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 SiO2 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 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.
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°.
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. 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. 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.
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 occuring 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.
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 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.
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. In many parts of the world, silica is the major constituent of sand.
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. 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).
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.
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.
Silica is used primarily in the production of glass, which is fused quartz. 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.
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.
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.
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.
In pharmaceutical products, silica aids powder flow when tablets are formed.
In cosmetics, it is useful for its light-diffusing properties and natural absorbency.
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.
A silica-based aerogel was used in the Stardust spacecraft to collect extraterrestrial particles.
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 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 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. Idealized equation involving a trisilicate and sulfuric acid is shown:
Approximately one billion kilograms/year (1999) of silica was 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 or wet oxidation with O2
The depth of the layer of silicon which dioxide replaces is 44% of the depth of the silicon dioxide layer produced.
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 silicate esters, the best known being tetraethyl orthosilicate (TEOS). Simply heating TEOS at 680–730 °C gives the dioxide:
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.
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.
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 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 milligram-per-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. It increases the risk of 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, 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.
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.
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. 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, called polymorphs.
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