Oxygen

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Oxygen is the chemical element in the periodic table that has the symbol O and atomic number 8. The element is very common, found not only on Earth but throughout the universe, usually bound with other elements. Unbound oxygen (usually called molecular oxygen, O₂) made its initial appearance on Earth as a product of the metabolic action of early anaerobes (archaea and bacteria). The atmospheric abundance of free oxygen in later geological epochs and up to the present has been largely driven by terrestrial plants, which release oxygen during photosynthesis.

Contents

1 Characteristics 2 Applications 3 History 4 Occurrence 5 Compounds 6 Isotopes 7 Precautions

Characteristics

At standard temperature and pressure, oxygen is predominantly found as a gas consisting of a diatomic molecule with the chemical formula O₂. O₂ itself has two energetic forms: the low-energy, predominant single-bonded diradical triplet oxygen, and the high-energy double-bonded molecule singlet oxygen. This native diradical quality of oxygen contributes to its destructive chemical nature.

Oxygen is a major component of air, produced by plants during photosynthesis, and is necessary for aerobic respiration in animals. The word oxygen derives from two words in Greek, oxys (acid) and geinomai (engender). The name "oxygen" was chosen because, at the time it was discovered in the late 18th century, it was believed that all acids contained oxygen. The definition of acid has since been revised to not require oxygen in the molecular structure.

Liquid O₂ and solid O₂ have a light blue color and both are highly paramagnetic. Liquid O₂ is usually obtained by the fractional distillation of liquid air. Both liquid and solid O₃ (ozone) have a deeper color of blue.

Another recently discovered allotrope of oxygen, tetraoxygen (O₄), is a deep red solid that is created by pressurizing O₂ to the order of 20 GPa. Its properties are being studied for use in rocket fuels and similar applications, as it is a much more powerful oxidizer than either O₂ or O₃.

Applications

Oxygen finds considerable use as an oxidizer, with only fluorine having a higher electronegativity. Liquid oxygen finds use as an oxidizer in rocket propulsion. Oxygen is essential to respiration, so oxygen supplementation has found use in medicine. People who climb mountains or fly in airplanes sometimes have supplemental oxygen supplies (as air). Oxygen is used in welding, and in the making of steel and methanol.

Oxygen, as a mild euphoric, has a history of recreational use that extends into modern times. Oxygen bars can be seen at parties to this day. In the 19th century, oxygen was often mixed with nitrous oxide to promote a kind of analgesic effect; indeed such a mixture (Entonox) is commonly used in medicine today.

History

Oxygen was first discovered by [[Michal Sedziw󪼍ichał Sędziw󪝝, Polish alchemist and philosopher in late 16th century. Sędziw󪠡ssumed the existence of oxygen by warming nitre (saltpetre). He thought of the gas given off as "the elixir of life".

Oxygen was again discovered by the Swedish pharmacist Carl Wilhelm Scheele sometime before 1773, but the discovery was not published until after the independent discovery by Joseph Priestley on August 1 1774. Priestley published his findings in 1775 and Scheele in 1777; consequently Priestley is usually given the credit. It was named by Antoine Laurent Lavoisier after Priestley's publication in 1775.

Occurrence

Oxygen is the second largest single component of the Earth's atmosphere (20.947% by volume).

Compounds

Due to its electronegativity, oxygen forms chemical bonds with almost all other elements (which is the origin of the original definition of oxidation). The only elements to escape the possibility of oxidation are a few of the noble gases. The most famous of these oxides is dihydrogen oxide, or water (H₂O). Other well known examples include compounds of carbon and oxygen, such as carbon dioxide (CO₂), alcohols (R-OH), aldehydes, (R-CHO), and carboxylic acids (R-COOH). Oxygenated radicals such as chlorates (ClO₃⁻), perchlorates (ClO₄⁻), chromates (CrO₄²⁻), dichromates (Cr₂O₇²⁻), permanganates (MnO₄⁻), and nitrates (NO₃⁻)are strong oxidizing agents in and of themselves. Many metals such as Iron bond with oxygen atoms, iron (III) oxide (Fe₂O₃). Ozone (O₃) is formed by electrostatic discharge in the presence of molecular oxygen. A double oxygen molecule (O₂)₂ is known, found as a minor component of liquid oxygen. Epoxides are ethers in which the oxygen atom is part of a ring of three atoms.

Isotopes

Oxygen has three stable isotopes and ten known radioactive isotopes. The radioisotopes all have half lives of less than three minutes.

Precautions

Oxygen can be toxic at elevated partial pressures.

Certain derivatives of oxygen, such as ozone (O₃), singlet oxygen, hydrogen peroxide, hydroxyl radicals and superoxide, are also highly toxic. The body has developed mechanisms to protect against these toxic species. For instance, the naturally-occurring glutathione can act as an antioxidant, as can bilirubin which is normally a breakdown product of hemoglobin. Highly concentrated sources of oxygen promote rapid combustion and therefore are fire and explosion hazards in the presence of fuels. This is true as well of compounds of oxygen such as chlorates, perchlorates, dichromates, etc. Compounds with a high oxidative potential can often cause chemical burns.

The fire that killed the Apollo 1 crew on a test launchpad spread so rapidly because the pure oxygen atmosphere was at normal atmospheric pressure instead of the one third pressure that would be used during an actual launch. (See partial pressure.)

Oxygen derivatives are prone to form free radicals, especially in metabolic processes. Because they can cause severe damage to cells and their DNA, they are thought to be related to cancer and aging.

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