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Oxygen,  8O
General properties
Name, symbol Oxygen, O
Allotropes Dioxygen, ozone, tetraoxygen
Appearance Colorless (gas)
Blueish liquid (liquid)
Oxygen in the periodic table


Atomic number 8
Standard atomic weight (Ar) 15.999
Group, block (chalcogens); p-block
Period period 2
Electron configuration [He] 2s2 2p4
per shell
2, 6
Physical properties
Phase Gas
Melting point 54.36 K ​(−218.79 °C, ​−361.82 °F)
Boiling point 90.188 K ​(−182.962 °C, ​−297.332 °F)
Density at  (0 °C and 101.325 kPa) 1.429 g/L
when liquid, at  1.141 g/cm3
Triple point 54.361 K, ​0.1463 kPa
Critical point 154.581 K, 5.043 MPa
Heat of fusion 0.444 kJ/mol
Heat of 6.82 kJ/mol
Molar heat capacity 29.378 J/(mol·K)
Atomic properties
Oxidation states 2, 1, −1, −2
Electronegativity Pauling scale: 3.44
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
Crystal structure ​Cubic
Speed of sound 330 m/s (gas, at 27 °C)
Thermal conductivity 26.58×10−3 W/(m·K)
Magnetic ordering Paramagnetic
CAS Registry Number 7782-44-7
Discovery Carl Wilhelm Scheele (1771)
Named by Antoine Lavoisier (1777)
· references

Oxygen (symbol: O) is the 8th element on the periodic table and is the second strongest oxidizer, second to fluorine. It has the atomic weight of ~16 (15.9949), but as a gas it is diatomic with a molar mass of ~32.



Oxygen has 6 valence electrons and typically exists with two lone pairs of electrons.

It is highly electronegative, with a 3.44 on the Pauling scale. It reacts directly with almost every element, a major exception being most of the noble metals and gases (except for xenon).

It reacts with many metals to form oxides. These reactions can be slow and gradual (as is the rusting of iron) or extremely fast, as in the combustion of cesium.


Although in gas form it is indistinguishable from other common gases, in liquid form it is pale blue (and highly reactive). Oxygen is paramagnetic.

Oxygen has several allotrope forms, with dioxygen and ozone (trioxygen) being the most important.

Under normal conditions dioxygen exists in a triplet state denoted by the term symbol 3Σ-g or simplified chemically 3O2. However, it is possible to create dioxygen in the singlet state usually denoted by 1Δg (or simplified 1O2) photochemically or from, for example, the reaction of hypochlorite and hydrogen peroxide. At high concentrations of singlet oxygen, a red chemiluminescence is observed when the oxygen returns to the ground state:

NaOCl + H2O2 → NaCl + H2O + 1O2
1O23O2 + hv

To be precise, two transitions cause the red glow:[1]

2 O2 (1Δg) → 2 O2 (3Σ-g) + hv (λ = 633.4 nm)
O2 (3Σ+g) → O2 (3Σ-g) + hv (λ = 703.2 nm)


Gaseous oxygen makes up 21% of the atmosphere. For higher concentrations, compressed oxygen can be procured from the companies that sell welding products as well as scuba diving stores. Cryogenic liquid oxygen is harder to get because unlike liquid nitrogen it is a fire and explosive hazard, when it (accidentally) enters in contact with organic materials.


Oxygen is extracted from air by fractional distillation on an industrial scale.

In the lab, it can be isolated by:

  • electrolysis of water with containing ions
  • heating potassium chlorate with manganese dioxide,
  • decomposing hydrogen peroxide with manganese dioxide
  • Decomposition of hypochlorite solution (laundry bleach) using a small amount of cobalt chloride as a catalyst
  • heating potassium permanganate.

A way to make liquid oxygen involves liquifying normal air with liquid nitrogen and collecting the oxygen-rich layer on top.



Atmospheric oxygen is not a hazard to health, however at high concentration it becomes dangerous to the lungs and can cause blindness. At concentrations over 50%, it will greatly amplify any exothermic reaction. Liquid oxygen is a fire and explosive hazard when in contact with organic materials and a fire source. It can also cause frostbites if it touches the skin.


Oxygen tanks and dewars should be kept in dark and cool places, away from any combustible materials. Periodically check the valves for any leaks.


Oxygen can be safely released in open air, but never in closed places.


See also


  1. P. Lechtken, Chemie in unserer Zeit 1974, 8, 11-16, doi:10.1002/ciuz.19740080103

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