oxygen

oxygen
Name	oxygen
Category	nonmetal
Group	16
Appearance	gas: colorless, liquid: pale blue
Atomic weight standard	15.999
Electron configuration	1s² 2s² 2p⁴
Electrons per shell	2, 6
Phase at STP	gas
Melting point K	54.36
Boiling point K	90.188
Density at STP g/L	1.429
Triple point K	54.361
Triple point kPa	0.1463
Critical point K	154.581
Critical point MPa	5.043
Heat of fusion	(O₂) 0.444 kJ/mol
Heat of vaporization	(O₂) 6.82 kJ/mol
Molar heat capacity	(O₂) 29.378 J/(mol·K)
Oxidation states	−2, −1, 0, +1, +2
Electronegativity	Pauling scale: 3.44
Ionization energies	1st: 1313.9 kJ/mol, 2nd: 3388.3 kJ/mol, 3rd: 5300.5 kJ/mol
Covalent radius	66±2 pm
Van der Waals radius	152 pm
CAS number	7782-44-7
Discoverer	Carl Wilhelm Scheele, Joseph Priestley
Discovered	1771–1774
Named by	Antoine Lavoisier
Etymology	Greek oxy + genes (acid-forming)

oxygen. It is a member of the chalcogen group on the periodic table and is a highly reactive nonmetal. As the most abundant element in the Earth's crust and a critical component of water and the Earth's atmosphere, it is essential for aerobic respiration in most living organisms. The element was independently discovered in the 1770s by Carl Wilhelm Scheele and Joseph Priestley, with its nature later elucidated by Antoine Lavoisier.

Properties

At standard temperature and pressure, two atoms of the element bond to form dioxygen, a colorless, odorless diatomic gas with the formula O₂. It has a melting point of 54.36 K and a boiling point of 90.188 K, condensing into a distinctive pale blue liquid at cryogenic temperatures. The element exhibits several allotropes; the common O₂ form is paramagnetic, while another, ozone (O₃), is a powerful oxidizing agent in the stratosphere. Its high electronegativity, second only to fluorine, underpins its strong reactivity, forming compounds with almost all other elements except the lighter noble gases.

Occurrence and production

It is the most abundant element in the Earth's crust, largely present in silicate minerals like quartz and feldspar, and constitutes about 46% of the crust by mass. The Earth's atmosphere contains approximately 21% by volume as O₂, primarily sustained by photosynthesis from organisms like cyanobacteria and terrestrial plants. Commercially, the vast majority is produced via fractional distillation of liquid air, a process pioneered by companies like Linde and Air Liquide. Other methods include pressure swing adsorption and the electrolysis of water, the latter being used in niche applications and supported by research at institutions like the Massachusetts Institute of Technology.

Biological role and medical use

In cellular respiration, O₂ serves as the terminal electron acceptor in the mitochondrial electron transport chain, enabling efficient ATP production in organisms from Escherichia coli to Homo sapiens. Medical applications are extensive; supplemental O₂ is a cornerstone therapy for conditions like chronic obstructive pulmonary disease, pneumonia, and during cardiac arrest. It is administered via devices such as nasal cannulas, simple face masks, and in critical care, through mechanical ventilation. Hyperbaric oxygen therapy, used for treating decompression sickness and serious infections, is a standard protocol at facilities like the Undersea and Hyperbaric Medical Society.

Chemical compounds and reactions

The element forms a vast array of compounds, most notably water (H₂O), silicon dioxide (SiO₂), and iron(III) oxide (Fe₂O₃). Its reactivity is central to combustion and corrosion processes; for instance, the rusting of iron is an electrochemical oxidation reaction. In organic chemistry, it is a key constituent of functional groups such as alcohols, ketones, and carboxylic acids. Important industrial reactions involving it include the Haber process for ammonia and the basic oxygen steelmaking process developed by companies like Voestalpine. The study of its reaction kinetics is a major focus within the American Chemical Society.

History and discovery

While the effects of air were noted by ancients like Empedocles, the isolation of the element is credited to the work of Carl Wilhelm Scheele in Uppsala around 1771 and Joseph Priestley in Wiltshire in 1774, who produced it by heating mercury(II) oxide. Priestley published his findings in the journal Philosophical Transactions of the Royal Society. However, it was Antoine Lavoisier in Paris who conducted extensive experiments, disproved the phlogiston theory, named the element "oxygen" from Greek roots, and correctly identified its role in combustion and respiration, as detailed in his work Traité Élémentaire de Chimie.

Applications

Beyond medical uses, its primary applications are in metallurgy, such as in blast furnaces and electric arc furnaces for steel production by companies like ArcelorMittal. It is crucial in chemical synthesis for producing ethylene oxide and polyesters. In aerospace, liquid O₂ is a standard oxidizer in rocket propellants, used in engines like those on the Saturn V and the Space Shuttle. Other significant uses include oxy-fuel welding and cutting, water treatment in facilities managed by the Environmental Protection Agency, and life support systems in submarines and spacecraft.

Category:Chemical elements Category:Nonmetals Category:Chalcogens