Oxygen

For other uses, see Oxygen (disambiguation).

8 nitrogen ← oxygen → fluorine - ↑ O ↓ S periodic table
General
Name, Symbol, Number	oxygen, O, 8
Chemical series	Nonmetals, chalcogens
Group, Period, Block	16, 2, p
Appearance	colorless
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
Melting point	54.36 K (-218.79 °C, -361.82 °F)
Boiling point	90.20 K (-182.95 °C, -297.31 °F)
Critical point	154.59 K, 5.043 MPa
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)
Vapor pressure P/Pa 1 10 100 1 k 10 k 100 k at T/K       61 73 90
Atomic properties
Crystal structure	cubic
Oxidation states	−2, −1 (neutral oxide)
Electronegativity	3.44 (Pauling scale)
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 mW/(m·K)
Speed of sound	(gas, 27 °C) 330 m/s
CAS registry number	7782-44-7
Notable isotopes
Main article: Isotopes of oxygen iso NA half-life DM DE (MeV) DP 16O 99.76% O is stable with 8 neutrons 17O 0.038% O is stable with 9 neutrons 18O 0.21% O is stable with 10 neutrons

Oxygen is a chemical element in the periodic table. It has the symbol O and atomic number 8. Oxygen is the second most common element on Earth composing around 46% of the mass of Earth's crust and 28% of the mass of Earth as a whole, and is the third most common element in the universe. On Earth, it is usually covalently or ionically bonded to other elements. Unbound oxygen (usually called molecular dioxygen, O₂, a diatomic molecule) first appeared in significant quantities on Earth during the Paleoproterozoic era (between 2500 million years ago and 1600 million years ago) as a product of the metabolic action of early anaerobes (archaea and bacteria). According to most experts, this new presence of large amounts 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. Hydrochloric acid (HCl) does not contain oxygen.

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. Liquid and solid O₃ (ozone) have a deeper color 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 (to increase the inspired Oxygen partial pressure nearer to that found at sea-level requires increasing the proportion as a percentage of 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; a stable 50% gaseous mixture (Entonox) is commonly used in medicine today as an analgesic, and 30% oxygen with 70% Nitrous Oxide is the common basic anaesthetic mixture.

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), the second most common component of the Earth as a whole (28.2% by mass), and the second most common component of the Earth's atmosphere (20.947% by volume).

See also Silicate minerals, Oxide minerals.

Compounds

File:Oxygen -Molecular-.JPG

A common form of oxygen, O₂, which is a gas and consists of 2 oxygen atoms.

Ozone, a form of oxygen, O₃, which consists of 3 oxygen atoms.

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 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⁻. It was discovered when Neil Bartlett was studying the properties of PtF₆. He noticed a change in color when this compound was exposed to atmospheric air. Bartlett reasoned that xenon should also be oxidized by PtF₆. This led him to the discovery of xenon hexafluoroplatinate Xe⁺PtF₆⁻.

See also Oxygen compounds.

Isotopes

Oxygen has seventeen known isotopes with atomic masses ranging from 12.03 u to 28.06 u. Three are stable, ¹⁶O, ¹⁷O, and ¹⁸O, of which ¹⁶O is the most abundant (over 99.7%). The radioisotopes all have half-lives of less than three minutes.

An atomic weight of 16 was assigned to oxygen prior to the definition of the unified atomic mass unit based upon ¹²C. Since physicists referred to ¹⁶O only, while chemists meant the naturally abundant mixture of isotopes, this led to slightly different atomic weight scales.

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 compounds. 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. The fire that killed the Apollo 1 crew on a test launchpad spread so rapidly because the capsule was pressurized with pure oxygen as would be usual in an actual flight, but to maintain positive pressure in the capsule, this was at slightly more than atmospheric pressure instead of the 1/3 pressure that would be used in flight. (See partial pressure.) Similar hazards also apply to compounds of oxygen with a high oxidative potential, such as chlorates, perchlorates, and dichromates; they also can often cause chemical burns.

Oxygen derivatives are prone to form free radicals, especially in metabolic processes. Because they can cause severe damage to cells and their DNA, they form part of theories of carcinogenesis and aging.

See also

• Winkler test for dissolved oxygen for instructions on how to determine the amount of oxygen dissolved in fresh water.
• Combustion
• Oxidation
• Oxygen Catastrophe in geology
• The role of oxygen as a diving breathing gas
• Oxygen depletion aquatic ecology
• Ozone layer
• Oxygen isotope ratio cycle
• Aerobic

References

ISBN links support NWE through referral fees

• Los Alamos National Laboratory – Oxygen
• Nist atomic spectra database
• Nuclides and Isotopes Fourteenth Edition: Chart of the Nuclides, General Electric Company, 1989

External links

• Priestley Society, Dedicated to Joseph Priestley the man who discovered oxygen
• Joseph Priestley Information Website, about the man who discovered oxygen
• Los Alamos National Laboratory – Oxygen
• WebElements.com – Oxygen
• It's Elemental – Oxygen
• Oxygen (O2) Properties, Uses, Applications
• Computational Chemistry Wiki
• Oxidizing Agents > Oxygen

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