oxygen - Applications, Compounds, Isotopes

O, element 8, boiling point ?183°C. By far the commonest element in the Earth's crust, of which it makes up nearly 50% in various combined forms. It also constitutes 21% of the atmosphere as diatomic molecules (O₂). Although this is very reactive, it is constantly replenished by photosynthesis. All higher forms of life depend on oxygen. The element boils at 13°C higher than nitrogen, and is isolated by the fractional distillation of air. Liquid oxygen is pale blue in colour, and is used as an oxidant. Oxygen occurs widely in organic and inorganic compounds, mainly showing oxidation state ?2.

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8 nitrogen ← oxygen → fluorine - ↑ O ↓ S Periodic Table - Extended Periodic Table
General
Name, Symbol, Number	oxygen, O, 8
Chemical series	Nonmetals, chalcogens
Group, Period, Block	16, 2, p
Appearance	colorless (gas) very pale blue (liquid)
Atomic mass	15.9994(3) g/mol
Electron configuration	1s 2p4
Electrons per shell	2, 6
Physical properties
Phase	gas
Density	(0 °C, 101.325 kPa) 1.429 g/L
Melting point	54.36 K (-218.79 °C, -361.82 °F)
Boiling point	90.20 K (-182.95 °C, -297.31 °F)
Critical point	154.59 K, 5.043 MPa
Heat of fusion	(O2) 0.444 kJ·mol−1
Heat of vaporization	(O2) 6.82 kJ·mol−1
Heat capacity	(25 °C) (O2) 29.378 J·mol
Vapor pressure P/Pa 1 10 100 1 k 10 k 100 k at T/K   61 73 90
Atomic properties
Crystal structure	cubic
Oxidation states	−2, −1 (neutral oxide)
Electronegativity	3.44 (Pauling scale)
Ionization energies (more)	1st: 1313.9 kJ·mol−1
	2nd: 3388.3 kJ·mol−1
	3rd: 5300.5 kJ·mol−1
Atomic radius	60 pm
Atomic radius (calc.)	48 pm
Covalent radius	73 pm
Van der Waals radius	152 pm
Miscellaneous
Magnetic ordering	paramagnetic
Thermal conductivity	(300 K) 26.58 mW·m
Speed of sound	(gas, 27 °C) 330 m/s
CAS registry number	7782-44-7
Selected isotopes
Main article: Isotopes of oxygen iso NA half-life DM DE (MeV) DP 16O 99.76% O is stable with 8 neutrons 17O 0.038% O is stable with 9 neutrons 18O 0.21% O is stable with 10 neutrons
References

Oxygen (IPA: /ˈɔksidʒən/) is a chemical element with the chemical symbol O and atomic number 8.

At standard temperature and pressure, oxygen exists as a diatomic molecule with the formula O2, in which the two oxygen atoms are doubly bonded to each other. In its most stable form, oxygen exists as a diradical (triplet oxygen) with two unpaired electrons in molecular orbitals of the O2 molecule.

Singlet oxygen, a name given to several higher energy species of molecular oxygen in which all the electron spins are paired, is much more reactive towards common organic molecules. Liquid oxygen is attracted to a magnet to a sufficient extent that a bridge of liquid oxygen may be supported against its own weight between the poles of a powerful magnet, in laboratory demonstrations.

Applications

Oxygen is essential to respiration, so oxygen supplementation has found use in medicine (as oxygen therapy). the reason is that increasing the proportion of oxygen in the breathing gas at low pressure acts to increase the inspired oxygen partial pressure nearer to that found at sea-level. A notable application of oxygen as a very low-pressure breathing gas, is in modern spacesuits, where use of nearly pure oxygen at a total pressure of about 1/3rd normal, results in normal blood partial pressures of oxygen.

Oxygen, as a supposed mild euphoric, has a history of recreational use (see oxygen bar), however the reality of this effect is doubtful.

See also Silicate minerals, Oxide minerals.

Compounds

Due to its electronegativity, oxygen forms chemical bonds with almost all other elements hence the origin of the original definition of oxidation. A double oxygen molecule (O2)2 is known and is found as a minor component of liquid oxygen.

See also Oxygen compounds.

Isotopes

Oxygen has seventeen known isotopes with atomic masses ranging from 12.03 u to 28.06 u. Since oxygen partial pressure is the fraction of oxygen times the total pressure, elevated partial pressures can occur either from high oxygen fraction in breathing gas, or from high breathing gas pressure, or a combination of both. Oxygen toxicity usually begins to occur at partial pressures more than 0.5 atmospheres, or 2.5 times the normal sea-level oxygen partial pressure of about 0.2 atmospheres or bars. However in medical applications (such as in ventilation gas mixtures in hospital applications) mixtures containing more than 50% oxygen can be expected to show lung toxicity, causing slow damage to the lungs over periods of days, with the rate of damage rising rapidly from mixtures between 50% and 100% oxygen. In the case of spacesuits, oxygen partial pressure in the breathing gas is typically about 0.30 bar (1.4 times normal), and oxygen partial pressure in the astronaut's blood (due to downward adjustments due to water vapor and CO2 in the alveoli) is close to sea-level normal of 0.14 bar.

In deep scuba diving and surface supplied diving and when using equipment which can provide high partial pressures of oxygen, such as rebreathers, oxygen toxicity to the lungs can occur, just as in medical applications. This form of oxygen toxicity usually occurs after several hours exposure to oxygen partial pressures over about 1.4 atmospheres (bars) (i.e.

Toxicity and antibacterial use of other chemical oxygen forms

Certain derivatives of oxygen, such as ozone (O3), singlet oxygen, hydrogen peroxide, hydroxyl radicals and superoxide, are also highly toxic. Not only do antibodies catalyze production of peroxide from oxygen, it is now known that immune cells produce peroxide, superoxide, and singlet oxygen in the course of an immune response. Recently, singlet oxygen has been found to be a source of biologically-produced ozone: this reaction proceeds through an unusual compound dihydrogen trioxide, also known as trioxidane, (HOOOH) which is an antibody-catalyzed product of singlet oxygen and water.