Silicon dioxide

Silicon dioxide
Names
IUPAC name
Silicon dioxide
Other names
Quartz

Silica
Silicic oxide
Silicon(IV) oxide

Crystalline silica
Identifiers
ChEBI
ChemSpider
ECHA InfoCard 100.028.678
EC Number 231-545-4
E number E551 (acidity regulators, ...)
200274
KEGG
MeSH Silicon+dioxide
RTECS number VV7565000
UNII
Properties
SiO2
Molar mass 60.08 g/mol
Appearance Transparent 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
Flammability code 0: Will not burn. E.g., water Health code 0: Exposure under fire conditions would offer no hazard beyond that of ordinary combustible material. E.g., sodium chloride Reactivity code 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g., liquid nitrogen Special hazards (white): no codeNFPA 704 four-colored diamond
0
0
0
US health exposure limits (NIOSH):
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

Germanium dioxide
Tin dioxide
Lead dioxide

Related compounds
 Silicon monoxide

Silicon sulfide

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).
N verify (what is YesYN ?)
Infobox references

Silicon dioxide, also known as silica (from the Latin silex), is an oxide of silicon with the chemical formula SiO2, 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 component in the food and pharmaceutical industry.

Inhaling finely divided crystalline silica is toxic and can lead to silicosis, bronchitis, lung cancer and systemic autoimmune diseases, such as lupus and rheumatoid arthritis.

Structure

Structural motif found in α-quartz, but also found in almost all forms of silicon dioxide

In the majority of silicates, the Si atom shows tetrahedral coordination, with four oxygen atoms surrounding a central Si atom. The most common example is seen in the quartzite polymorph.

Relationship between refractive index and density for some SiO2 forms[7]

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 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°.[8]

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.[9] 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.[10] The change in the coordination increases the ionicity of the Si-O bond.[11] 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.

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, 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.

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.[12]

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.[13] Its density decreases from 2.08 g/cm3 at 1950 °C to 2.03 g/cm3 at 2200 °C.[14]

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.[15]

Natural occurrence

Geology

Silica with the chemical formula SiO2 is most commonly found in nature as quartz SiO4, which comprises more than 10% by mass of the earth's crust.[16] In many parts of the world, silica is the major constituent of sand.

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.[17][18]

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.[19]

Uses

Structural use

An estimated 95% of silicon dioxide produced is consumed in the construction industry, e.g. for the production of Portland cement.[16]

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.[20]

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.[21] 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.[22]

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.[9]

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:[23]

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 the primary component of diatomaceous earth. Colloidal silica is also used as a wine, beer, and juice fining agent.[16]

In pharmaceutical products, silica aids powder flow when tablets are formed.

Personal care

In cosmetics, its useful for its light-diffusing properties and natural absorbency.

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

Other

Hydrophobic silica is used as a defoamer component.

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

It is used as a thermal enhancement compound in the ground source heat pump industry.

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.

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

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.

Silica fume

Silica fume is obtained as byproduct from hot processes like ferro-silicon 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 are all different 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.[9] Idealized equation involving a trisilicate and sulfuric acid is shown:

Na2Si3O7 + H2SO4 → 3 SiO2 + Na2SO4 + H2O

Approximately one billion kilograms/year (1999) of silica was produced in this manner, mainly for use for polymer composites – tires and shoe soles.[16]

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.[24] 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 or wet oxidation with O2

Si + O2 → SiO2

or H2O, respectively.[25][26]

Si + 2 H2O → SiO2 + 2 H2

The depth of the layer of silicon which dioxide replaces is 44% of the depth of the silicon dioxide layer produced.[25]

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.[27]

Laboratory or special methods

From silicate esters

Many routes to silicon dioxide start with silicate esters, the best known being tetraethyl orthosilicate (TEOS). Simply heating TEOS at 680–730 °C gives the dioxide:

Si(OC2H5)4 → SiO2 + 2 O(C2H5)2

Similarly TEOS combusts around 400 °C:

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:[28]

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:

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 200-500oC.[29]

Chemical reactions

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.[9]

Silicon dioxide is attacked by hydrofluoric acid (HF) to produce hexafluorosilicic acid:[8]

SiO2 + 6 HF → H2SiF6 + 2 H2O.

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

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.[30]

Silicon dioxide dissolves in hot concentrated alkali or fused hydroxide, as described in this idealized equation:[9]

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

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

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

Silicon dioxide reacts with elemental silicon at high temperatures to produce SiO:[8]

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 than quartz; as a function of temperature, it peaks around 340°C.[32] 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.[8] These crystals are a source of very pure quartz for use in electronic applications.[9]

Health effects

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).[16] 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 milligram-per-day of silica in drinking water was associated with a decreased risk of dementia of 11%.[33]

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.[34] It increases the risk of of systemic autoimmune diseases such as lupus[35] and rheumatoid arthritis compared to expected rates in the general population.[36]

Occupational hazard

Silica is an occupational hazard for people who do sandblasting, 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.[37]

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.[38] Crystalline silica used in hydraulic fracturing presents a health hazard to workers.[20]

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.[39][40][41]

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.[42][43][44][45][46]

In 2013, the U.S. Occupational Safety and Health Administration reduced the exposure limit to 50µg per cubic meter of air. Prior to 2013 it had allowed 100  µg/m3 and in construction workers even 250 µg/m3.[20] In 2013, OSHA also required "green completion" of fracked wells to reduce exposure to crystalline silica besides restricting the limit of exposure.[20]

Crystalline forms

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

Crystalline forms of SiO2[8]
Form Crystal symmetry
Pearson symbol, group No.		 ρ
g/cm3		 Notes Structure
α-quartz rhombohedral (trigonal)
hP9, P3121 No.152[47]		 2.648 Helical chains making individual single crystals optically active; α-quartz converts to β-quartz at 846 K
β-quartz hexagonal
hP18, P6222, No. 180[48]		 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
α-tridymite orthorhombic
oS24, C2221, No.20[49]		 2.265 Metastable form under normal pressure
β-tridymite hexagonal
hP12, P63/mmc, No. 194[49]		 Closely related to α-tridymite; β-tridymite converts to β-cristobalite at 2010 K
α-cristobalite tetragonal
tP12, P41212, No. 92[50]		 2.334 Metastable form under normal pressure
β-cristobalite cubic
cF104, Fd3m, No.227[51]		 Closely related to α-cristobalite; melts at 1978 K
keatite tetragonal
tP36, P41212, No. 92[52]		 3.011 Si5O10, Si4O14, Si8O16 rings; synthesised from glassy silica and alkali at 600–900 K and 40–400 MPa
moganite monoclinic
mS46, C2/c, No.15[53]		 Si4O8 and Si6O12 rings
coesite monoclinic
mS48, C2/c, No.15[54]		 2.911 Si4O8 and Si8O16 rings; 900 K and 3–3.5 GPa
stishovite tetragonal
tP6, P42/mnm, No.136[55]		 4.287 One 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[56]		 4.294 One of the densest (together with stishovite) polymorphs of silica; is produced at pressures above 40 GPa.[57]
melanophlogite cubic (cP*, P4232, No.208)[7] or tetragonal (P42/nbc)[58] 2.04 Si5O10, Si6O12 rings; mineral always found with hydrocarbons in interstitial spaces - a clathrasil[59]
fibrous
W-silica[9]		 orthorhombic
oI12, Ibam, No.72[60]		 1.97 Like SiS2 consisting of edge sharing chains, melts at ~1700 K
2D silica[61] hexagonal Sheet-like bilayer structure

See also

References

  1. 1 2 3 4 5 Haynes, William M., ed. (2011). CRC Handbook of Chemistry and Physics (92nd ed.). Boca Raton, FL: CRC Press. ISBN 1439855110.
  2. 1 2 3 "NIOSH Pocket Guide to Chemical Hazards #0552". National Institute for Occupational Safety and Health (NIOSH).
  3. 1 2 "NIOSH Pocket Guide to Chemical Hazards #0682". National Institute for Occupational Safety and Health (NIOSH).
  4. 1 2 Zumdahl, Steven S. (2009). Chemical Principles 6th Ed. Houghton Mifflin Company. p. A22. ISBN 0-618-94690-X.
  5. ↑ Iler, R. K. (1979). The Chemistry of Silica. Plenum Press. ISBN 0-471-02404-X.
  6. 1 2 Fernández LD, Lara E, Mitchell EA (2015). "Checklist, diversity and distribution of testate amoebae in Chile". Eur. J. Protistol. 51 (5): 409–24. PMID 26340665. doi:10.1016/j.ejop.2015.07.001.
  7. 1 2 Skinner B. J.; Appleman D. E. (1963). "Melanophlogite, a cubic polymorph of silica" (PDF). American Mineralogist. 48: 854–867.
  8. 1 2 3 4 5 6 7 Holleman, A. F.; Wiberg, E. (2001), Inorganic Chemistry, San Diego: Academic Press, ISBN 0-12-352651-5
  9. 1 2 3 4 5 6 7 8 Greenwood, Norman N.; Earnshaw, Alan (1984). Chemistry of the Elements. Oxford: Pergamon Press. pp. 393–99. ISBN 0-08-022057-6.
  10. ↑ Wells A. F. (1984). Structural Inorganic Chemistry. Oxford Science Publications. ISBN 0-19-855370-6.
  11. ↑ Kirfel, A.; Krane, H. G.; Blaha, P.; Schwarz, K.; Lippmann, T. (2001). "Electron-density distribution in stishovite, SiO2: a new high-energy synchrotron-radiation study". Acta Crystallographica A. 57 (6): 663. doi:10.1107/S0108767301010698.
  12. ↑ Scherzer, J. (1978). "Dealuminated faujasite-type structures with SiO2/Al2O3 ratios over 100". Journal of Catalysis. 54 (2): 285. doi:10.1016/0021-9517(78)90051-9.
  13. ↑ Shell, Scott M.; Debenedetti, Pablo G.; Panagiotopoulos, Athanassios Z. (2002). "Molecular structural order and anomalies in liquid silica" (PDF). Physical Review E. 66: 011202. Bibcode:2002PhRvE..66a1202S. arXiv:cond-mat/0203383 Freely accessible. doi:10.1103/PhysRevE.66.011202.
  14. ↑ Aksay, I. A.; Pask, J. A. & Davis, R. F. (1979). "Densities of SiO2-Al2O3 Melts" (PDF). Journal of the American Ceramic Society. 62 (7–8): 332–336. doi:10.1111/j.1151-2916.1979.tb19071.x.
  15. ↑ Jutzi, Peter & Schubert, Ulrich (2003). Silicon chemistry: from the atom to extended systems. Wiley-VCH. ISBN 3-527-30647-1.
  16. 1 2 3 4 5 Flörke, Otto W. et al. (2008) "Silica" in Ullmann's Encyclopedia of Industrial Chemistry, Weinheim: Wiley-VCH, . doi:10.1002/14356007.a23_583.pub3.
  17. ↑ Massey, Fergus P.; Ennos, A. Roland; Hartley, Sue E. (2006). "Silica in grasses as a defence against insect herbivores: Contrasting effects on folivores and a phloem feeder". Journal of Animal Ecology. 75 (2): 595–603. PMID 16638012. doi:10.1111/j.1365-2656.2006.01082.x.
  18. ↑ Keeping, Malcolm G.; Kvedaras, Olivia L. (2008). "Silicon as a plant defence against insect herbivory: Response to Massey, Ennos and Hartley". Journal of Animal Ecology. 77 (3): 631–3. PMID 18341561. doi:10.1111/j.1365-2656.2008.01380.x.
  19. ↑ Carlisle, EM (1986). "Silicon as an essential trace element in animal nutrition". Ciba Foundation Symposium. Novartis Foundation Symposia. 121: 123–39. ISBN 978-0-470-51332-3. PMID 3743227. doi:10.1002/9780470513323.ch8.
  20. 1 2 3 4 Greenhouse, Steven (August 23, 2013). "New Rules Would Cut Silica Dust Exposure". The New York Times. Retrieved August 24, 2013.
  21. ↑ Ojovan, M. I. (2004). "Glass formation in amorphous SiO2 as a percolation phase transition in a system of network defects". Journal of Experimental and Theoretical Physics Letters. 79 (12): 632–634. Bibcode:2004JETPL..79..632O. doi:10.1134/1.1790021.
  22. ↑ Elliott, S. R. (1991). "Medium-range structural order in covalent amorphous solids". Nature. 354 (6353): 445–452. Bibcode:1991Natur.354..445E. doi:10.1038/354445a0.
  23. ↑ Shriver and Atkins. Inorganic Chemistry (5th Edition). W. H. Freeman and Company, New York, 2010, p. 354.
  24. ↑ Doering, Robert & Nishi, Yoshio (2007). Handbook of Semiconductor Manufacturing Technology. CRC Press. ISBN 1-57444-675-4.
  25. 1 2 Lee, Sunggyu (2006). Encyclopedia of chemical processing. CRC Press. ISBN 0-8247-5563-4.
  26. ↑ Morgan, D. V.; Board, K. (1991). An Introduction To Semiconductor Microtechnology (2nd ed.). Chichester, West Sussex, England: John Wiley & Sons. p. 72. ISBN 0471924784.
  27. ↑ Riordan, Michael (2007) "The Silicon Dioxide Solution: How physicist Jean Hoerni built the bridge from the transistor to the integrated circuit" IEEE Spectrum.
  28. ↑ Nandiyanto, A. B. D.; Kim, S. G.; Iskandar, F.; Okuyama, K. (2009). "Synthesis of spherical mesoporous silica nanoparticles with nanometer-size controllable pores and outer diameters". Microporous and Mesoporous Materials. 120 (3): 447. doi:10.1016/j.micromeso.2008.12.019.
  29. ↑ Morgan, D. V.; Board, K. (1991). An Introduction To Semiconductor Microtechnology (2nd ed.). Chichester, West Sussex, England: John Wiley & Sons. p. 27. ISBN 0471924784.
  30. ↑ Glen E. Rodgers (19 January 2011). Descriptive Inorganic, Coordination, and Solid State Chemistry. Cengage Learning. pp. 421–2. ISBN 1-133-17248-2.
  31. ↑ Laine, Richard M.; Blohowiak, Kay Youngdahl; Robinson, Timothy R.; Hoppe, Martin L.; Nardi, Paola; Kampf, Jeffrey; Uhm, Jackie (17 October 1991). "Synthesis of pentacoordinate silicon complexes from SiO2". Nature. 353 (353): 642–644. Bibcode:1991Natur.353..642L. doi:10.1038/353642a0.
  32. ↑ Fournier R. O.; Rowe J. J. (1977). "The solubility of amorphous silica in water at high temperatures and high pressures" (PDF). American Mineralogist. 62: 1052–1056.
  33. ↑ Rondeau, V.; Jacqmin-Gadda, H.; Commenges, D.; Helmer, C.; Dartigues, J.-F. (2008). "Aluminum and Silica in Drinking Water and the Risk of Alzheimer's Disease or Cognitive Decline: Findings from 15-Year Follow-up of the PAQUID Cohort". American Journal of Epidemiology. 169 (4): 489–96. PMC 2809081 Freely accessible. PMID 19064650. doi:10.1093/aje/kwn348.
  34. ↑ CPWR-Center for Construction Research and Training — Work Safely with Silica: "What are the Health Effects?
  35. ↑ The Future Directions of Lupus Research. niams.nih.gov
  36. ↑ Armando Meyer, Dale P. Sandler, Laura E. Beane Freeman, et alExposure and Risk of Rheumatoid Arthritis among Licensed Male Pesticide Applicators in the Agricultural Health Study Environ Health Perspect; DOI:10.1289/EHP1013
  37. ↑ Reuzel, P. G.; Bruijntjes, J. P.; Feron, V. J.; Woutersen, R. A. (1991). "Subchronic inhalation toxicity of amorphous silicas and quartz dust in rats". Food and Chemical Toxicology. 29 (5): 341–54. PMID 1648030. doi:10.1016/0278-6915(91)90205-L.
  38. ↑ "Worker Exposure to Silica during Countertop Manufacturing, Finishing and Installation" (PDF). National Institute for Occupational Safety and Health (NIOSH) and Occupational Safety and Health Administration (OSHA). February 2015. Retrieved 26 February 2015.
  39. ↑ Hornung, Veit; Bauernfeind, Franz; Halle, Annett; et al. (2008). "Silica crystals and aluminum salts activate the NALP3 inflammasome through phagosomal destabilization". Nature Immunology. 9 (8): 847–856. PMC 2834784 Freely accessible. PMID 18604214. doi:10.1038/ni.1631.
  40. ↑ NIOSH (1986) Occupational respiratory diseases. Cincinnati, OH: United States Department of Health and Human Services|U.S. Department of Health and Human Services]], U.S. Public Health Service, [[Centers for Disease Control and Prevention|Centers for Disease Control, National Institute for Occupational Safety and Health, DHHS (NIOSH) Publication No. 86-102.
  41. ↑ NIOSH (2002) Hazard Review, Health Effects of Occupational Exposure to Respirable Crystalline Silica. Cincinnati, OH: U.S. Department of Health and Human Services, U.S. Public Health Service, Centers for Disease Control, National Institute for Occupational Safety and Health, DHHS (NIOSH) Publication No. 2002-129.
  42. ↑ (PDF) https://www.osha.gov/OshDoc/data_General_Facts/crystalline-factsheet.pdf. Retrieved 3 August 2017. Missing or empty |title= (help)
  43. ↑ https://www.osha.gov/dsg/topics/silicacrystalline/. Retrieved 3 August 2017. Missing or empty |title= (help)
  44. ↑ http://www.silica-safe.org/ask-a-question/faq. Retrieved 3 August 2017. Missing or empty |title= (help)
  45. ↑ (PDF) http://www.ehs.uconn.edu/Word%20Docs/Silica%20fact%20sheet.pdf. Retrieved 3 August 2017. Missing or empty |title= (help)
  46. ↑ (PDF) http://osha.oregon.gov/OSHAPubs/3301.pdf. Retrieved 3 August 2017. Missing or empty |title= (help)
  47. ↑ Lager G. A.; Jorgensen J. D.; Rotella F.J. (1982). "Crystal structure and thermal expansion of a-quartz SiO2 at low temperature". Journal of Applied Physics. 53 (10): 6751–6756. Bibcode:1982JAP....53.6751L. doi:10.1063/1.330062.
  48. ↑ Wright, A. F.; Lehmann, M. S. (1981). "The structure of quartz at 25 and 590 °C determined by neutron diffraction". Journal of Solid State Chemistry. 36 (3): 371–80. Bibcode:1981JSSCh..36..371W. doi:10.1016/0022-4596(81)90449-7.
  49. 1 2 Kihara, Kuniaki; Matsumoto, Takeo; Imamura, Moritaka (1986). "Structural change of orthorhombic-Itridymite with temperature: A study based on second-order thermal-vibrational parameters". Zeitschrift für Kristallographie. 177: 27–38. Bibcode:1986ZK....177...27K. doi:10.1524/zkri.1986.177.1-2.27.
  50. ↑ Downs R. T.; Palmer D. C. (1994). "The pressure behavior of a cristobalite" (PDF). American Mineralogist. 79: 9–14.
  51. ↑ Wright, A. F.; Leadbetter, A. J. (1975). "The structures of the β-cristobalite phases of SiO2 and AlPO4". Philosophical Magazine. 31 (6): 1391–401. Bibcode:1975PMag...31.1391W. doi:10.1080/00318087508228690.
  52. ↑ Shropshire, Joseph; Keat, Paul P.; Vaughan, Philip A. (1959). "The crystal structure of keatite, a new form of silica". Zeitschrift für Kristallographie. 112: 409–13. Bibcode:1959ZK....112..409S. doi:10.1524/zkri.1959.112.1-6.409.
  53. ↑ Miehe, Gerhard; Graetsch, Heribert (1992). "Crystal structure of moganite: a new structure type for silica". European Journal of Mineralogy. 4 (4): 693–706. doi:10.1127/ejm/4/4/0693.
  54. ↑ Levien L.; Prewitt C. T. (1981). "High-pressure crystal structure and compressibility of coesite" (PDF). American Mineralogist. 66: 324–333.
  55. ↑ Smyth J. R.; Swope R. J.; Pawley A. R. (1995). "H in rutile-type compounds: II. Crystal chemistry of Al substitution in H-bearing stishovite" (PDF). American Mineralogist. 80: 454–456.
  56. ↑ Dera P.; Prewitt C. T.; Boctor N. Z.; Hemley R. J. (2002). "Characterization of a high-pressure phase of silica from the Martian meteorite Shergotty". American Mineralogist. 87: 1018.
  57. ↑ Seifertite. Mindat.org.
  58. ↑ Nakagawa T.; Kihara K.; Harada K. (2001). "The crystal structure of low melanophlogite". American Mineralogist. 86: 1506.
  59. ↑ Rosemarie Szostak (1998). Molecular sieves: Principles of Synthesis and Identification. Springer. ISBN 0-7514-0480-2.
  60. ↑ Weiss, Alarich; Weiss, Armin (1954). "Über Siliciumchalkogenide. VI. Zur Kenntnis der faserigen Siliciumdioxyd-Modifikation". Zeitschrift für anorganische und allgemeine Chemie. 276: 95–112. doi:10.1002/zaac.19542760110.
  61. ↑ Björkman, T; Kurasch, S; Lehtinen, O; Kotakoski, J; Yazyev, O. V.; Srivastava, A; Skakalova, V; Smet, J. H.; Kaiser, U; Krasheninnikov, A. V. (2013). "Defects in bilayer silica and graphene: common trends in diverse hexagonal two-dimensional systems". Scientific Reports. 3: 3482. Bibcode:2013NatSR...3E3482B. PMC 3863822 Freely accessible. PMID 24336488. doi:10.1038/srep03482.

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