Silicon dioxide

Overview
The chemical compound silicon dioxide, also known as silica or silox (from the Latin "silex"), is an oxide of silicon, chemical formula, and has been known for its hardness since the 9th century. Silica is most commonly found in nature as sand or quartz, as well as in the cell walls of diatoms. It is a principal component of most types of glass and substances such as concrete.

Manufactured forms
Silica is manufactured in several forms including:
 * glass (a colorless, high-purity form is called fused silica)
 * synthetic amorphous silica
 * silica gel (used e.g. as desiccants in new clothes and leather goods)

It is used in the production of various products.
 * Inexpensive soda-lime glass is the most common and typically found in drinking glasses, bottles, and windows.
 * A raw material for many whiteware ceramics such as earthenware, stoneware and porcelain.
 * A raw material for the production of Portland cement.
 * A food additive, primarily as a flow agent in powdered foods, or to absorb water (see the ingredients list for).
 * The natural ("native") oxide coating that grows on silicon is hugely beneficial in microelectronics. It is a superior electric insulator, possessing high chemical stability. In electrical applications, it can protect the silicon, store charge, block current, and even act as a controlled pathway to allow small currents to flow through a device. At room temperature, however, it grows extremely slowly, and so to manufacture such oxide layers on silicon, the traditional method has been the deliberate heating of silicon in high temperature furnaces within an oxygen ambient (thermal oxidation).
 * Raw material for aerogel in the Stardust spacecraft
 * Used in the extraction of DNA and RNA due to its ability to bind to the nucleic acids under the presence of chaotropes.
 * Added to medicinal anti-foaming agent, like Simethicone, in a small proportion to enhance defoaming activity.
 * As hydrated silica in Toothpaste (abrasive to fight away plaque.)

Health effects
Inhaling finely divided crystalline silica dust in very small quantities (OSHA allows 0.1mg/m3) over time can lead to silicosis, bronchitis or (much more rarely) cancer, as the dust becomes lodged in the lungs and continuously irritates them, reducing lung capacities (silica does not dissolve over time). This effect can be an occupational hazard for people working with sandblasting equipment, products that contain powdered silica, and so on. But children, asthmatics of any age, allergy sufferers and the elderly, all of whom have reduced lung capacity, can be affected in much shorter periods of time.

In all other respects, silicon dioxide is inert and harmless. When silica is ingested orally, it passes unchanged through the gastrointestinal tract, exiting in the feces, leaving no trace behind. Small pieces of silicon dioxide are equally harmless, as long as they are not large enough to mechanically obstruct the GI tract, or jagged enough to lacerate its lining. Silicon dioxide produces no fumes and is insoluble in vivo. It is indigestible, with zero nutritional value and zero toxicity.

Chemistry
Silicon dioxide is formed when silicon is exposed to oxygen (or air). A very thin layer (approximately 1 nm or 10 Å) of so-called 'native oxide' is formed on the surface when silicon is exposed to air under ambient conditions.

Higher temperatures and alternate environments are used to grow well-controlled layers of silicon dioxide on silicon, for example at temperatures of 600 -1200 °C so-called "dry" or "wet" oxidation using O2 or H2O respectively. The thickness of the layer of silicon replaced by the dioxide is 44% of the thickness of the silicon dioxide layer produced.

Alternative methods used to deposit a layer of SiO2 include:
 * Low temperature oxidation (LTO) of silane
 * SiH4 + 2O2 → SiO2 + 2H2O (at 400-450 °C)


 * Decomposition of tetraethyl orthosilicate (TEOS) at 680 – 730°C


 * Si(OC2H5)4 → SiO2 + H2O + 2C2H4


 * Plasma enhanced chemical vapour deposition using TEOS at approximately 400°C
 * Si(OC2H5)4 + 12O2 → SiO2 + 10H2O + 8CO2

Pyrogenic silica(sometimes called fumed silica or silica fume), which is a very fine particulate form of silicon dioxide, is prepared by burning SiCl4 in an oxygen rich hydrocarbon flame to produce a "smoke" of SiO2:
 * SiCl4 + 2H2 + O2 → SiO2 + 4HCl

Quartz exhibits a maximum solubility in water at 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 which 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 sourcer of very pure quartz for use in electronic applications.

Fluorine reacts with silicon dioxide to form SiF4 and O2 whereas the other halogen gases (Cl2, Br2, I2) react much less readily.

Silicon dioxide is attacked by hydrofluoric acid (HF) to produce "hexafluorosilicic acid":
 * SiO2 + 6HF → H2SiF6 + 2H2O

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

Silicon dioxide dissolves in hot concentrated alkali or fused hydroxide (e.g):
 * SiO2 + NaOH → Na2SiO3 + H2O

Silicon dioxide reacts with basic metal oxides (e.g. sodium oxide, potassium oxide, lead(II) oxide, zinc oxide or mixtures of oxides forming silicates and glasses as the Si-O-Si bonds in silica are broken successively. As an example the reaction of sodium oxide and SiO2 can produce sodium orthosilicate, sodium silicate and glasses, depending on the proportions of reactants:
 * 2Na2O + SiO2 → Na4SiO4


 * Na2O + SiO2 → Na2SiO3


 * (0.25 - 0.8)Na2O + SiO2 → glasses

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.

With silicon at high temperatures gaseous SiO is produced:
 * SiO2 + Si → 2SiO (gas)

Structure and properties
SiO2 has a number of distinct crystalline forms in addition to amorphous forms. With the exception of stishovite and fibrous sulfur, 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 sulfur has a structure similar to that of SiS2 with chains of edge-sharing SiO4 tetrahedra.

Stishovite, the highest 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.

Note that the only stable form under normal conditions is α-quartz and this is the form in which crystalline silicon dioxide is usually encountered. In nature impurities in crystalline α-quartz can give rise to colours (see quartz for a list).

Molecular forms of silicon dioxide
When molecular silicon monoxide, SiO is condensed in an argon matrix cooled with helium along with oxygen atoms generated by microwave discharge molecular SiO2 is produced which has a linear structure. 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.

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 148.2pm.