Allotropes of oxygen

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There are several known allotropes of oxygen:


The common allotrope of elemental oxygen on Earth, O2, is known as dioxygen. Elemental oxygen is most commonly encountered in this form, as about 21% (by volume) of Earth's atmosphere. O2 has a bond length of 121 pm and a bond energy of 498 kJ/mol.[1]

Oxygen itself is a colourless gas with a boiling point of -183°C. It can be condensed out of air by cooling with liquid nitrogen, which has a boiling point of -196°C. Liquid oxygen is pale blue in colour, and is quite markedly paramagnetic : liquid oxygen contained in a flask suspended by a string is attracted to a magnet.


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.[2] Traces of it can be detected as a sharp, chlorine-like smell coming from electric motors, laser printers, and photocopiers. It was named "ozone" by Christian Friedrich Schönbein, in 1840, from the Greek word ÖĮώ (ozo) for smell.[3]

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.[3] 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).[3] Ozone is formed near the earth's surface by the photochemical disintegration of nitrogen dioxide from the exhaust of automobiles.[4] 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.[5] The immune system produces ozone as an antimicrobial (see below).[6] Liquid and solid O3 have a deeper-blue color than ordinary oxygen and they are unstable and explosive.[7][3]

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 for those living in urban environments, of subways - the so-called 'electrical odour'.

Electrical discharges cause dioxygen to split into oxygen radicals. Most of these recombine to form dioxygen, but a few react with dioxygen to give ozone:

O2 + O· → O3

The ozone molecules themselves can also react with oxygen free radicals, to reform dioxygen, and so the actual concentration of atmospheric ozone is quite small. It is believed that ozone is formed in the upper atmosphere by the photodissociation of dioxygen by the intense ultraviolet radiation from the sun. This light energy is thus absorbed, otherwise it would reach the Earth and destroy all life quite rapidly. Ozone is a greenhouse gas and, as such, would contribute to global warming if present in the lower atmosphere.


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

Phases of solid oxygen

There are 6 known distinct phases of solid oxygen. One of them is a dark-red O
cluster. When oxygen is subjected to a pressure of 96 GPa, it becomes metallic, in a similar manner as hydrogen,[8] 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. Chieh, Chung. "Bond Lengths and Energies". University of Waterloo. Retrieved 2007-12-16.
  2. Stwertka 1998, p.48
  3. 3.0 3.1 3.2 3.3 Mellor 1939
  4. Stwertka 1998, p.49
  5. "Who is most at risk from ozone?". Retrieved 2008-01-06.
  6. 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: 2195 - 219. doi:10.1126/science.1077642. Retrieved on 2007-12-16. 
  7. 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.
  8. Peter P. Edwards and Friedrich Hensel (2002-01-14). "Metallic Oxygen". ChemPhysChem 3 (1): 53 - 56. Retrieved on 2007-12-16. 


  • Parks, G. D. (1939). Mellor's Modern Inorganic Chemistry (6th edition ed.). London: Longmans, Green and Co. Unknown parameter |coauthors= ignored (help)
  • Stwertka, Albert (1998). Guide to the Elements (Revised Edition ed.). Oxford University Press. ISBN 0-19-508083-1.