Allotropes of oxygen
There are several known allotropes of oxygen:
- Free radicals O1 - unstable
- dioxygen, O2 - colorless
- ozone, O3 - blue
- tetraoxygen, O4 - metastable
- solid oxygen exists in 6 variously colored phases - of which one is O
8 and another one metallic
Dioxygen
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.
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 properties than triplet oxygen, including Diels-Alder reaction, or 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
Ozone
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 ὠζώ (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.[3][7]
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'.
Tetraoxygen
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
4 was thought to be in one of the phases of solid oxygen later identified as O
8. Cacace's team think that O
4 probably consists of two dumbbell-like O
2 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
8 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.
References
- ^ Chieh, Chung. "Bond Lengths and Energies". University of Waterloo. Retrieved 2007-12-16.
- ^ Stwertka 1998, p.48
- ^ a b c d Mellor 1939
- ^ Stwertka 1998, p.49
- ^ "Who is most at risk from ozone?". airnow.gov. Retrieved 2008-01-06.
- ^ 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): 2195–219. doi:10.1126/science.1077642. PMID 12434011.
{{cite journal}}:|access-date=requires|url=(help)CS1 maint: multiple names: authors list (link) - ^ 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.
- ^ Peter P. Edwards and 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. Retrieved 2007-12-16.
Further reading
- Parks, G. D. (1939). Mellor's Modern Inorganic Chemistry (6th ed.). London: Longmans, Green and Co.
{{cite book}}: Unknown parameter|coauthors=ignored (|author=suggested) (help) - Stwertka, Albert (1998). Guide to the Elements (Revised ed.). Oxford University Press. ISBN 0-19-508083-1.