Oxygen

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nitrogen ← oxygen → fluorine
–
↑
O
↓
S
Atomic properties
Atomic number 8
Standard atomic weight 15.9994(3)
Electron configuration [He] 2s2 2p4
Physical properties (O2)[1][2]
Melting point 54.8(2) K (−218.8 °C)
Boiling point 90.2(2) K (−183.0 °C)
Triple point 54.35 K, 1.52 mbar
Critical point 154.58 K, 50.43 bar
Density 1.354 kg m−3 (1 atm, 15 °C)
4.475 kg m−3 (1 atm, 90.2 K)
1.141 g cm−3 (l, 90.2 K)
Chemical properties[3]
Electronegativity 3.44 (Pauling)
Solubility in water 48.9 cm3 dm−3 (1 atm, 0 °C)
Ionization energies[4][5]
1st 13.618 06 eV,
1313.943 kJ mol−1
2nd 35.1211 eV,
3388.67 kJ mol−1
3rd 54.9355 eV,
5300.47 kJ mol−1
4th 77.4135 eV,
7469.27 kJ mol−1
5th 113.8989 eV,
10 989.57 kJ mol−1
6th 138.1196 eV,
13 326.52 kJ mol−1
7th 739.3268 eV,
71 334.20 kJ mol−1
8th 871.4097 eV,
84 078.26 kJ mol−1
Total 2043.8432 eV,
197 200.9 kJ mol−1
Electron affinity[4]
140.9759(42) kJ mol−1
Atomic radii [6][7][8]
Covalent radius 66(2) pm
Ionic radius 140 pm (O2−, Oh)
Van der Waals radius 152 pm
Thermodynamic properties (O2)[1][9]
Standard entropy 205.152(5) J K−1 mol−1
Enthalpy change of atomization 249.18(10) kJ mol−1
Entropy change of atomization −44.093(6) J K−1 mol−1
Enthalpy change of fusion 0.444 kJ mol−1
Enthalpy change of vaporization 6.82 kJ mol−1
Molar heat capacity (Cp) 29.378 J K−1 mol−1
Miscellaneous
CAS number 7782-44-7 (O2)
17778-80-2 (atomic)
EC number 231-956-9
Where appropriate, and unless otherwise stated, data are given for 100 kPa (1 bar) and 298.15 K (25 °C).
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Oxygen is a colourless gas which makes up about one fifth of the Earth's atmosphere.

History

The discovery of oxygen is often credited to English chemist Joseph Priestley, although the full story is somewhat more involved. That there is a component of air which is necessary for combustion and respiration was recognized by Leonardo da Vinci in the fifteenth century, and confirmed by English chemist John Mayow in the mid-seventeenth century.[10] However, the interpretation of these results was hampered by the rise of phlogiston theory, which stated that substances gave off phlogiston during combustion: air that was no longer capable of supporting combustion was said to be saturated with phlogiston. The unequivocal identification of oxygen as a chemical substance would have to wait for its preparation by chemical means.

The preparation was first carried out by Swedish chemist Carl Wilhelm Scheele on several occasions during the period 1771–73. Scheele heated various compounds such as KNO3, Mg(NO3)2 and HgO and found that they gave off a gas that he named "vitriol air", which supported combustion better than normal air.[11] Scheele's results, however, were not published until 1777.[12] In the meantime, Priestley had isolated the gas given off by heating HgO and named it "dephlogisticated air", and published his results in 1775 after proving that the gas was different from nitrous oxide.[13]

Priestley's work certainly had the greater impact, as he was able to discuss it with French chemist Antoine Lavoisier in October 1774 during a visit to Paris with his mentor and employer the Earl of Shelbourne. Lavoisier repeated and extended Priestley's work, and published his results in 1777 as Mémoire sur la combustion en général[14] and Considérations générales sur la nature des acides et sur les principes dont ils sont composés.[15] It was Lavoisier who proposed the name "oxygen" for the new gas, from the Greek ὀξύς (oxys; acid, literally "sharp", from the taste of acids) and -γενής (-genēs; producer, literally "begetter"), as he (incorrectly) believed that all acids contained oxygen.

Occurance and production

Oxygen is almost ubiquitous at the surface of the Earth. The approximate mass fractions of oxygen are: crustal rocks 46%; the human body 61%; sea water 86%. The amount fraction of oxygen in the Earth's atmosphere (at sea level) is 20.9476%,[16] amounting to some 1015 tonnes. Apart from the atmosphere, the vast majority of this oxygen is chemically combined in a wide variety of inorganic and organic compounds: in the atmosphere, oxygen exists almost entirely as dioxygen (O2) molecules, its normal elemental state, with a small (but very important) amount of ozone (O3).

While it is difficult to obtain precise statistics, oxygen is believed to be the third most important bulk industrial chemical, after sulfuric acid and lime but ahead of ammonia and nitrogen, with production of at least 100 million tonnes per year. It is produced by the fractional distillation of liquid air. As the transport costs for the raw material (air) are zero, it is often economic to locate the oxygen plant close to the consumer, even for relatively small consumers such as hospitals.

Use

The majority of oxygen produced (ca. 70%) is used in the basic oxygen process for steelmaking.[11] A streem of pure oxygen is blown into the molton iron, where it burns off the residual carbon to produce low-carbon (0.1–1.0%) steel. Another large scale use is in oil refining, particularly in fluid catalytic cracking (FCC) units that treat roughly one-third of the total crude oil. The oxygen does not play a direct role in the cracking process, but is used to regenerate the catalyst, which rapidly gets contaminated with byproduct coke. The oxygen is mixed with air to give a mixture of around 28% O2,[2] and this is used to burn the coke off the catalyst: the heat released from the combustion of the coke compensates for the heat lost in the endothermic cracking reactions, hence the need for a strict control of the rate of combustion.

Oxygen is also used in other bulk chemical processes, such as the manufactire of ethylene oxide, propylene oxide, ethylene dichloride, vinyl chloride, vinyl acetate, titanium dioxide and ferric sulfate. It can also be used to enrich the air stream in processes that normally use air, such as the production of acrylonitrile and terephthalic acid, and is used in the production of synthesis gas (CO/H2).[2][11]

Allotropes

Main articles: Dioxygen and Ozone

Oxygen forms two allotropes: dioxygen (O2) and ozone (O3). Dioxygen is by far the commonest form. It is a colourless and odourless gas at ambient temperature and pressure, and condenses to a pale blue liquid at −183.0 °C. Unusually for a diatomic molecule, it is paramagnetic, with two unpaired electrons occupying π* orbitals.

Ozone (trioxygen) is an unstable, toxic gas with a characteristic odour. It is formed from dioxygen by the action of an electric discharge (e.g., during lighteneing storms) or ultraviolet light (e.g., in photochemical smog). The O3 molecule is bent, with an angle of 116.8(5)° at the central oxygen and bond lengths of 127.8(3) pm.

Chemical properties

Water and hydrogen peroxide

Main articles: Water and Hydrogen peroxide

Oxides

Main article: Oxide

Oxygen forms binary compounds with all the elements except the noble gases (although xenon forms two well-defined oxides). Oxides cover the full range of bonding types from almost completely ionic (e.g., caesium oxide) to almost completely covalent (e.g., chlorine dioxide), and show a huge variety of structural types.

Oxygen fluorides

Main article: Oxygen fluorides

As fluorine is more electronegative than oxygen, their binary compounds are properly referred to as oxygen fluorides rather than fluorine oxides. The most stable is oxygen difluoride, OF2, first prepared in 1929, a colourless, toxic gas (bp = −145.3 °C) that is a powerful oxidizing and fluorinating agent. Dioxygen difluoride, O2F2, and other unstable species are also known.[11]

Dioxygen as a ligand

Main article: Dioxygen complexes

Organic compunds of oxygen

Biological role

Safety

Notes and references

Notes

References

  1. ↑ 1.0 1.1 Oxygen. In NIST Chemistry WebBook; National Institute for Standards and Technology, <http://webbook.nist.gov/cgi/inchi/InChI%3D1S/O2/c1-2>. (accessed 15 March 2010).
  2. ↑ 2.0 2.1 2.2 Oxygen. In Gas Encyclopedia; Air Liquide, <http://encyclopedia.airliquide.com/encyclopedia.asp?GasID=48>. (accessed 3 April 2010).
  3. ↑ Allred, A. L. Electronegativity values from thermochemical data. J. Inorg. Nucl. Chem. 1961, 17 (3–4), 215–21. DOI: 10.1016/0022-1902(61)80142-5.
  4. ↑ 4.0 4.1 Oxygen, atomic. In NIST Chemistry WebBook; National Institute for Standards and Technology, <http://webbook.nist.gov/cgi/inchi/InChI%3D1S/O>. (accessed 15 March 2010).
  5. ↑ Mohr, Peter J.; Taylor, Barry N. CODATA recommended values of the fundamental physical constants: 2002. Rev. Mod. Phys. 2005, 77 (1), 1–107. DOI: 10.1103/RevModPhys.77.1.
  6. ↑ Cordero, Beatriz; Gómez, Verónica; Platero-Prats, Ana E.; Revés, Marc; Echeverría, Jorge; Cremades, Eduard; Barragán, Flavia; Alvarez, Santiago Covalent radii revisited. Dalton Trans. 2008 (5), 2832–38. DOI: 10.1039/b801115j.
  7. ↑ Shannon, R. D. Revised effective ionic radii and systematic studies of interatomic distances in halids and chalcogenides. Acta Crystallogr. A 1976, 32 (5), 751–67. DOI: 10.1107/S0567739476001551.
  8. ↑ Bondi, A. van der Waals Volumes and Radii. J. Phys. Chem. 1964, 68 (3), 441–51. DOI: 10.1021/j100785a001.
  9. ↑ Cox, J. D.; Wagman, D. D.; Medvedev, V. A. CODATA Key Values for Thermodynamics; Hemisphere: New York, 1989. ISBN 0891167587, <http://www.codata.org/resources/databases/key1.html>.
  10. ↑ Mayow, John De sal-nitro et spiritu nitro-aereo. In Tractatus quinque medico-physici; Sheldonian: Oxford, 1674, summarized in Phil. Trans. R. Soc. London 1674, 9, 101–13, <http://gallica.bnf.fr/ark:/12148/bpt6k558143.image.f110.langEN>.
  11. ↑ 11.0 11.1 11.2 11.3 Greenwood, Norman N.; Earnshaw, A. Chemistry of the Elements; Pergamon: Oxford, 1984; pp 698–756. ISBN 0-08-022057-6.
  12. ↑ Scheele, Carl Wilhelm Chemische Abhandlung von der Luft und dem Feuer; Siverderus: Uppsala and Leipzig, 1777, <http://books.google.co.uk/books?id=beoTAAAAQAAJ>; Translated extracts.
  13. ↑ Priestley, Joseph Experiments and Observations on Different Kinds of Air. Volume II, 2nd ed.; J. Johnson: London, 1775, <http://books.google.co.uk/books?id=gB0UAAAAQAAJ>.
  14. ↑ Lavoisier, Antoine Mémoire sur la combustion en général. Mem. Acad. Sci. Paris 1777, 592, <http://www.lavoisier.cnrs.fr/ice/modules/ice2pdf/pdf/extraitPDF06-04-2010_9-53-25.pdf>.
  15. ↑ Lavoisier, Antoine Considérations générales sur la nature des acides et sur les principes dont ils sont composés. Mem. Acad. Sci. Paris 1778, <http://books.google.co.uk/books?id=qbc-AAAAcAAJ&hl=en&pg=PA248#v=onepage>.
  16. ↑ U.S. Standard Atmosphere, 1976; National Oceanic and Atmospheric Administration: Washington, D.C., 1976; p 3, <http://ntrs.nasa.gov/archive/nasa/casi.ntrs.nasa.gov/19770009539_1977009539.pdf>.

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