# Silicon dioxide

Last updated
Silicon dioxide
Names
IUPAC name
Silicon dioxide
Other names
Quartz

Silica
Silicic oxide
Silicon(IV) oxide
Crystalline silica
Pure Silica
Silicea

Silica sand

## Contents

Identifiers
•
ChEBI
•
ChemSpider
•
ECHA InfoCard 100.028.678
EC Number
• 231-545-4
E number E551 (acidity regulators, ...)
200274
KEGG
•
MeSH
PubChem CID
RTECS number
• VV7565000
UNII
•
Properties
SiO2
Molar mass 60.08 g/mol
AppearanceTransparent solid (Amorphous) White/Whitish Yellow (Powder/Sand)
Density 2.648 (α-quartz), 2.196 (amorphous) g·cm−3 [1]
Melting point 1,713 °C (3,115 °F; 1,986 K)(amorphous) [1] (p4.88) to
Boiling point 2,950 °C (5,340 °F; 3,220 K) [1]
29.6·10−6 cm3/mol
Thermal conductivity 12 (|| c-axis), 6.8 (⊥ c-axis), 1.4 (am.) W/(m⋅K) [1] (p12.213)
1.544 (o), 1.553 (e) [1] (p4.143)
Hazards
NFPA 704 (fire diamond)
NIOSH (US health exposure limits):
PEL (Permissible)
TWA 20 mppcf (80 mg/m3/%SiO2) (amorphous) [2]
REL (Recommended)
TWA 6 mg/m3 (amorphous) [2]
Ca TWA 0.05 mg/m3 [3]
IDLH (Immediate danger)
3000 mg/m3 (amorphous) [2]
Ca [25 mg/m3 (cristobalite, tridymite); 50 mg/m3 (quartz)] [3]
Related compounds
Related diones
Carbon dioxide
Related compounds
Silicon monoxide
Thermochemistry
42 J·mol−1·K−1 [4]
−911 kJ·mol−1 [4]
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
(what is   ?)
Infobox references

Silicon dioxide, also known as silica, is an oxide of silicon with the chemical formula Si O 2, most commonly found in nature as quartz and in various living organisms. [5] [6] In many parts of the world, silica is the major constituent of sand. Silica is one of the most complex and most abundant families of materials, existing as a compound of several minerals and as synthetic product. Notable examples include fused quartz, fumed silica, silica gel, and aerogels. It is used in structural materials, microelectronics (as an electrical insulator), and as components in the food and pharmaceutical industries.

Inhaling finely divided crystalline silica is toxic and can lead to severe inflammation of the lung tissue, silicosis, bronchitis, lung cancer, and systemic autoimmune diseases, such as lupus and rheumatoid arthritis. Inhalation of amorphous silicon dioxide, in high doses, leads to non-permanent short-term inflammation, where all effects heal. [7]

## Structure

In the majority of silicates, the silicon atom shows tetrahedral coordination, with four oxygen atoms surrounding a central Si atom. The most common example is seen in the quartz polymorphs. It is a 3 dimensional network solid in which each silicon atom is covalently bonded in a tetrahedral manner to 4 oxygen atoms.

For example, in the unit cell of α-quartz, the central tetrahedron shares all four of its corner O atoms, the two face-centered tetrahedra share two of their corner O atoms, and the four 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 seven 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. Silicon–oxygen bond lengths vary between the various 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°. [9]

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. [10] 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. [11] The change in the coordination increases the ionicity of the Si-O bond. [12] More importantly, 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 alpha quartz, 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 lower densities and indices 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 are, due to the increased vibration energy.[ citation needed ]

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. [13]

The high-pressure minerals, seifertite, stishovite, and coesite, though, have higher densities and indices of refraction than quartz. This is probably due to the intense compression of the atoms occurring during their formation, resulting in more condensed structure. [14]

Faujasite silica is another form of crystalline silica. It is obtained by dealumination of a low-sodium, ultra-stable Y zeolite with combined acid and thermal treatment. The resulting product contains over 99% silica, and has high crystallinity and 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. [15]

Molten silica exhibits several peculiar physical characteristics that are similar to those observed in liquid water: negative temperature expansion, density maximum at temperatures ~5000 °C, and a heat capacity minimum. [16] Its density decreases from 2.08 g/cm3 at 1950 °C to 2.03 g/cm3 at 2200 °C. [17]

Molecular SiO2 with a linear structure is produced when molecular silicon monoxide, SiO, is condensed in an argon matrix cooled with helium along with oxygen atoms generated by microwave discharge. 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. [18]

## Natural occurrence

### Geology

Silica with the chemical formula Si O 2 is most commonly found in nature as quartz, which comprises more than 10% by mass of the earth's crust. [19] Quartz is the only polymorph of silica stable at the Earth's surface. Metastable occurrences of the high-pressure forms coesite and stishovite have been found around impact structures and associated with eclogites formed during ultra-high-pressure metamorphism. The high-temperature forms of tridymite and cristobalite are known from silica-rich volcanic rocks. In many parts of the world, silica is the major constituent of sand. [20] . The various forms of silicon dioxide can be converted from one form to another by heating and changes in pressure.

### Biology

Even though it is poorly soluble, silica occurs in many plants. Plant materials with high silica phytolith content appear to be of importance to grazing animals, from chewing insects to ungulates. Silica accelerates tooth wear, and high levels of silica in plants frequently eaten by herbivores may have developed as a defense mechanism against predation. [21] [22]

Silica is also the primary component of rice husk ash, which is used, for example, in filtration and cement manufacturing.

For well over a billion years, silicification in and by cells has been common in the biological world. In the modern world it occurs in bacteria, single-celled organisms, plants, and animals (invertebrates and vertebrates). Prominent examples include:

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. [23]

## Uses

### Structural use

About 95% of the commercial use of silicon dioxide (sand) occurs in the construction industry, e.g. for the production of concrete (Portland cement concrete). [19]

Silica, in the form of sand is used as the main ingredient in sand casting for the manufacture of metallic components in engineering and other applications. The high melting point of silica enables it to be used in such applications.

Crystalline silica is used in hydraulic fracturing of formations which contain tight oil and shale gas. [24]

### Precursor to glass and silicon

Silica is the primary ingredient in the production of most glass. The glass transition temperature of pure SiO2 is about 1475 K. [25] When molten silicon dioxide SiO2 is rapidly cooled, it does not crystallize, but solidifies as a glass.

The structural geometry of silicon and oxygen in glass is similar to that in quartz and most other crystalline forms of silicon and oxygen with silicon surrounded by regular tetrahedra of oxygen centers. The difference between the glass and crystalline forms arises from the connectivity of the tetrahedral units: Although there is no long range periodicity in the glassy network ordering remains at length scales well beyond the SiO bond length. One example of this ordering is the preference to form rings of 6-tetrahedra. [26]

### Fumed silica

Fumed silica also known as pyrogenic silica is a very fine particulate or colloidal form of silicon dioxide. It is prepared by burning SiCl4 in an oxygen-rich hydrogen flame to produce a "smoke" of SiO2. [10]

${\displaystyle {\ce {SiCl4 + 2 H2 + O2 -> SiO2 + 4 HCl}}}$

The majority of optical fibers for telecommunication are also made from silica. It is a primary raw material for many ceramics such as earthenware, stoneware, and porcelain.

Silicon dioxide is used to produce elemental silicon. The process involves carbothermic reduction in an electric arc furnace: [27]

${\displaystyle {\ce {SiO2 + 2 C -> Si + 2 CO}}}$

### Food and pharmaceutical applications

Silica is a common additive in food production, where it is used primarily as a flow agent in powdered foods, or to adsorb water in hygroscopic applications. It is used as an anti-caking agent in powdered foods such as spices and non-dairy coffee creamer. It is the primary component of diatomaceous earth. Colloidal silica is also used as a wine, beer, and juice fining agent. [19] It has the E number reference E551.

In pharmaceutical products, silica aids powder flow when tablets are formed.[ citation needed ]

### Personal care

In cosmetics, it is useful for its light-diffusing properties [28] and natural absorbency. [29]

Hydrated silica is used in toothpaste as a hard abrasive to remove tooth plaque.

### Semiconductors

Silicon dioxide could be grown on a silicon semiconductor surface. [30] Silicon oxide layers could protect silicon surfaces during diffusion processes, and could be used for diffusion masking. [31] [32]

Surface passivation is the process by which a semiconductor surface is rendered inert, and does not change semiconductor properties as a result of interaction with air or other materials in contact with the surface or edge of the crystal. [33] [34] [35] The formation of a thermally grown silicon dioxide layer greatly reduces the concentration of electronic states at the silicon surface. [35] SiO2 films preserve the electrical characteristics of p–n junctions and prevent these electrical characteristics from deteriorating by the gaseous ambient environment. [32] Silicon oxide layers could be used to electrically stabilize silicon surfaces. [31] The surface passivation process is an important method of semiconductor device fabrication that involves coating a silicon wafer with an insulating layer of silicon oxide so that electricity could reliably penetrate to the conducting silicon below. Growing a layer of silicon dioxide on top of a silicon wafer enables it to overcome the surface states that otherwise prevent electricity from reaching the semiconducting layer. [34] [36]

The process of silicon surface passivation by thermal oxidation (silicon dioxide) is critical to the semiconductor industry. It is commonly used to manufacture metal-oxide-semiconductor field-effect transistors (MOSFETs) and silicon integrated circuit chips (with the planar process). [34] [36]

### Other

Hydrophobic silica is used as a defoamer component. [37]

In its capacity as a refractory, it is useful in fiber form as a high-temperature thermal protection fabric.[ citation needed ]

Silica is used in the extraction of DNA and RNA due to its ability to bind to the nucleic acids under the presence of chaotropes. [38]

A silica-based aerogel was used in the Stardust spacecraft to collect extraterrestrial particles. [39]

Pure silica (silicon dioxide), when cooled as fused quartz into a glass with no true melting point, can be used as a glass fiber for fiberglass.

## Production

Silicon dioxide is mostly obtained by mining, including sand mining and purification of quartz. Quartz is suitable for many purposes, while chemical processing is required to make a purer or otherwise more suitable (e.g. more reactive or fine-grained) product.[ citation needed ]

### Silica fume

Silica fume is obtained as byproduct from hot processes like ferrosilicon production. It is less pure than fumed silica and should not be confused with that product. The production process, particle characteristics, and fields of application of fumed silica differ from those of silica fume.

### Precipitated silica

Precipitated silica or amorphous silica is produced by the acidification of solutions of sodium silicate. The gelatinous precipitate or silica gel, is first washed and then dehydrated to produce colorless microporous silica. [10] The idealized equation involving a trisilicate and sulfuric acid is:

${\displaystyle {\ce {Na2Si3O7 + H2SO4 -> 3 SiO2 + Na2SO4 + H2O}}}$

Approximately one billion kilograms/year (1999) of silica were produced in this manner, mainly for use for polymer composites – tires and shoe soles. [19]

### On microchips

Thin films of silica grow spontaneously on silicon wafers via thermal oxidation, producing a very shallow layer of about 1 nm or 10 Å of so-called native oxide. [40] Higher temperatures and alternative environments are used to grow well-controlled layers of silicon dioxide on silicon, for example at temperatures between 600 and 1200 °C, using so-called dry oxidation with O2

${\displaystyle {\ce {Si + O2 -> SiO2}}}$

or wet oxidation with H2O. [41] [42]

${\displaystyle {\ce {Si + 2 H2O -> SiO2 + 2 H2}}}$

The native oxide layer is beneficial in microelectronics, where it acts as electric insulator with high chemical stability. It can protect the silicon, store charge, block current, and even act as a controlled pathway to limit current flow. [43]

### Laboratory or special methods

#### From organosilicon compounds

Many routes to silicon dioxide start with an organosilicon compound, e.g., HMDSO [44] , TEOS. Synthesis of silica is illustrated below using tetraethyl orthosilicate (TEOS). Simply heating TEOS at 680–730 °C results in the oxide:

${\displaystyle {\ce {Si(OC2H5)4 -> SiO2 + 2 O(C2H5)2}}}$

Similarly TEOS combusts around 400 °C:

${\displaystyle {\ce {Si(OC2H5)4 + 12 O2 -> SiO2 + 10 H2O + 8 CO2}}}$

TEOS undergoes hydrolysis via the so-called sol-gel process. The course of the reaction and nature of the product are affected by catalysts, but the idealized equation is: [45]

${\displaystyle {\ce {Si(OC2H5)4 + 2 H2O -> SiO2 + 4 HOCH2CH3}}}$

#### Other methods

Being highly stable, silicon dioxide arises from many methods. Conceptually simple, but of little practical value, combustion of silane gives silicon dioxide. This reaction is analogous to the combustion of methane:

${\displaystyle {\ce {SiH4 + 2 O2 -> SiO2 + 2 H2O}}}$

However the chemical vapor deposition of silicon dioxide onto crystal surface from silane had been used using nitrogen as a carrier gas at 200–500 °C. [46]

## Chemical reactions

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. [10]

Silicon dioxide is attacked by hydrofluoric acid (HF) to produce hexafluorosilicic acid: [9]

${\displaystyle {\ce {SiO2 + 6 HF -> H2SiF6 + 2 H2O}}}$

HF is used to remove or pattern silicon dioxide in the semiconductor industry.

Under normal conditions, silicon does not react with most acids but is dissolved by hydrofluoric acid.

${\displaystyle {\ce {Si(s) + 6HF(aq) -> [SiF6]^{2-}(aq) + 2H+(aq) + 2H2(g)}}}$

Silicon is attacked by bases such as aqueous sodium hydroxide to give silicates.

${\displaystyle {\ce {Si(s) + 4NaOH(aq) -> [SiO4]^{4-}(aq) + 4Na+(aq) + 2H2(g)}}}$

Silicon dioxide acts as a Lux–Flood acid, being able to react with bases under certain conditions. As it does not contain any hydrogen, it cannot act as a Brønsted–Lowry acid. While not soluble in water, some strong bases will react with glass and have to be stored in plastic bottles as a result. [47]

Silicon dioxide dissolves in hot concentrated alkali or fused hydroxide, as described in this idealized equation: [10]

${\displaystyle {\ce {SiO2 + 2 NaOH -> Na2SiO3 + H2O}}}$

Silicon dioxide will neutralise 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). [9] As an example the reaction of sodium oxide and SiO2 can produce sodium orthosilicate, sodium silicate, and glasses, dependent on the proportions of reactants: [10]

${\displaystyle {\ce {2 Na2O + SiO2 -> Na4SiO4;}}}$
${\displaystyle {\ce {Na2O + SiO2 -> Na2SiO3;}}}$
(0.25–0.8) ${\displaystyle {\ce {Na2O + SiO2 -> glass}}}$.

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. [9] The reaction is also used in blast furnaces to remove sand impurities in the ore by neutralisation with calcium oxide, forming calcium silicate slag.

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. [48] The silicates are essentially insoluble in all polar solvent except methanol.

Silicon dioxide reacts with elemental silicon at high temperatures to produce SiO: [9]

${\displaystyle {\ce {SiO2 + Si -> 2 SiO}}}$

### Water solubility

The solubility of silicon dioxide in water strongly depends on its crystalline form and is three-four times higher for silica[ clarification needed ] than quartz; as a function of temperature, it peaks around 340 °C. [49] 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. [9] These crystals are a source of very pure quartz for use in electronic applications. [10]

## Health effects

Silica ingested orally is essentially nontoxic, with an LD50 of 5000 mg/kg (5 g/kg). [19] A 2008 study following subjects for 15 years found that higher levels of silica in water appeared to decrease the risk of dementia. An increase of 10 mg/day of silica in drinking water was associated with a decreased risk of dementia of 11%. [50]

Inhaling finely divided crystalline silica dust can lead to silicosis, bronchitis, or lung cancer, as the dust becomes lodged in the lungs and continuously irritates the tissue, reducing lung capacities. [51] When fine silica particles are inhaled in large enough quantities (such as through occupational exposure), it increases the risk of systemic autoimmune diseases such as lupus [52] and rheumatoid arthritis compared to expected rates in the general population. [53]

### Occupational hazard

Silica is an occupational hazard for people who do sandblasting, or work with products that contain powdered crystalline silica. Amorphous silica, such as fumed silica, may cause irreversible lung damage in some cases, but is not associated with development of silicosis. Children, asthmatics of any age, those with allergies, and the elderly (all of whom have reduced lung capacity) can be affected in less time. [54]

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. [55] Crystalline silica used in hydraulic fracturing presents a health hazard to workers. [24]

### Pathophysiology

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 production of interleukin, a highly pro-inflammatory cytokine in the immune system. [56] [57] [58]

### Regulation

Regulations restricting silica exposure 'with respect to the silicosis hazard' specify that they are concerned only with silica, which is both crystalline and dust-forming. [59] [60] [61] [62] [63] [64]

In 2013, the U.S. Occupational Safety and Health Administration reduced the exposure limit to 50 µg/m3 of air. Prior to 2013, it had allowed 100  µg/m3 and in construction workers even 250 µg/m3. [24] In 2013, OSHA also required "green completion" of fracked wells to reduce exposure to crystalline silica besides restricting the limit of exposure. [24]

## Crystalline forms

SiO2, more so than almost any material, exists in many crystalline forms. These forms are called polymorphs.

Crystalline forms of SiO2 [9]
FormCrystal symmetry
Pearson symbol, group No.
ρ
g/cm3
NotesStructure
α-quartz rhombohedral (trigonal)
hP9, P3121 No.152 [65]
2.648Helical chains making individual single crystals optically active; α-quartz converts to β-quartz at 846 K
β-quartz hexagonal
hP18, P6222, No. 180 [66]
2.533Closely related to α-quartz (with an Si-O-Si angle of 155°) and optically active; β-quartz converts to β-tridymite at 1140 K
α-tridymite orthorhombic
oS24, C2221, No.20 [67]
2.265Metastable form under normal pressure
β-tridymitehexagonal
hP12, P63/mmc, No. 194 [67]
Closely related to α-tridymite; β-tridymite converts to β-cristobalite at 2010 K
α-cristobalite tetragonal
tP12, P41212, No. 92 [68]
2.334Metastable form under normal pressure
β-cristobalite cubic
cF104, Fd3m, No.227 [69]
Closely related to α-cristobalite; melts at 1978 K
keatite tetragonal
tP36, P41212, No. 92 [70]
3.011Si5O10, Si4O8, Si8O16 rings; synthesised from glassy silica and alkali at 600–900 K and 40–400 MPa
moganite monoclinic
mS46, C2/c, No.15 [71]
Si4O8 and Si6O12 rings
coesite monoclinic
mS48, C2/c, No.15 [72]
2.911Si4O8 and Si8O16 rings; 900 K and 3–3.5 GPa
stishovite tetragonal
tP6, P42/mnm, No.136 [73]
4.287One of the densest (together with seifertite) polymorphs of silica; rutile-like with 6-fold coordinated Si; 7.5–8.5 GPa
seifertite orthorhombic
oP, Pbcn [74]
4.294One of the densest (together with stishovite) polymorphs of silica; is produced at pressures above 40 GPa. [75]
melanophlogite cubic (cP*, P4232, No.208) [8] or tetragonal (P42/nbc) [76] 2.04Si5O10, Si6O12 rings; mineral always found with hydrocarbons in interstitial spaces - a clathrasil [77]
fibrous
W-silica [10]
orthorhombic
oI12, Ibam, No.72 [78]
1.97Like SiS2 consisting of edge sharing chains, melts at ~1700 K
2D silica [79] hexagonalSheet-like bilayer structure

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Germanium dioxide, also called germanium oxide, germania, and salt of germanium, is an inorganic compound with the chemical formula GeO2. It is the main commercial source of germanium. It also forms as a passivation layer on pure germanium in contact with atmospheric oxygen.

Tetraethyl orthosilicate, formally named tetraethoxysilane and abbreviated TEOS, is the chemical compound with the formula Si(OC2H5)4. TEOS is a colorless liquid that degrades in water. TEOS is the ethyl ester of orthosilicic acid, Si(OH)4. It is the most prevalent alkoxide of silicon.

Silicon monoxide is the chemical compound with the formula SiO where silicon is present in the oxidation state +2. In the vapour phase it is a diatomic molecule..

The structure of liquids, glasses and other non-crystalline solids is characterized by the absence of long-range order which defines crystalline materials. Liquids and amorphous solids do, however, possess a rich and varied array of short to medium range order, which originates from chemical bonding and related interactions. Metallic glasses, for example, are typically well described by the dense random packing of hard spheres, whereas covalent systems, such as silicate glasses, have sparsely packed, strongly bound, tetrahedral network structures. These very different structures result in materials with very different physical properties and applications.

Silicon oxynitride is a ceramic material with the chemical formula SiOxNy. While in amorphous forms its composition can continuously vary between SiO2 (silica) and Si3N4 (silicon nitride), the only known intermediate crystalline phase is Si2N2O. It is found in nature as the rare mineral sinoite in some meteorites and can be synthesized in the laboratory.

Silicon carbonate is a crystalline substance formed under pressure from silica and carbon dioxide. The formula of the substance is SiCO4. To produce it silicalite is compressed with carbon dioxide at a pressure of 18 Gpa and a temperature around 740K. The silicon carbonate made this way has carbonate linked to silicon by way of oxygen in unidentate, bidentate, or bridged positions. However a stable crystal structure is not formed in these conditions. The phase produced is amorphous, but it has carbon in three-fold coordination, and silicon in six-fold coordination. When decompressed, not all carbon is released as carbon dioxide. If this really exists, the substance should be dynamically stable when reduced to atmospheric pressure.

Mineral alteration refers to the various natural processes that alter a mineral's chemical composition or crystallography.

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