Oxygen
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| General | |||||||||||||||||||||||||
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| Name, Symbol, Number | oxygen, O, 8 | ||||||||||||||||||||||||
| Chemical series | Chalcogens | ||||||||||||||||||||||||
| Group, Period, Block | 16, 2, p | ||||||||||||||||||||||||
| Appearance | colorless |
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| Atomic mass | 15.9994 (3) g/mol | ||||||||||||||||||||||||
| Electron configuration | 1s2 2s2 2p4 | ||||||||||||||||||||||||
| Electrons per shell | 2, 6 | ||||||||||||||||||||||||
| Physical properties | |||||||||||||||||||||||||
| Phase | gas | ||||||||||||||||||||||||
| Density | (0 °C, 101.325
kPa) 1.429 g/L |
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| Melting point | 54.36
K (-218.79 ° C, -361.82 ° F) |
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| Boiling point | 90.20
K (-182.95 ° C, -297.31 ° F) |
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| Heat of fusion | (O2) 0.444 kJ/mol | ||||||||||||||||||||||||
| Heat of vaporization | (O2) 6.82 kJ/mol | ||||||||||||||||||||||||
| Heat capacity | (25 °C)
(O2) 29.378 J/(mol·K) |
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| Atomic properties | |||||||||||||||||||||||||
| Crystal structure | cubic | ||||||||||||||||||||||||
| Oxidation states | −2,
−1 (neutral oxide) |
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| Electronegativity | 3.44 ( Pauling scale) | ||||||||||||||||||||||||
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Ionization
energies ( more) |
1st: 1313.9 kJ/mol | ||||||||||||||||||||||||
| 2nd: 3388.3 kJ/mol | |||||||||||||||||||||||||
| 3rd: 5300.5 kJ/mol | |||||||||||||||||||||||||
| Atomic radius | 60 pm | ||||||||||||||||||||||||
| Atomic radius (calc.) | 48 pm | ||||||||||||||||||||||||
| Covalent radius | 73 pm | ||||||||||||||||||||||||
| Van der Waals radius | 152 pm | ||||||||||||||||||||||||
| Miscellaneous | |||||||||||||||||||||||||
| Magnetic ordering | paramagnetic | ||||||||||||||||||||||||
| Thermal conductivity | (300 K) 26.58 m W/(m·K) | ||||||||||||||||||||||||
| Speed of sound | (gas, 27 °C) 330 m/s | ||||||||||||||||||||||||
| CAS registry number | 7782-44-7 | ||||||||||||||||||||||||
| Notable isotopes | |||||||||||||||||||||||||
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| References | |||||||||||||||||||||||||
Oxygen is a chemical element in the periodic table. It has the symbol O and atomic number 8. The element is the most common on Earth, composing around 46% of the mass of Earth's crust, and is the third most common element in the universe, usually covalently bonded with other elements. Unbound oxygen (usually called molecular oxygen, O2, a diatomic molecule) first appeared on Earth during the Paleoproterozoic era (between 2500 million years ago and 1600 million years ago) and as a product of the metabolic action of early anaerobes ( archaea and bacteria). The presence of free oxygen drove most of the organisms then living to extinction. The atmospheric abundance of free oxygen in later geological epochs and up to the present has been largely driven by photosynthetic organisms, roughly three quarters by phytoplankton and algae in the oceans and one quarter from terrestrial plants.
Characteristics
At standard temperature and pressure, oxygen exists as a diatomic molecule with the formula O2, in which the two oxygen atoms are doubly bonded to each other. In its most stable form, oxygen exists as a diradical ( triplet oxygen). Though radicals are commonly associated with highly reactive compounds, triplet oxygen is surprisingly (and fortunately) unreactive towards most compounds. Singlet oxygen, a name given to several higher energy species in which all the electron spins are paired, is much more reactive towards common organic molecules. Carotenoids effectively absorb energy from singlet oxygen and convert it back into the unexcited ground state.
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, οξυς (oxys) (acid, sharp) and γεινομαι (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 colour and both are highly paramagnetic. Liquid O2 is usually obtained by the fractional distillation of liquid air. Liquid and solid O3 ( ozone) have a deeper colour of blue.
A 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
Liquid oxygen finds use as an oxidizer in rocket propulsion. Oxygen is essential to respiration, so oxygen supplementation has found use in medicine (as oxygen therapy). People who climb mountains or fly in airplanes sometimes have supplemental oxygen supplies (as air). Oxygen is used in welding (such as the oxyacetylene torch), and in the making of steel and methanol.
Oxygen presents two absorption bands centered in the wavelengths 687 and 760 nanometers. Some scientists have proposed to use the measurement of the radiance coming from vegetation canopies in those oxygen bands to characterize plant health status from a satellite platform. This is because in those bands, it is possible to discriminate the vegetation's reflectance from the vegetation's fluorescence, which is much weaker. The measurement presents several technical difficulties due to the low signal to noise ratio and due to the vegetation's architecture, but it has been proposed as a possibility to monitor the carbon cycle from satellites on a global scale.
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 an analgesic effect; indeed, such a mixture ( Entonox) is commonly used in medicine today.
History
Oxygen was first discovered by Michał Sędziwój, Polish alchemist and philosopher in late 16th century. Sędziwój assumed the existence of oxygen by warming nitre (saltpeter). 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, who called the gas dephlogisticated air (see phlogiston theory). Priestley published his discoveries 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 most common component of the Earth's crust (46.6% by mass) and the second most common component of the Earth's atmosphere (20.947% by volume).
Compounds
Due to its electronegativity, oxygen forms chemical bonds with almost all other elements hence 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 monoxide, 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 and is found as a minor component of liquid oxygen. Epoxides are ethers in which the oxygen atom is part of a ring of three atoms.
One unexpected oxygen compound is dioxygen hexafluoroplatinate O2+PtF6−, which resulted when someone tried to make neutral PtF6 in the presence of atmospheric air. This led to xenon hexafluoroplatinate being made.
Isotopes
Oxygen has fifteen known isotopes with atomic masses ranging from 12 to 26. Three of them are stable and twelve are radioactive. The radioisotopes all have half lives of less than three minutes. The stable isotopes have mass numbers of 16, 17 and 18, of which 16O is the most common (over 99%).
Precautions
Oxygen can be toxic at elevated partial pressures (i.e. high relative concentrations). This is important in some forms of scuba diving, such as with a rebreather.
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. To protect against the destructive nature of peroxides, nearly every organism on earth has developed some form of the enzyme catalase, which very quickly disproportionates peroxide into water and dioxygen. 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.