm. The Si-O-Si angle also varies between a low value of 140° in α-tridymite, up to 180° in β-tridymite. In α-quartz the Si-O-Si angle is 144°.[15]
Fibrous silica has a structure similar to that of SiS2 with chains of edge-sharing SiO4 tetrahedra. Stishovite, the higher-pressure form, in contrast has a rutile-like structure where silicon is 6-coordinate. The density of stishovite is 4.287 g/cm3, which compares to α-quartz, the densest of the low-pressure forms, which has a density of 2.648 g/cm3.[8] The difference in density can be ascribed to the increase in coordination as the six shortest Si-O bond lengths in stishovite (four Si-O bond lengths of 176 pm and two others of 181 pm) are greater than the Si-O bond length (161 pm) in α-quartz.[16] The change in the coordination increases the ionicity of the Si-O bond.[17] But more important is the observation that any deviations from these standard parameters constitute microstructural differences or variations, which represent an approach to an amorphous, vitreous or glassy solid.
The only stable form under normal conditions is α-quartz and this is the form in which crystalline silicon dioxide is usually encountered. In nature impurities in crystalline α-quartz can give rise to colors (see list). The high temperature minerals, cristobalite and tridymite, have both a lower density and index of refraction than quartz. Since the composition is identical, the reason for the discrepancies must be in the increased spacing in the high temperature minerals. As is common with many substances, the higher the temperature the farther apart the atoms due to the increased vibration energy.
The transformation from α-quartz to beta-quartz takes place abruptly at 573 C. Since the transformation is accompanied by a significant change in volume it can easily induce fracturing of ceramics or rocks passing through this temperature limit.
The high-pressure minerals, seifertite, stishovite, and coesite, on the other hand, have a higher density and index of refraction when compared to quartz. This is probably due to the intense compression of the atoms that must occur during their formation, resulting in a more condensed structure.
Faujasite silica is another form of crystalline silica. It is obtained by dealumination of a low-sodium, ultra-stable Y zeolite with a combined acid and thermal treatment. The resulting product contains over 99% silica, has high crystallinity and high surface area (over 800 m2/g). Faujasite-silica has very high thermal and acid stability. For example, it maintains a high degree of long-range molecular order (or crystallinity) even after boiling in concentrated hydrochloric acid.[18]
Molten silica exhibits several peculiar physical characteristics that are similar to the ones observed in liquid water: negative temperature expansion, density maximum (at temperatures ~5000 °C), and a heat capacity minimum.[19] Its density decreases from 2.08 g/cm3 at 1950 °C to 2.03 g/cm3 at 2200 °C.[20] When molecular silicon monoxide, SiO, is condensed in an argon matrix cooled with helium along with oxygen atoms generated by microwave discharge, molecular SiO2 is produced with a linear structure. Dimeric silicon dioxide, (SiO2)2 has been prepared by reacting O2 with matrix isolated dimeric silicon monoxide, (Si2O2). In dimeric silicon dioxide there are two oxygen atoms bridging between the silicon atoms with an Si-O-Si angle of 94° and bond length of 164.6 pm and the terminal Si-O bond length is 150.2 pm. The Si-O bond length is 148.3 pm, which compares with the length of 161 pm in α-quartz. The bond energy is estimated at 621.7 kJ/mol.[21]
Fused quartz[edit]
Main article: Fused quartz
When molten silicon dioxide SiO2 is rapidly cooled, it does not crystallize but solidifies as a glass. The geometry of the silicon and oxygen centers in glass is similar to that in quartz and most other crystalline forms of the same composition, i.e., silicon is surrounded by a regular tetrahedra of oxygen centers. The difference between the glass and the crystalline forms arise from the connectivity of these tetrahedral units. Although there is no long range periodicity in the glassy network there remains significant ordering at length scales well beyond the SiO bond length. One example of this ordering is found in the preference of the network to form rings of 6-tetrahedra.[22]
The glass transition temperature of pure SiO2 is about 1475 K.[23]
Chemical reactions[edit]
Manufactured silica fume at maximum surface area of 380 m2/g
Silica is converted to silicon by reduction with carbon.
Fluorine reacts with silicon dioxide to form SiF4 and O2 whereas the other halogen gases (Cl2, Br2, I2) are essentially unreactive.[8]
Silicon dioxide is attacked by hydrofluoric acid (HF) to produce hexafluorosilicic acid:[15]
SiO2 + 6 HF → H2SiF6 + 2 H2O.
HF is used to remove or pattern silicon dioxide in the semiconductor industry.
Silicon dioxide dissolves in hot concentrated alkali or fused hydroxide, as described in this idealized equation:[8]
SiO2 + 2 NaOH → Na2SiO3 + H2O.
Silicon dioxide reacts with basic metal oxides (e.g. sodium oxide, potassium oxide, lead(II) oxide, zinc oxide, or mixtures of oxides forming silicates and glasses as the Si-O-Si bonds in silica are broken successively).[15] As an example the reaction of sodium oxide and SiO2 can produce sodium orthosilicate, sodium silicate, and glasses, dependent on the proportions of reactants:[8]
2 Na2O + SiO2 → Na4SiO4;
Na2O + SiO2 → Na2SiO3;
(0.25–0.8)Na2O + SiO2 → glass.
Examples of such glasses have commercial significance, e.g. soda-lime glass, borosilicate glass, lead glass. In these glasses, silica is termed the network former or lattice former.[15]
Bundle of optical fibers composed of high purity silica.
Silicon dioxide reacts in heated reflux under dinitrogen with ethylene glycol and an alkali metal base to produce highly reactive, pentacoordinate silicates which provide access to a wide variety of new silicon compounds.[24] The silicates are essentially insoluble in all polar solvent except methanol.
Silicon dioxide reacts with elemental silicon at high temperatures to produce SiO:[15]
SiO2 + Si → 2 SiO
Solubility in water[edit]
The solubility of silicon dioxide in water strongly depends on its crystalline form and is 3–4 times higher for silica than quartz; as a function of temperature, it peaks at about 340 °C.[25] This property is used to grow single crystals of quartz in a hydrothermal process where natural quartz is dissolved in superheated water in a pressure vessel that is cooler at the top. Crystals of 0.5–1 kg can be grown over a period of 1–2 months.[15] These crystals are a source of very pure quartz for use in electronic applications.[8]
Occurrence[edit]
Biology[edit]
Even though it is poorly soluble, silica occurs widely in many plants. Plant materials with high silica phytolith content appear to be of importance to grazing animals, from chewing insects to ungulates. Studies have shown that it accelerates tooth wear, and high levels of silica in plants frequently eaten by herbivores may have developed as a defense mechanism against predation.[26][27]
It is also the primary component of rice husk ash, which is used, for example, in filtration and cement manufacturing.
Silicification in and by cells has been common in the biological world for well over a billion years. In the modern world it occurs in bacteria, single-celled organisms, plants, and animals (invertebrates and vertebrates). Prominent examples include:
Tests or frustules (i.e. shells) of diatoms, Radiolaria and testate amoebae.[6]
Silica phytoliths in the cells of many plants, including Equisetaceae, practically all grasses, and a wide range of dicotyledons.
The spicules forming the skeleton of many sponges.
Crystalline minerals formed in the physiological environment often show exceptional physical properties (e.g., strength, hardness, fracture toughness) and tend to form hierarchical structures that exhibit microstructural order over a range of scales. The minerals are crystallized from an environment that is undersaturated with respect to silicon, and under conditions of neutral pH and low temperature (0–40 °C).
Formation of the mineral may occur either within the cell wall of an organism (such as with phytoliths), or outside the cell wall, as typically happens with tests. Specific biochemical reactions exist for mineral deposition. Such reactions include those that involve lipids, proteins, and carbohydrates.
It is unclear in what ways silica is important in the nutrition of animals. This field of research is challenging because silica is ubiquitous and in most circumstances dissolves in trace quantities only. All the same it certainly does occur in the living body, leaving us with the problem that it is hard to create proper silica-free controls for purposes of research. This makes it difficult to be sure when the silica present has had operative beneficial effects, and when its presence is coincidental, or even harmful. The current consensus is that it certainly seems important in the growth, strength, and management of many connective tissues. This is true not only for hard connective tissues such as bone and tooth but possibly in the biochemistry of the subcellular enzyme-containing structures as well.[28]
Health effects[edit]
Quartz sand (silica) as main raw material for commercial glass production
Silica ingested orally is essentially nontoxic, with an LD50 of 5000 mg/kg (5 g/kg).[7] On the other hand, inhaling finely divided crystalline silica dust can lead to silicosis, bronchitis, or cancer, as the dust becomes lodged in the lungs and continuously irritates them, reducing lung capacities.[29] Studies of workers with exposure to crystalline silica have shown 10-fold higher than expected rates of lupus and other systemic autoimmune diseases compared to expected rates in the general population.[30] Prior to new rules issued in 2013, OSHA allowed 100 µg per cubic meter of air. The new regulations reduce the amount to 50 µg/m3 down from 100 µg/m3. The exposure limit for the construction industry is also set at 50 µg/m3 down from 250 µg/m3.[31]
In the body crystalline silica particles do not dissolve over clinically relevant periods. Silica crystals inside the lungs can activate the NLRP3 inflammasome inside macrophages and dendritic cells and thereby result in processing of pro-Interleukin 1 beta into its mature form. Chronic exposure to silica may thereby account for some of its health hazards, as interleukin-1 is a highly pro-inflammatory cytokine in the immune system.[32][33][34] This effect can create an occupational hazard for people working with sandblasting equipment, products that contain powdered crystalline silica and so on. Children, asthmatics of any age, allergy sufferers, and the elderly (all of whom have reduced lung capacity) can be affected in much less time. Amorphous silica, such as fumed silica is not associated with development of silicosis, but may cause irreversible lung damage in some cases.[35] Laws restricting silica exposure with respect to the silicosis hazard specify that they are concerned only with silica that is both crystalline and dust-forming.
A study that followed subjects for 15 years found that higher levels of silica in water appeared to decrease the risk of dementia. The study found an association between an increase of 10 milligram-per-day of the intake of silica in drinking water with a decreased risk of dementia of 11%.[36]
Crystalline silica is used in hydraulic fracturing of formation which contain tight oil and shale gas, a use which presents a health hazard to workers. In 2013 OSHA announced tightened restrictions on the amount of crystalline silica which could be present and required "green completion" of fracked wells to reduce exposure.[31] Crystalline silica is an occupational hazard for those working with stone countertops, because the process of cutting and installing the countertops creates large amounts of airborne silica.[37]
Crystalline forms[edit]
SiO2, more so than almost any material, exists in many crystalline forms (called polymorphs).
Crystalline forms of SiO2[15]
Form Crystal symmetry
Pearson symbol, group No. ρ
g/cm3 Notes Structure
α-quartz rhombohedral (trigonal)
hP9, P3121 No.152[38] 2.648 Helical chains making individual single crystals optically active; α-quartz converts to β-quartz at 846 K A-quartz.png
β-quartz hexagonal
hP18, P6222, No. 180[39] 2.533 Closely related to α-quartz (with an Si-O-Si angle of 155°) and optically active; β-quartz converts to β-tridymite at 1140 K B-quartz.png
α-tridymite orthorhombic
oS24, C2221, No.20[40] 2.265 Metastable form under normal pressure A-tridymite.png
β-tridymite hexagonal
hP12, P63/mmc, No. 194[40] Closely related to α-tridymite; β-tridymite converts to β-cristobalite at 2010 K B-tridymite.png
α-cristobalite tetragonal
tP12, P41212, No. 92[41] 2.334 Metastable form under normal pressure A-cristobalite.png
β-cristobalite cubic
cF104, Fd3m, No.227[42] Closely related to α-cristobalite; melts at 1978 K B-cristobalite.png
keatite tetragonal
tP36, P41212, No. 92[43] 3.011 Si5O10, Si4O14, Si8O16 rings; synthesised from glassy silica and alkali at 600–900 K and 40–400 MPa Keatite.png
moganite monoclinic
mS46, C2/c, No.15[44] Si4O8 and Si6O12 rings Moganite.png
coesite monoclinic
mS48, C2/c, No.15[45] 2.911 Si4O8 and Si8O16 rings; 900 K and 3–3.5 GPa Coesite.png
stishovite Tetragonal
tP6, P42/mnm, No.136[46] 4.287 One of the densest (together with seifertite) polymorphs of silica; rutile-like with 6-fold coordinated Si; 7.5–8.5 GPa Stishovite.png
seifertite orthorhombic
oP, Pbcn[47] 4.294 One of the densest (together with stishovite) polymorphs of silica; is produced at pressures above 40 GPa.[48] SeifertiteStructure.png
melanophlogite cubic (cP*, P4232, No.208)[14] or tetragonal (P42/nbc)[49] 2.04 Si5O10, Si6O12 rings; mineral always found with hydrocarbons in interstitial spaces - a clathrasil[50] MelanophlogiteStucture.png
faujasite cubic
cF576, Fd3m, No.227[51] 1.92 Sodalite cages connected by hexagonal prisms; 12-membered ring pore opening; faujasite structure.[18] Faujasite structure.svg
fibrous
W-silica[8] orthorhombic
oI12, Ibam, No.72[52] 1.97 Like SiS2 consisting of edge sharing chains, melts at ~1700 K SiS2typeSilica.png
2D silica[53] hexagonal Sheet-like bilayer structure
Tridymite
From Wikipedia, the free encyclopedia
Tridymite
Tridymite tabulars - Ochtendung, Eifel, Germany.jpg
tabular tridymite crystals from Ochtendung, Eifel, Germany
General
Category Oxide mineral
Formula
(repeating unit) SiO2
Identification
Formula mass 60.08
Color Colorless, white
Crystal habit Platy – sheet forms
Crystal system several coexisting phases
Cleavage {0001} indistinct, {1010} imperfect
Fracture Brittle – conchoidal
Mohs scale hardness 7
Luster Vitreous
Streak white
Specific gravity 2.25–2.28
Optical properties Biaxial (+), 2V=40–86°
Refractive index nα=1.468–1.482 nβ=1.470–1.484 nγ=1.474–1.486
Birefringence δ < 0.004
Pleochroism Colorless
Other characteristics non-radioactive, non-magnetic; fluorescent, short UV=dark red
References [1]
Tridymite is a high-temperature polymorph of silica and usually occurs as minute tabular white or colorless pseudo-hexagonal crystals, or scales, in cavities in felsic volcanic rocks. Its chemical formula is SiO2. Tridymite was first described in 1868 and the type location is in Hidalgo, Mexico. The name is from the Greek tridymos for triplet as tridymite commonly occurs as twinned crystal trillings.[1]
Contents [hide]
1 Structure
2 Mars
3 See also
4 References
5 External links
Tridymite
From Wikipedia, the free encyclopedia
Tridymite
Tridymite tabulars - Ochtendung, Eifel, Germany.jpg
tabular tridymite crystals from Ochtendung, Eifel, Germany
General
Category Oxide mineral
Formula
(repeating unit) SiO2
Identification
Formula mass 60.08
Color Colorless, white
Crystal habit Platy – sheet forms
Crystal system several coexisting phases
Cleavage {0001} indistinct, {1010} imperfect
Fracture Brittle – conchoidal
Mohs scale hardness 7
Luster Vitreous
Streak white
Specific gravity 2.25–2.28
Optical properties Biaxial (+), 2V=40–86°
Refractive index nα=1.468–1.482 nβ=1.470–1.484 nγ=1.474–1.486
Birefringence δ < 0.004
Pleochroism Colorless
Other characteristics non-radioactive, non-magnetic; fluorescent, short UV=dark red
References [1]
Tridymite is a high-temperature polymorph of silica and usually occurs as minute tabular white or colorless pseudo-hexagonal crystals, or scales, in cavities in felsic volcanic rocks. Its chemical formula is SiO2. Tridymite was first described in 1868 and the type location is in Hidalgo, Mexico. The name is from the Greek tridymos for triplet as tridymite commonly occurs as twinned crystal trillings.[1]
Contents [hide]
1 Structure
2 Mars
3 See also
4 References
5 External links
Tridymite can occur in seven crystalline forms. Two of the most common at standard pressure are known as α and β. The α-tridymite phase is favored at elevated temperatures (>870 °C) and it converts to β-cristobalite at 1470 °C.[2][3] However, tridymite does usually not form from pure β-quartz, one needs to add trace amounts of certain compounds to achieve this.[4] Otherwise the β-quartz-tridymite transition is skipped and β-quartz transitions directly to cristobalite at 1050 °C without occurrence of the tridymite phase.
Crystal phases of tridymite[3]
Name Symmetry Space group T (°C)
HP (β) Hexagonal P63/mmc 460
LHP Hexagonal P6322 400
OC (α) Orthorhombic C2221 220
OS Orthorhombic 100–200
OP Orthorhombic P212121 155
MC Monoclinic Cc 22
MX Monoclinic C1 22
In the table, M, O, H, C, P, L and S stand for monoclinic, orthorhombic, hexagonal, centered, primitive, low (temperature) and superlattice. T indicates the temperature, at which the corresponding phase is relatively stable, though the interconversions between phases are complex and sample dependent, and all these forms can coexist at ambient conditions.[3] Mineralogy handbooks often arbitrarily assign tridymite to the triclinic crystal system, yet use hexagonal Miller indices because of the hexagonal crystal shape (see infobox image).[1]
In December 2015, the team behind NASA's Mars Science Laboratory announced the discovery of large amounts of tridymite in Marias pass on the slope of Aeolis Mons, popularly known as Mount Sharp, on the planet Mars.[5] This discovery was unexpected given the rarity of the mineral on Earth and the apparent lack of volcanic activity where it was discovered, and at the time of discovery no explanation for how it was formed was forthcoming. Its discovery was serendipitous: two teams, responsible for two different instruments on the Curiosity rover, both happened to report what in isolation were relatively uninteresting findings related to their instruments: the ChemCam team reported a region of high silica while the DAN team reported high neutron readings in what turned out to be the same area. Neither team would have been aware of the other's findings had it not been for a fortuitous Mars conjunction in July 2015, during which the various international teams took advantage of the downtime to meet in Paris and discuss their scientific findings. DAN's high neutron readings would normally have been interpreted as meaning the region was hydrogen-rich, and ChemCam's high-silica readings were not surprising given the ubiquity of silica-rich deposits on Mars, but taken together it was clear that further study of the region was needed. Following conjunction, NASA directed the Curiosity rover back to the area where the readings had been taken and discovered that large amounts of tridymite were present. How they were formed remains a mystery.[6]
QMRThe mineral cristobalite is a high-temperature polymorph of silica, meaning that it has the same chemical formula as quartz, SiO2, but a distinct crystal structure. Both quartz and cristobalite are polymorphs with all the members of the quartz group, which also include coesite, tridymite and stishovite. Cristobalite occurs as white octahedra or spherulites in acidic volcanic rocks and in converted diatomaceous deposits in the Monterey Formation of the US state of California and similar areas. Cristobalite is stable only above 1470 °C, but can crystallize and persist metastably at lower temperatures. It is named after Cerro San Cristóbal in Pachuca Municipality, Hidalgo, Mexico.
The persistence of cristobalite outside of its thermodynamic stability range occurs because the transition from cristobalite to quartz or tridymite is "reconstructive", requiring the breaking up and reforming of the silica framework. These frameworks are composed of SiO4 tetrahedra in which every oxygen atom is shared with a neighbouring tetrahedron, so that the chemical formula of silica is SiO2. The breaking of these bonds required to convert cristobalite to tridymite and quartz requires considerable activation energy and may not happen on a human time frame. Framework silicates are also known as tectosilicates.
The mineral cristobalite is a high-temperature polymorph of silica, meaning that it has the same chemical formula as quartz, SiO2, but a distinct crystal structure. Both quartz and cristobalite are polymorphs with all the members of the quartz group, which also include coesite, tridymite and stishovite. Cristobalite occurs as white octahedra or spherulites in acidic volcanic rocks and in converted diatomaceous deposits in the Monterey Formation of the US state of California and similar areas. Cristobalite is stable only above 1470 °C, but can crystallize and persist metastably at lower temperatures. It is named after Cerro San Cristóbal in Pachuca Municipality, Hidalgo, Mexico.
The persistence of cristobalite outside of its thermodynamic stability range occurs because the transition from cristobalite to quartz or tridymite is "reconstructive", requiring the breaking up and reforming of the silica framework. These frameworks are composed of SiO4 tetrahedra in which every oxygen atom is shared with a neighbouring tetrahedron, so that the chemical formula of silica is SiO2. The breaking of these bonds required to convert cristobalite to tridymite and quartz requires considerable activation energy and may not happen on a human time frame. Framework silicates are also known as tectosilicates.
There is more than one form of the cristobalite framework. At high temperatures, the structure is cubic, Fd3m, No.227, Pearson symbol cF104.[5] A tetragonal form of cristobalite (P41212, No. 92, Pearson symbol tP12[6]) occurs on cooling below ca. 250 °C at ambient pressure, and is related to the cubic form by a static tilting of the silica tetrahedra in the framework. This transition is variously called the low-high or \alpha -\beta transition. It may be termed "displacive"; i.e., it is not generally possible to prevent the cubic β-form from becoming tetragonal by rapid cooling. Under rare circumstances the cubic form may be preserved if the crystal grain is pinned in a matrix that does not allow for the considerable spontaneous strain that is involved in the transition, which causes a change in shape of the crystal. This transition is highly discontinuous. The exact transition temperature depends on the crystallinity of the cristobalite sample, which itself depends on factors such as how long it has been annealed at a particular temperature.
The cubic β phase consists of dynamically disordered silica tetrahedra. The tetrahedra remain fairly regular and are displaced from their ideal static orientations due to the action of a class of low-frequency phonons called rigid unit modes. It is the "freezing" of one of these rigid unit modes that is the soft mode for the α–β transition.
In the α–β phase transition only one of the three degenerate cubic crystallographic axes retains a fourfold rotational axis in the tetragonal form. The choice of axis is arbitrary, so that various twins can form within the same grain. These different twin orientations coupled with the discontinuous nature of the transition can cause considerable mechanical damage to materials in which cristobalite is present and that pass repeatedly through the transition temperature, such as refractory bricks.
When devitrifying silica, cristobalite is usually the first phase to form, even when well outside of its thermodynamic stability range. The dynamically disordered nature of the β-phase is partly responsible for the low enthalpy of fusion of silica.
The micrometre-scale spheres that make up precious opal exhibit some x-ray diffraction patterns that are similar to that of cristobalite, but lack any long-range order so they are not considered true cristobalite. In addition, the presence of structural water in opal makes it doubtful that opal consists of cristobalite.
QMRIn crystallography, the tetragonal crystal system is one of the 7 lattice point groups. Tetragonal crystal lattices result from stretching a cubic lattice along one of its lattice vectors, so that the cube becomes a rectangular prism with a square base (a by a) and height (c, which is different from a).
There are two tetragonal crystal structure types. Bravais lattices: the simple tetragonal (from stretching the simple-cubic lattice) and the centered tetragonal (from stretching either the face-centered or the body-centered cubic lattice). One might suppose stretching face-centered cubic would result in face-centered tetragonal, but face-centered tetragonal is equivalent to body-centered tetragonal, BCT (with a smaller lattice spacing). BCT is considered more fundamental, so that is the standard terminology.[1]
Crystal classes[edit]
The point groups that fall under this crystal system are listed below, followed by their representations in international notation, Schoenflies notation, orbifold notation, Coxeter notation and mineral examples.[2][3]
# Point group Example Space groups
Class Intl Schoen. Orb. Cox.
75–80 Tetragonal pyramidal 4 C4 44 [4]+ pinnoite,
piypite P4, P41, P42, P43
I4, I41
81–82 Tetragonal disphenoidal 4 S4 2× [2+,4+] cahnite, tugtupite P4
I4
83–88 Tetragonal dipyramidal 4/m C4h 4* [2,4+] scheelite, wulfenite, leucite P4/m, P42/m, P4/n, P42/n
I4/m, I41/a
89–98 Tetragonal trapezohedral 422 D4 224 [2,4]+ cristobalite, wardite P422, P4212, P4122, P41212, P4222, P42212, P4322, P43212
I422, I4122
99–110 Ditetragonal pyramidal 4mm C4v *44 [4] diaboleite P4mm, P4bm, P42cm, P42nm, P4cc, P4nc, P42mc, P42bc
I4mm, I4cm, I41md, I41cd
111–122 Tetragonal scalenohedral 42m D2d 2*2 [2+,4] chalcopyrite, stannite P42m, P42c, P421m, P421c, P4m2, P4c2, P4b2, P4n2
I4m2, I4c2, I42m, I42d
123–142 Ditetragonal dipyramidal 4/mmm D4h *224 [2,4] rutile, pyrolusite, zircon P4/mmm, P4/mcc, P4/nbm, P4/nnc, P4/mbm, P4/mnc, P4/nmm, P4/ncc, P42/mmc, P42/mcm, P42/nbc, P42/nnm, P42/mbc, P42/mnm, P42/nmc, P42/ncm
I4/mmm, I4/mcm, I41/amd, I41/acd
QMRSilicon is a chemical element with symbol Si and atomic number 14. It is a tetravalent metalloid, more reactive than germanium, the metalloid directly below it in the table. Controversy about silicon's character dates to its discovery. It was first prepared and characterized in pure form in 1823. In 1808, it was given the name silicium (from Latin: silex, hard stone or flint), with an -ium word-ending to suggest a metal, a name which the element retains in several non-English languages. However, its final English name, first suggested in 1817, reflects the more physically similar elements carbon and boron.
Silicon is the eighth most common element in the universe by mass, but very rarely occurs as the pure free element in nature. 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.[9]
Most silicon is used commercially without being separated, and indeed often with little processing of compounds from nature. These include direct industrial building-use of clays, silica sand and stone. Silicate goes into Portland cement for mortar and stucco, and mixed with silica sand and gravel to make concrete. Silicates are also in whiteware ceramics such as porcelain, and in traditional quartz-based soda-lime glass and many other specialty glasses. More modern silicon compounds such as silicon carbide form abrasives and 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. Although most free silicon is used in the steel refining, aluminium-casting, and fine chemical industries (often to make fumed silica), the relatively small portion of very highly purified silicon that is used in semiconductor electronics (< 10%) is perhaps even more critical. Because of wide use of silicon in integrated circuits, the basis of most computers, a great deal of modern technology depends on it.
Silicon is an essential element in biology, although only tiny traces of it appear to be required by animals.[10] However, various sea sponges as well as microorganisms like diatoms and radiolaria secrete skeletal structures made of silica. Silica is often deposited in 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.[11]
Characteristics
Physical
Silicon crystallizes in a diamond cubic crystal structure
Further information: Monocrystalline silicon
Silicon is a solid at room temperature, with relatively high melting and boiling points of 1414 and 3265 °C, respectively. Like water, it has a greater density in a liquid state than in a solid state, and so, like water but unlike most substances, it does not contract when it freezes, but expands. With a relatively high thermal conductivity of 149 W·m−1·K−1, silicon conducts heat well.
In its crystalline form, pure silicon has a gray color and a metallic luster. Like germanium, silicon is rather strong,[vague] very brittle, and prone to chipping. Silicon, like carbon and germanium, crystallizes in a diamond cubic crystal structure, with a lattice spacing of 0.5430710 nm (5.430710 Å).[12]
The outer electron orbital of silicon, like that of carbon, has four valence electrons. The 1s, 2s, 2p and 3s subshells are completely filled while the 3p subshell contains two electrons out of a possible six.
Silicon is a semiconductor. It has a negative temperature coefficient of resistance, since the number of free charge carriers increases with temperature. The electrical resistance of single crystal silicon significantly changes under the application of mechanical stress due to the piezoresistive effect.[13]
Chemical
Silicon powder
Silicon is a metalloid, readily either donating or sharing its four outer electrons and it typically forms four bonds. Like carbon, its four bonding electrons give it opportunities to combine with many other elements or compounds to form a wide range of compounds. Unlike carbon, it can accept additional electrons and form five or six bonds in a sometimes more labile silicate form. Tetra-valent silicon is relatively inert, but still reacts with halogens and dilute alkalis, but most acids (except for some hyper-reactive combinations of nitric acid and hydrofluoric acid) have no known effect on it.
Isotopes
Main article: isotopes of silicon
Naturally occurring silicon is composed of three stable isotopes, silicon-28, silicon-29, and silicon-30, with silicon-28 being the most abundant (92% natural abundance).[14] Out of these, only silicon-29 is of use in NMR and EPR spectroscopy.[15] Twenty radioisotopes have been characterized, with the most stable being silicon-32 with a half-life of 170 years, and silicon-31 with a half-life of 157.3 minutes.[14] 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.[14] Silicon does not have any known nuclear isomers.[14]
The isotopes of silicon range in mass number from 22 to 44.[14] The most common decay mode of six isotopes with mass numbers lower than the most abundant stable isotope, silicon-28, is β+, primarily forming aluminium isotopes (13 protons) as decay products.[14] The most common decay mode(s) for 16 isotopes with mass numbers higher than silicon-28 is β−, primarily forming phosphorus isotopes (15 protons) as decay products.[14]
History
Attention was first drawn to silica as the possible oxide of a fundamental chemical element by Antoine Lavoisier, in 1787.[16] 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 was a metal.[17] 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.[18] 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.[19] In 1823, Berzelius prepared amorphous silicon using approximately the same method as Gay-Lussac (potassium metal and potassium fluorosilicate), but purifying the product to a brown powder by repeatedly washing it.[20] As a result, he is usually given credit for the element's discovery.[21][22]
Silicon in its more common crystalline form was not prepared until 31 years later, by Deville.[23][24] By electrolyzing impure sodium-aluminium chloride containing approximately 10% silicon, he was able to obtain a slightly impure allotrope of silicon in 1854.[25] Later, more cost-effective methods have been developed to isolate silicon in several allotrope forms, the most recent being silicene.
Because silicon is an important element in semiconductors and high-technology devices, many places in the world bear its name. For example, Silicon Valley in California, bears the element's name since it is the base for a number of computer technology-related industries. Other geographic locations with connections to the industry have since been named after silicon as well. Examples include Silicon Forest in Oregon, Silicon Hills in Austin, Texas, Silicon Slopes in Salt Lake City, Utah, Silicon Saxony in Germany, Silicon Valley in India, Silicon Border in Mexicali, Mexico, Silicon Fen in Cambridge, England, Silicon Roundabout in London, Silicon Glen in Scotland, and Silicon Gorge in Bristol, England.
Occurrence
Quartz crystal cluster from Tibet. The naturally occurring mineral is a network solid with the formula SiO2.
See also: Silicate minerals
Measured by mass, silicon makes up 27.7% of the Earth's crust and is the second most abundant element in the crust, with only oxygen having a greater abundance.[26] Silicon is usually found in the form of complex silicate minerals, and less often as silicon dioxide (silica, a major component of common sand). Pure silicon crystals are very rarely found in nature.
The silicate minerals—various minerals containing silicon, oxygen and reactive metals—account for 90% of the mass of the Earth's crust. This is due to the fact that at the high temperatures characteristic of the formation of the inner solar system, silicon and oxygen readily combine chemically, forming network solids of silicon and oxygen in compounds of very low volatility. Since oxygen and silicon were the most common non-gaseous and non-metallic elements in the debris from supernova dust which formed the protoplanetary disk in the formation and evolution of the Solar System, they formed many complex silicates which accreted into larger rocky planetesimals that formed the terrestrial planets. Here, the reduced silicate mineral matrix entrapped the metals reactive enough to be oxidized (aluminium, calcium, sodium, potassium and magnesium). After loss of volatile gases, as well as carbon and sulfur via reaction with hydrogen, this silicate mixture of elements formed most of the Earth's crust.
These silicates were of relatively low density with respect to iron, nickel, and other metals non-reactive to oxygen and thus a residuum of uncombined metallic iron and nickel sank to the planet's core, leaving a thick mantle between core and crust, consisting mostly of magnesium and iron silicates. These are thought to be mostly silicate perovskites, followed in abundance by the magnesium/iron oxide ferropericlase.[27]
Examples of silicate minerals in the crust include those in the pyroxene, amphibole, mica, and feldspar groups. These minerals occur in clay and various types of rock such as granite and sandstone. In the crust, silica occurs in minerals consisting of very pure silicon dioxide in different crystalline forms of quartz and opal. The crystals have the empirical formula of silicon dioxide, but do not consist of separate silicon dioxide molecules in the manner of solid carbon dioxide. Rather, silica is structurally a network solid consisting of silicon and oxygen in three-dimensional crystals, like diamond. Less pure silica forms the natural glass obsidian. Biogenic silica occurs in the structure of diatoms, radiolaria and siliceous sponges.
Silicon is also a principal component of many meteorites, and is a component of tektites, a silicate mineral of possibly lunar origin, or (if Earth-derived) which has been subjected to unusual temperatures and pressures, possibly from meteorite strike.
Production
Alloys
Ferrosilicon alloy
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).[28] Ferrosilicon is primarily used by the steel industry (see below).
Aluminium-silicon alloys (called silumin alloys) are heavily used in the aluminium alloy casting industry, where silicon is the single most important additive to aluminium to improve its casting properties. Since cast aluminium is widely used in the automobile industry, this use of silicon is thus the single largest industrial use (about 55% of the total) of "metallurgical grade" pure silicon (as this purified silicon is added to pure aluminium, whereas ferrosilicon is never purified before being added to steel).[29]
Metallurgical grade
Elemental silicon not alloyed with significant quantities of other elements, and usually > 95%, is often referred to loosely as silicon metal. It makes up about 20% of the world total elemental silicon production, with less than 1 to 2% of total elemental silicon (5–10% of metallurgical grade silicon) ever purified to higher grades for use in electronics. Metallurgical grade 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 1,900 °C (3,450 °F), the carbon in the aforementioned materials and the silicon undergo the chemical reaction:
SiO2 + 2 C → Si + 2 CO
Liquid silicon collects in the bottom of the furnace, which is then drained and cooled. The silicon produced in this manner is called metallurgical grade silicon and is at least 98% pure. Using this method, silicon carbide (SiC) may also form from an excess of carbon in one or both of the following ways:
SiO2 + C → SiO + CO
SiO + 2 C → SiC + CO
However, provided the concentration of SiO2 is kept high, the silicon carbide can be eliminated by the chemical reaction:
2 SiC + SiO2 → 3 Si + 2 CO
As noted above, metallurgical grade silicon "metal" has its primary use in the aluminium casting industry to make aluminium-silicon alloy parts. The remainder (about 45%) is used by the chemical industry, where it is primarily employed to make fumed silica, with the rest used in production of other fine chemicals such as silanes and some types of silicones.[30]
As of September 2008, metallurgical grade silicon costs about US$1.45 per pound ($3.20/kg),[31] up from $0.77 per pound ($1.70/kg) in 2005.[32]
Polysilicon
Main article: Polycrystalline silicon
Today's purification processes involve the conversion of silicon into volatile liquids, such as trichlorosilane (HSiCl3) and silicon tetrachloride (SiCl4) or into the gaseous silane (SiH4). These compounds are then separated by a distillation and transformed into high-purity silicon, either by a redox reaction or by chemical decomposition at high temperatures.
In the late 1950s, the American chemical company DuPont patented a method for the production of 99.99% pure silicon, using the metal zinc as a reductant to transform redistilled silicon tetrachloride into high-purity silicon by a vapor phase reaction at 900 °C. This technique, however, was plagued with practical problems, as the byproduct zinc chloride (ZnCl2) solidified and clogged lines, and was eventually abandoned in favor of more sophisticated processes.
Siemens process and alternatives
The best known technique is the so-called Siemens process. This technique does not require a reductant such as zinc, as it grows high-purity silicon crystallites directly on the surface of (pre-existing) pure silicon seed rods by a chemical decomposition that takes place when the gasous trichlorosilane is blown over the rod's surface at 1150 °C. A common name for this type of technique is chemical vapor deposition (CVD) and produces high-purity polycrystalline silicon, also known as polysilicon. While the conventional Siemens process produces electronic grade polysilicon at typically 9N–11N purities, that is, it contains impurity levels of less than one part per billion (ppb), the modified Siemens process is a dedicated process-route for the production of solar grade silicon (SoG-Si) with purities of 6N (99.9999%) and less energy demand.[34][35][36]
A more recent alternative for the production of polysilicon is the fluidized bed reactor (FBR) manufacturing technology. Compared to the traditional Siemens process, FBR features a number of advantages that lead to cheaper polysilicon demanded by the fast-growing photovoltaic industry. Contrary to Siemens' batch process, FBR runs continuously, wasting fewer resources and requires less setup and downtime. It uses about 10 percent of the electricity consumed by a conventional rod reactor in the established Siemens process, as it does not waste energy by placing heated gas and silicon in contact with cold surfaces. In the FBR, silane (SiH4) is injected into the reactor from below and forms a fluidized bed together with the silicon seed particles that are fed from above. The gaseous silane then decomposes and deposits silicon on the seed particles. When the particles have grown to larger granules, they eventually sink to the bottom of the reactor where they are continuously withdrawn from the process.
The FBR manufacturing technology outputs polysilicon at 6N to 9N, a purity still higher than the 5N to 6N of upgraded metallurgical silicon (UMG-Si), a third technology used by the photovoltaic industry, that dispenses altogether with chemical purification, using metallurgical techniques instead. Currently most silicon for the photovoltaic market is produced by the Siemens process and only about 10 percent by the FBR technology, while UMG-Si accounts for about 2 percent. By 2020, however, IHS Technology predicts that market shares for FBR technology and UMG-Si will grow to 16.7 and 5.4 percent, respectively.[37]
The company REC is one of the leading producers of silane and polysilicon using FBR technology. The three-step chemical reaction involves (last step occurs inside the FB-reactor): (1.) 3 SiCl4 + Si + 2 H2 → 4 HSiCl3, followed by (2.) 4 HSiCl3 → 3 SiCl4 + SiH4, and (3.) SiH4 → Si + 2 H2.[38] Other precursors such as tribromosilane had been used by other companies as well.
Electronic grade
Monocrystalline silicon ingot grown by the Czochralski process
Main article: Monocrystalline silicon
The use of silicon in semiconductor devices demands a much greater purity than afforded by metallurgical grade silicon. Very pure silicon (>99.9%) can be extracted directly from solid silica or other silicon compounds by molten salt electrolysis.[39][40] This method, known as early as 1854[41] (see also FFC Cambridge process), has the potential to directly produce solar-grade silicon without any carbon dioxide emission at much lower energy consumption.
Solar grade silicon cannot be used for microelectronics. To properly control the quantum mechanical properties, the purity of the silicon must be very high. Bulk silicon wafers used at the beginning of the integrated circuit making process must first be refined to a purity of 99.9999999% often referred to as "9N" for "9 nines", a process which requires repeated applications of refining technology.
The majority of silicon crystals grown for device production are produced by the Czochralski process, (Cz-Si) It was the cheapest method available. However, single crystals grown by the Czochralski process contain impurities because the crucible containing the melt often dissolves. Historically, a number of methods have been used to produce ultra-high-purity silicon.
Early purification techniques
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.[42]
Compounds
See also: Category:Silicon compounds.
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