Oxygen

Oxygen, ₈O
	Liquid oxygen (O2 at below −183 °C)
Oxygen
Allotropes	O2, O3 (ozone) and more (see Allotropes of oxygen)
Appearance	gas: colorless; liquid and solid: pale blue
Standard atomic weight Ar°(O)
	[15.99903, 15.99977]; 15.999±0.001 (abridged);
Abundance
in the Earth's crust	461000 ppm
Oxygen in the periodic table
	nitrogen ← oxygen → fluorine
Atomic number (Z)	8
Group	group 16 (chalcogens)
Period	period 2
Block	p-block
Electron configuration	[He] 2s2 2p4
Electrons per shell	2, 6
Physical properties
Phase at STP	gas
Melting point	(O2) 54.36 K (−218.79 °C, −361.82 °F)
Boiling point	(O2) 90.188 K (−182.962 °C, −297.332 °F)
Density (at STP)	1.429 g/L
when liquid (at b.p.)	1.141 g/cm3
Triple point	54.361 K, 0.1463 kPa
Critical point	154.581 K, 5.043 MPa
Heat of fusion	(O2) 0.444 kJ/mol
Heat of vaporization	(O2) 6.82 kJ/mol
Molar heat capacity	(O2) 29.378 J/(mol·K)
Atomic properties
Oxidation states	common: −2; −1, 0, +1, +2
Electronegativity	Pauling scale: 3.44
Ionization energies	1st: 1313.9 kJ/mol ; 2nd: 3388.3 kJ/mol ; 3rd: 5300.5 kJ/mol ; (more) ;
Covalent radius	66±2 pm
Van der Waals radius	152 pm
	Spectral lines of oxygen
Other properties
Natural occurrence	primordial
Crystal structure	cubic (cP16)
Lattice constant	a = 678.28 pm (at t.p.)
Thermal conductivity	26.58×10−3 W/(m⋅K)
Magnetic ordering	paramagnetic
Molar magnetic susceptibility	+3449.0×10−6 cm3/mol (293 K)
Speed of sound	330 m/s (gas, at 27 °C)
CAS Number	7782-44-7
History
Discovery	Michael Sendivogius; Carl Wilhelm Scheele (1604, 1771)
Named by	Antoine Lavoisier (1777)
Isotopes of oxygen v; e;
	Category: Oxygen; view; talk; edit; \| references

Oxygen is a chemical element with the symbol O and atomic number 8. It is a member of the chalcogen group in the periodic table, a highly reactive nonmetal, and a potent oxidizing agent that readily forms oxides with most elements as well as with other compounds. Oxygen is the most abundant element in Earth's crust, and the third-most abundant element in the universe after hydrogen and helium.

At standard temperature and pressure, two oxygen atoms will bind covalently to form dioxygen, a colorless and odorless diatomic gas with the chemical formula O
₂. Dioxygen gas currently constitutes 20.95% molar fraction of the Earth's atmosphere, though this has changed considerably over long periods of time in Earth's history. Oxygen makes up almost half of the Earth's crust in the form of various oxides such as water, carbon dioxide, iron oxides and silicates.^[6]

All eukaryotic organisms, including plants, animals, fungi, algae and most protists, need oxygen for cellular respiration, which extracts chemical energy by the reaction of oxygen with organic molecules derived from food and releases carbon dioxide as a waste product. In aquatic animals, dissolved oxygen in water is absorbed by specialized respiratory organs called gills, through the skin or via the gut; in terrestrial animals such as tetrapods, oxygen in air is actively taken into the body via specialized organs known as lungs, where gas exchange takes place to diffuse oxygen into the blood and carbon dioxide out, and the body's circulatory system then transports the oxygen to other tissues where cellular respiration takes place.^[7]^[8] However in insects, the most successful and biodiverse terrestrial clade, oxygen is directly conducted to the internal tissues via a deep network of airways.

Many major classes of organic molecules in living organisms contain oxygen atoms, such as proteins, nucleic acids, carbohydrates and fats, as do the major constituent inorganic compounds of animal shells, teeth, and bone. Most of the mass of living organisms is oxygen as a component of water, the major constituent of lifeforms. Oxygen in Earth's atmosphere is produced by biotic photosynthesis, in which photon energy in sunlight is captured by chlorophyll to split water molecules and then react with carbon dioxide to produce carbohydrates and oxygen is released as a byproduct. Oxygen is too chemically reactive to remain a free element in air without being continuously replenished by the photosynthetic activities of autotrophs such as cyanobacteria, chloroplast-bearing algae and plants. A much rarer triatomic allotrope of oxygen, ozone (O
₃), strongly absorbs the UVB and UVC wavelengths and forms a protective ozone layer at the lower stratosphere, which shields the biosphere from ionizing ultraviolet radiation. However, ozone present at the surface is a corrosive byproduct of smog and thus an air pollutant.

Oxygen was isolated by Michael Sendivogius before 1604, but it is commonly believed that the element was discovered independently by Carl Wilhelm Scheele, in Uppsala, in 1773 or earlier, and Joseph Priestley in Wiltshire, in 1774. Priority is often given for Priestley because his work was published first. Priestley, however, called oxygen "dephlogisticated air", and did not recognize it as a chemical element. The name oxygen was coined in 1777 by Antoine Lavoisier, who first recognized oxygen as a chemical element and correctly characterized the role it plays in combustion.

Common industrial uses of oxygen include production of steel, plastics and textiles, brazing, welding and cutting of steels and other metals, rocket propellant, oxygen therapy, and life support systems in aircraft, submarines, spaceflight and diving.

^ "Standard Atomic Weights: Oxygen". CIAAW. 2009.
^ Prohaska, Thomas; Irrgeher, Johanna; Benefield, Jacqueline; Böhlke, John K.; Chesson, Lesley A.; Coplen, Tyler B.; Ding, Tiping; Dunn, Philip J. H.; Gröning, Manfred; Holden, Norman E.; Meijer, Harro A. J. (May 4, 2022). "Standard atomic weights of the elements 2021 (IUPAC Technical Report)". Pure and Applied Chemistry. doi:10.1515/pac-2019-0603. ISSN 1365-3075.
^ ^a ^b ^c Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. p. 28. ISBN 978-0-08-037941-8.
^ Arblaster, John W. (2018). Selected Values of the Crystallographic Properties of Elements. Materials Park, Ohio: ASM International. ISBN 978-1-62708-155-9.
^ Weast, Robert (1984). CRC, Handbook of Chemistry and Physics. Boca Raton, Florida: Chemical Rubber Company Publishing. pp. E110. ISBN 0-8493-0464-4.
^ Atkins, P.; Jones, L.; Laverman, L. (2016).Chemical Principles, 7th edition. Freeman. ISBN 978-1-4641-8395-9
^ Hall, John (2011). Guyton and Hall textbook of medical physiology (12th ed.). Philadelphia, Pa.: Saunders/Elsevier. p. 5. ISBN 978-1-4160-4574-8.
^ Pocock, Gillian; Richards, Christopher D. (2006). Human physiology : the basis of medicine (3rd ed.). Oxford: Oxford University Press. p. 311. ISBN 978-0-19-856878-0.

[1] "Standard Atomic Weights: Oxygen". CIAAW. 2009.

[CIAAW2021-2] Prohaska, Thomas; Irrgeher, Johanna; Benefield, Jacqueline; Böhlke, John K.; Chesson, Lesley A.; Coplen, Tyler B.; Ding, Tiping; Dunn, Philip J. H.; Gröning, Manfred; Holden, Norman E.; Meijer, Harro A. J. (May 4, 2022). "Standard atomic weights of the elements 2021 (IUPAC Technical Report)". Pure and Applied Chemistry. doi:10.1515/pac-2019-0603. ISSN 1365-3075.

[GE28-3] Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. p. 28. ISBN 978-0-08-037941-8.

[Arblaster_2018-4] Arblaster, John W. (2018). Selected Values of the Crystallographic Properties of Elements. Materials Park, Ohio: ASM International. ISBN 978-1-62708-155-9.

[5] Weast, Robert (1984). CRC, Handbook of Chemistry and Physics. Boca Raton, Florida: Chemical Rubber Company Publishing. pp. E110. ISBN 0-8493-0464-4.

[Atkins7th-6] Atkins, P.; Jones, L.; Laverman, L. (2016).Chemical Principles, 7th edition. Freeman. ISBN 978-1-4641-8395-9

[7] Hall, John (2011). Guyton and Hall textbook of medical physiology (12th ed.). Philadelphia, Pa.: Saunders/Elsevier. p. 5. ISBN 978-1-4160-4574-8.

[Pocock2-8] Pocock, Gillian; Richards, Christopher D. (2006). Human physiology : the basis of medicine (3rd ed.). Oxford: Oxford University Press. p. 311. ISBN 978-0-19-856878-0.

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