|E number||E551 (acidity regulators, ...)|
CompTox Dashboard (EPA)
|Molar mass||60.08 g/mol|
|Appearance||Transparent solid (Amorphous) White/Whitish Yellow (Powder/Sand)|
|Density||2.648 (?-quartz), 2.196 (amorphous) g·cm-3|
|Melting point||1,713 °C (3,115 °F; 1,986 K) (amorphous)(p4.88) to|
|Boiling point||2,950 °C (5,340 °F; 3,220 K)|
|Thermal conductivity||12 (|| c-axis), 6.8 (? c-axis), 1.4 (am.) W/(m?K)(p12.213)|
Refractive index (nD)
|1.544 (o), 1.553 (e)(p4.143)|
|NFPA 704 (fire diamond)|
|NIOSH (US health exposure limits):|
|TWA 20 mppcf (80 mg/m3/%SiO2) (amorphous)|
|TWA 6 mg/m3 (amorphous)|
Ca TWA 0.05 mg/m3
IDLH (Immediate danger)
|3000 mg/m3 (amorphous)|
Ca [25 mg/m3 (cristobalite, tridymite); 50 mg/m3 (quartz)]
Std enthalpy of
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
|what is ?)(|
Silicon dioxide, also known as silica, is an oxide of silicon with the chemical formula SiO2, most commonly found in nature as quartz and in various living organisms. 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 a 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.
In the majority of silicates, the silicon atom shows tetrahedral coordination, with four oxygen atoms surrounding a central Si atom (see 3-D Unit Cell). Thus, SiO2 forms 3 dimensional network solids in which each silicon atom is covalently bonded in a tetrahedral manner to 4 oxygen atoms. In contrast, CO2 is a linear molecule. The starkly different structures of the dioxides of carbon and silicon is a manifestation of the Double bond rule.
SiO2 has a number of distinct crystalline forms, but they almost always have the same local structure around Si and O. In ?-quartz the Si-O bond length is 161 pm, whereas in ?-tridymite it is in the range 154-171 pm. The Si-O-Si angle also varies between a low value of 140° in ?-tridymite, up to 180° in ?-tridymite. In ?-quartz, the Si-O-Si angle is 144°.
Alpha quartz is the stable form of solid SiO2 at room temperature. The high-temperature minerals, cristobalite and tridymite, have both lower densities and indices of refraction than quartz. The transformation from ?-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, though, have higher densities and indices of refraction than quartz.  Stishovite has a rutile-like structure where silicon is 6-coordinate. The density of stishovite is 4.287 g/cm3, which compares to ?-quartz, the densest of the low-pressure forms, which has a density of 2.648 g/cm3. 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 ?-quartz. The change in the coordination increases the ionicity of the Si-O bond. More importantly, any deviations from these standard parameters constitute microstructural differences or variations, which represent an approach to an amorphous, vitreous, or glassy solid.
Faujasite silica, another polymorph, 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.
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. Its density decreases from 2.08 g/cm3 at 1950 °C to 2.03 g/cm3 at 2200 °C.
Molecular SiO2 is linear structure. It has been produced by combining silicon monoxide with oxygen atoms in an argon matrix. Dimeric silicon dioxide, (SiO2)2 has been generated 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 ?-quartz. The bond energy is estimated at 621.7 kJ/mol.
SiO2 is most commonly found in nature as quartz, which comprises more than 10% by mass of the earth's crust. 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.
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.
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).
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, creating the challenge of creating 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.
Certain deposits of silica sand, with desirable particle size and shape and desirable clay and other mineral content, were important for sand casting of metallic products. The high melting point of silica enables it to be used in such applications such as iron casting; modern sand casting sometimes uses other minerals for other reasons.
Silica is the primary ingredient in the production of most glass. As other minerals are melted with silica, the principle of Freezing Point Depression lowers the melting point of the mixture and increases fluidity. The glass transition temperature of pure SiO2 is about 1475 K. When molten silicon dioxide SiO2 is rapidly cooled, it does not crystallize, but solidifies as a glass. Because of this, most ceramic glazes have silica as the main ingredient.
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.
It can also be produced by vaporizing quartz sand in a 3000 °C electric arc. Both processes result in microscopic droplets of amorphous silica fused into branched, chainlike, three-dimensional secondary particles which then agglomerate into tertiary particles, a white powder with extremely low bulk density (0.03-.15 g/cm3) and thus high surface area. The particles act as a thixotropic thickening agent, or as an anti-caking agent, and can be treated to make them hydrophilic or hydrophobic for either water or organic liquid applications
Silica fume is an ultrafine powder collected as a by-product of the silicon and ferrosilicon alloy production. It consists of amorphous (non-crystalline) spherical particles with an average particle diameter of 150 nm, without the branching of the pyrogenic product. The main use is as pozzolanic material for high performance concrete.
Silica, either colloidal, precipitated, or pyrogenic fumed, is a common additive in food production. It is used primarily as a flow or anti-caking agent in powdered foods such as spices and non-dairy coffee creamer, or powders to be formed into pharmaceutical tablets. It can adsorb water in hygroscopic applications. Colloidal silica is used as a fining agent for wine, beer, and juice, with the E number reference E551.
Diatomaceous earth, a mined product, has been used in food and cosmetics for centuries. It consists of the silica shells of microscopic diatoms; in a less processed form it was sold as "tooth powder". Manufactured or mined hydrated silica is used as the hard abrasive in toothpaste.
Silicon dioxide is widely used in the semiconductor technology
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. The formation of a thermally grown silicon dioxide layer greatly reduces the concentration of electronic states at the silicon surface. SiO2 films preserve the electrical characteristics of p-n junctions and prevent these electrical characteristics from deteriorating by the gaseous ambient environment. Silicon oxide layers could be used to electrically stabilize silicon surfaces. 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.
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).
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.
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.
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. The idealized equation involving a trisilicate and sulfuric acid is:
Approximately one billion kilograms/year (1999) of silica were produced in this manner, mainly for use for polymer composites - tires and shoe soles.
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. 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
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.
Many routes to silicon dioxide start with an organosilicon compound, e.g., HMDSO, TEOS. Synthesis of silica is illustrated below using tetraethyl orthosilicate (TEOS). Simply heating TEOS at 680-730 °C results in the oxide:
Similarly TEOS combusts around 400 °C:
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:
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.
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.
Silicon is attacked by bases such as aqueous sodium hydroxide to give silicates.
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 silicon dioxide is not soluble in water, some strong bases will react with glass and have to be stored in plastic bottles as a result.
Silicon dioxide dissolves in hot concentrated alkali or fused hydroxide, as described in this idealized equation:
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). As an example the reaction of sodium oxide and SiO2 can produce sodium orthosilicate, sodium silicate, and glasses, dependent on the proportions of reactants:
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. 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. The silicates are essentially insoluble in all polar solvent except methanol.
Silicon dioxide reacts with elemental silicon at high temperatures to produce SiO:
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. 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. These crystals are a source of very pure quartz for use in electronic applications.
Silica ingested orally is essentially nontoxic, with an LD50 of 5000 mg/kg (5 g/kg). 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%.
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. 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 and rheumatoid arthritis compared to expected rates in the general population.
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.
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. Crystalline silica used in hydraulic fracturing presents a health hazard to workers.
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.
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. In 2013, OSHA also required "green completion" of fracked wells to reduce exposure to crystalline silica besides restricting the limit of exposure.
SiO2, more so than almost any material, exists in many crystalline forms. These forms are called polymorphs.
Pearson symbol, group No.
hP9, P3121 No.152
|2.648||Helical chains making individual single crystals optically active; ?-quartz converts to ?-quartz at 846 K|
hP18, P6222, No. 180
|2.533||Closely related to ?-quartz (with an Si-O-Si angle of 155°) and optically active; ?-quartz converts to ?-tridymite at 1140 K|
oS24, C2221, No.20
|2.265||Metastable form under normal pressure|
hP12, P63/mmc, No. 194
|Closely related to ?-tridymite; ?-tridymite converts to ?-cristobalite at 2010 K|
tP12, P41212, No. 92
|2.334||Metastable form under normal pressure|
cF104, Fd3m, No.227
|Closely related to ?-cristobalite; melts at 1978 K|
tP36, P41212, No. 92
|3.011||Si5O10, Si4O8, Si8O16 rings; synthesised from glassy silica and alkali at 600-900 K and 40-400 MPa|
mS46, C2/c, No.15
|Si4O8 and Si6O12 rings|
mS48, C2/c, No.15
|2.911||Si4O8 and Si8O16 rings; 900 K and 3-3.5 GPa|
tP6, P42/mnm, No.136
|4.287||One of the densest (together with seifertite) polymorphs of silica; rutile-like with 6-fold coordinated Si; 7.5-8.5 GPa|
|4.294||One of the densest (together with stishovite) polymorphs of silica; is produced at pressures above 40 GPa.|
|melanophlogite||cubic (cP*, P4232, No.208) or tetragonal (P42/nbc)||2.04||Si5O10, Si6O12 rings; mineral always found with hydrocarbons in interstitial spaces - a clathrasil (silica clathrate)|
oI12, Ibam, No.72
|1.97||Like SiS2 consisting of edge sharing chains, melts at ~1700 K|
|2D silica||hexagonal||Sheet-like bilayer structure|
Inhaling finely divided crystalline silica 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.
This extended list enumerates synonyms for silicon dioxide; all of these values are from a single source; values in the source were presented capitalized.
These soft focus pigments, mainly composed of polymers, micas and talcs covered with rough or spherical particles of small diameters, such as silica or titanium dioxide, are used to optically reduce the appearance of wrinkles. These effects are obtained by optimizing outlines of wrinkles and reducing the difference of brightness due to diffuse reflection.
The silica is a multiporous ingredient, which absorbs the oil and sebum.