Allotropes of oxygen

  (Redirected from Dioxygen)

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 oxygenEdit

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 exists naturally for a very short time. In outer space, the presence of ample 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]

DioxygenEdit

 
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, O
2
, 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, O
3
. 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 (O
2
) or the dioxygenyl ion (O+
2
).

The ground state of O
2
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 oxygenEdit

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.

OzoneEdit

Triatomic oxygen (ozone, O3) is a very reactive allotrope of oxygen that is a pale blue gas at standard temperature and pressure. Liquid and solid O3 have a deeper blue color than ordinary O2, and they are unstable and explosive.[7][8] In its gas phase, ozone is destructive to materials like rubber and fabric and is damaging to lung tissue.[9] Traces of it can be detected as a pungent, chlorine-like smell,[6] coming from electric motors, laser printers, and photocopiers, as it is formed whenever air is subjected to an electrical discharge. It was named "ozon" in 1840 by Christian Friedrich Schönbein,[10] from ancient Greek ὄζειν (ozein: "to smell") plus the suffix -on, commonly used at the time to designate a derived compound and anglicized as -one.[11]

Ozone is thermodynamically unstable and tends to react toward the more common dioxygen form. It is formed by reaction of intact O2 with atomic oxygen produced when UV radiation in the upper atmosphere splits O2.[7] Ozone absorbs strongly in the ultraviolet and in the stratosphere functions as a shield for the biosphere against mutagenic and other damaging effects of solar UV radiation (see ozone layer).[7] Tropospheric ozone is formed near the Earth's surface by the photochemical disintegration of nitrogen dioxide in the exhaust of automobiles.[12] 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.[13] The immune system produces ozone as an antimicrobial (see below).[14]

Cyclic ozoneEdit

Cyclic ozone is a theoretically predicted O3 molecule in which its three atoms of oxygen bond in an equilateral triangle instead of an open angle.

TetraoxygenEdit

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 O
4
was thought to be in one of the phases of solid oxygen later identified as O
8
. Cacace's team suggested that O
4
probably consists of two dumbbell-like O
2
molecules loosely held together by induced dipole dispersion forces.

Phases of solid oxygenEdit

There are six known distinct phases of solid oxygen. One of them is a dark-red O
8
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 selenium (exhibiting a pink-red color in its elemental state), tellurium and polonium, both of which show significant metallic character. At very low temperatures, this phase also becomes superconducting.

ReferencesEdit

  1. ^ a b Ryan D. McCulla, Saint Louis University (2010). "Atomic Oxygen O(3P): Photogeneration and Reactions with Biomolecules".
  2. ^ "Out of Thin Air" Archived 2017-06-23 at the Wayback Machine. NASA.gov. February 17, 2011.
  3. ^ Bell, Kassandra (6 May 2016). "Flying observatory detects atomic oxygen in Martian Atmosphere". NASA. Archived from the original on 8 November 2020. Retrieved 30 September 2021.
  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 16 December 2007.
  6. ^ a b Chemistry Tutorial : Allotropes Archived 2021-11-17 at the Wayback Machine from AUS-e-TUTE.com.au
  7. ^ a b c Mellor 1939
  8. ^ 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.
  9. ^ Stwertka 1998, p.48
  10. ^ Christian Friedrich Schönbein, Über die Erzeugung des Ozons auf chemischen Wege Archived 2020-06-30 at the Wayback Machine, p. 3, Basel: Schweighauser'sche Buchhandlung, 1844.
  11. ^ "Ozone", Oxford English Dictionary online, retrieved 29 June 2020.
  12. ^ Stwertka 1998, p.49
  13. ^ "Who is most at risk from ozone?". airnow.gov. Archived from the original on 17 January 2008. Retrieved 2008-01-06.
  14. ^ 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.
  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 readingEdit

  • Parks, G. D.; Mellor, J. W. (1939). Mellor's Modern Inorganic Chemistry (6th ed.). London: Longmans, Green and Co.
  • Stwertka, Albert (1998). Guide to the Elements (Revised ed.). Oxford University Press. ISBN 0-19-508083-1.