Molecular Oxygen
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Molecular Oxygen

There are several known allotropes of oxygen. The most familiar is molecular oxygen (O2), present at significant levels in Earth's atmosphere and also known as dioxygen or triplet oxygen. Another is the highly reactive ozone (O3). Others are:

Atomic oxygen

Atomic oxygen, denoted O(3P) or O(3P),[1] is very reactive, as the single atoms of oxygen tend to quickly bond with nearby molecules. On Earth's surface, it does not exist naturally for very long, but in outer space, the presence of plenty of ultraviolet radiation results in a low Earth orbit atmosphere in which 96% of the oxygen occurs in atomic form.[1][2]

Atomic oxygen has been detected on Mars by Mariner, Viking, and the SOFIA observatory.[3]


The most commonly encountered allotrope of elemental oxygen is triplet dioxygen, a diradical. The unpaired electrons participate in three-electron bonding, shown here using dashed lines.

The common allotrope of elemental oxygen on Earth, , is generally known as oxygen, but may be called dioxygen, diatomic oxygen, molecular oxygen, or oxygen gas to distinguish it from the element itself and from the triatomic allotrope ozone, . As a major component (about 21% by volume) of Earth's atmosphere, elemental oxygen is most commonly encountered in the diatomic form. Aerobic organisms release the chemical energy stored in the weak sigma bond of atmospheric dioxygen, the terminal oxidant in cellular respiration.[4] The ground state of dioxygen is known as triplet oxygen, 3O2, because it has two unpaired electrons. The first excited state, singlet oxygen, 1O2, has no unpaired electrons and is metastable. The doublet state requires an odd number of electrons, and so cannot occur in dioxygen without gaining or losing electrons, such as in the superoxide ion or the dioxygenyl ion .

The ground state of has a bond length of 121 pm and a bond energy of 498 kJ/mol.[5] It is a colourless gas with a boiling point of -183 °C (90 K; -297 °F).[6] It can be condensed from air by cooling with liquid nitrogen, which has a boiling point of -196 °C (77 K; -321 °F). Liquid oxygen is pale blue in colour, and is quite markedly paramagnetic due to the unpaired electrons; liquid oxygen contained in a flask suspended by a string is attracted to a magnet.

Singlet oxygen

Singlet oxygen is the common name used for the two metastable states of molecular oxygen (O2) with higher energy than the ground state triplet oxygen. Because of the differences in their electron shells, singlet oxygen has different chemical and physical properties than triplet oxygen, including absorbing and emitting light at different wavelengths. It can be generated in a photosensitized process by energy transfer from dye molecules such as rose bengal, methylene blue or porphyrins, or by chemical processes such as spontaneous decomposition of hydrogen trioxide in water or the reaction of hydrogen peroxide with hypochlorite.


Triatomic oxygen (ozone, O3), is a very reactive allotrope of oxygen that is destructive to materials like rubber and fabrics and is also damaging to lung tissue.[7] Traces of it can be detected as a sharp, chlorine-like smell,[6] coming from electric motors, laser printers, and photocopiers. It was named "ozon" in 1840 by Christian Friedrich Schönbein,[8] from ancient Greek (ozein: "to smell") plus the suffix -on (in English -one) commonly used at the time to designate a derived compound.[9]

Ozone is thermodynamically unstable toward the more common dioxygen form, and is formed by reaction of O2 with atomic oxygen produced by splitting of O2 by UV radiation in the upper atmosphere.[10] Ozone absorbs strongly in the ultraviolet and functions as a shield for the biosphere against the mutagenic and other damaging effects of solar UV radiation (see ozone layer).[10] Ozone is formed near the Earth's surface by the photochemical disintegration of nitrogen dioxide from the exhaust of automobiles.[11] Ground-level ozone is an air pollutant that is especially harmful for senior citizens, children, and people with heart and lung conditions such as emphysema, bronchitis, and asthma.[12] The immune system produces ozone as an antimicrobial (see below).[13] Liquid and solid O3 have a deeper blue color than ordinary oxygen and they are unstable and explosive.[10][14]

Ozone is a pale blue gas condensable to a dark blue liquid. It is formed whenever air is subjected to an electrical discharge, and has the characteristic pungent odour of new-mown hay or subways - the so-called 'electrical odour'.

Cyclic ozone


Tetraoxygen had been suspected to exist since the early 1900s, when it was known as oxozone. It was identified in 2001 by a team led by Fulvio Cacace at the University of Rome.[15] The molecule was thought to be in one of the phases of solid oxygen later identified as . Cacace's team suggested that probably consists of two dumbbell-like molecules loosely held together by induced dipole dispersion forces.

Phases of solid oxygen

There are six known distinct phases of solid oxygen. One of them is a dark-red cluster. When oxygen is subjected to a pressure of 96 GPa, it becomes metallic, in a similar manner to hydrogen,[16] and becomes more similar to the heavier chalcogens, such as tellurium and polonium, both of which show significant metallic character. At very low temperatures, this phase also becomes superconducting.


  1. ^ a b Ryan D. McCulla, Saint Louis University (2010). "Atomic Oxygen O(3P): Photogeneration and Reactions with Biomolecules".
  2. ^ "Out of Thin Air". February 17, 2011.
  3. ^ [1]
  4. ^ Schmidt-Rohr, Klaus (2020). "Oxygen is the High-Energy Molecule Powering Complex Multicellular Life: Fundamental Corrections to Traditional Bioenergetics". ACS Omega. 5 (5): 2221-2233. doi:10.1021/acsomega.9b03352. PMC 7016920. PMID 32064383.
  5. ^ Chieh, Chung. "Bond Lengths and Energies". University of Waterloo. Archived from the original on 14 December 2007. Retrieved 2007.
  6. ^ a b Chemistry Tutorial : Allotropes from
  7. ^ Stwertka 1998, p.48
  8. ^ Christian Friedrich Schönbein, Über die Erzeugung des Ozons auf chemischen Wege, p. 3, Basel: Schweighauser'sche Buchhandlung, 1844.
  9. ^ "ozone", Oxford English Dictionary online, retrieved 29 June 2020.
  10. ^ a b c Mellor 1939
  11. ^ Stwertka 1998, p.49
  12. ^ "Who is most at risk from ozone?". Archived from the original on 17 January 2008. Retrieved .
  13. ^ Paul Wentworth Jr.; Jonathan E. McDunn; Anita D. Wentworth; Cindy Takeuchi; Jorge Nieva; Teresa Jones; Cristina Bautista; Julie M. Ruedi; Abel Gutierrez; Kim D. Janda; Bernard M. Babior; Albert Eschenmoser; Richard A. Lerner (2002-12-13). "Evidence for Antibody-Catalyzed Ozone Formation in Bacterial Killing and Inflammation". Science. 298 (5601): 2195-2199. Bibcode:2002Sci...298.2195W. doi:10.1126/science.1077642. PMID 12434011. S2CID 36537588.
  14. ^ Cotton, F. Albert and Wilkinson, Geoffrey (1972). Advanced Inorganic Chemistry: A comprehensive Text. (3rd Edition). New York, London, Sydney, Toronto: Interscience Publications. ISBN 0-471-17560-9.
  15. ^ Cacace, Fulvio (2001). "Experimental Detection of Tetraoxygen". Angewandte Chemie International Edition. 40 (21): 4062-4065. doi:10.1002/1521-3773(20011105)40:21<4062::AID-ANIE4062>3.0.CO;2-X. PMID 12404493.
  16. ^ Peter P. Edwards; Friedrich Hensel (2002-01-14). "Metallic Oxygen". ChemPhysChem. 3 (1): 53-56. doi:10.1002/1439-7641(20020118)3:1<53::AID-CPHC53>3.0.CO;2-2. PMID 12465476.

Further reading

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