Oxygen
Oxygen is a chemical element with symbol O and atomic number 8. It is a member of the chalcogen group on the periodic table and is a highly reactive nonmetal and oxidizing agent that readily forms compounds (notably oxides) with most elements. By mass, oxygen is the third-most abundant element in the universe, after hydrogen and helium.
Identity
- Atomic mass
- [15.99903, 15.99977] u
- Category
- Reactive non-metals
- Group
- 16
- Period
- 2
- Block
- p
- Oxidation States
- −2, −1, +1, +2
- Phase (25 °C)
- Gas
Electronic Structure
- Electron configuration
- 1s2 2s2 2p4
- Short configuration
- [He] 2s2 2p4
- Electrons per shell
- 2, 6
- Electronegativity (Pauling)
- 3.44
- Electron affinity
- 141 kJ/mol
- 1st ionization energy
- 1.31e+3 kJ/mol
- Ionization energies
- 1314 · 3388 · 5301 · 7469 · 10990… kJ/mol
Physical properties
- Melting point
- -218.8 °C (54 K)
- Boiling point
- -183.0 °C (90 K)
- Density
- 0.001429 g/cm³
- Molar heat capacity
- —
Chemical behavior of oxygen
Oxygen has six valence electrons ([He] 2s² 2p⁴) and is the second most electronegative element (3.44), so it almost always acts as -2 by capturing two electrons in ionic or polar covalent bonds. Only halogens and noble gases compete for electrons (states +1 and +2 in OF₂ or with xenon). Its O₂ molecule, with a double bond, is a powerful oxidant but with high kinetic barrier; the -1 oxidation state appears in peroxides such as H₂O₂.
Reactivity
2 Mg + O₂ → 2 MgO
With metals it forms oxides releasing great energy: magnesium burns with blinding white light. Almost all metals oxidise, from iron to passive aluminium.
CH₄ + 2 O₂ → CO₂ + 2 H₂O
It oxidises organic matter in combustion and respiration: it is the reaction that fuels boilers, engines and cells. Without oxygen there is no fire or aerobic metabolism.
2 H₂ + O₂ → 2 H₂O
With hydrogen it forms water with huge energy release (286 kJ/mol). The mixture is explosive (detonating gas) and propels cryogenic rockets such as Ariane.
2 F₂ + O₂ → 2 OF₂
Only fluorine, more electronegative, oxidises it: oxygen acts here as +2 in oxygen difluoride, an even more violent oxidant. The exception that confirms its oxidising power.
Notable compounds
| Compound | Formula | Use or importance |
|---|---|---|
| Water | H₂O | Life support; solvent, refrigerant and essential resource |
| Carbon dioxide | CO₂ | Photosynthesis; fire extinguishers, beverages and climate regulation |
| Calcium oxide | CaO | Quicklime; cement, steelmaking and soil correction |
| Hydrogen peroxide | H₂O₂ | Antiseptic, bleaching agent and rocket propellant |
| Ozone | O₃ | Stratospheric ultraviolet shield; oxidant and disinfectant |
Electron shell distribution
View in the atom builderIsotopes
Natural oxygen combines oxygen-16 (¹⁶O), with eight neutrons and 99.76%, oxygen-17 (¹⁷O), with nine neutrons and 0.04%, and oxygen-18 (¹⁸O), with ten neutrons and 0.20%. ¹⁸O reconstructs past temperatures in polar ice and ¹⁷O traces atmospheric processes; both are used in medicine and metabolism.
Occurrence and production
It is the most abundant element in the crust (46% by mass) in oxides and silicates, and forms 21% of air as O₂ and 86% of water by mass. It was released by photosynthesis about 2.4 billion years ago (Great Oxidation). Industrially it is obtained by distillation of liquid air (millions of tonnes) and in the laboratory by electrolysis or decomposition of chlorate.
History and origin of the name
Priestley obtained it in 1774 by heating mercury oxide and observed that it enhanced the flame; Scheele had isolated it earlier without publishing. Lavoisier baptised it oxygen («generator of acids») and overturned the phlogiston theory, demonstrating the role of oxygen in combustion and respiration. Ozone was discovered by Schönbein in 1840 by its smell after storms.
Applications and relevance of oxygen
Uses of the element
As O₂, it feeds the steel industry (converters that burn impurities), oxyacetylene welding, medical respiratory therapy and aquaculture. Liquid oxygen propels rockets and ozone disinfects waters without chlorine.
Uses of its compounds
Water is a resource and universal solvent; lime (CaO) makes cement and steel; peroxide disinfects and bleaches; metal oxides provide ceramics, pigments and glass; CO₂ feeds greenhouses and extinguishes fires.
Health, safety and environment
Biological role
It is essential: aerobic respiration extracts about 18 times more energy from glucose than without oxygen. But in excess it is toxic: at high pressures it damages lungs and the nervous system (Paul Bert effect in diving).
Safety
It violently accelerates any fire: oxygen leaks turn a spark into a fire and fats burn spontaneously in pure O₂. Liquid oxygen (−183 °C) causes cold burns and ozone irritates the respiratory tract.
Environment
Oxygen sustains the ozone layer that blocks ultraviolet light, but CFCs thinned it (Antarctic ozone hole, recovery after Montreal 1987). Excess nutrients deplete O₂ in water (eutrophication) and burning fossil fuels consumes oxygen, though not at risk of depletion.
Frequently asked questions about Oxygen
Why is oxygen so reactive if we breathe 21% without burning?
Because O₂ molecule, although a powerful oxidant, has a high kinetic barrier: it needs a spark, an enzyme or high temperature to start reacting. At 20°C and without a catalyst, neither paper nor our skin burns in air. Cells use enzymes that release their energy step by step in respiration; a flame releases it all at once. It is a sleeping oxidant that only wakes with sufficient activation.
Is pure oxygen better for breathing?
No: breathing pure O₂ at normal pressure irritates lungs in hours and at high pressure causes convulsions (Paul Bert effect known in diving). Air with its 21% is an evolutionary balance; more oxygen generates free radicals that damage tissues. In medicine it is dosed with precision and in technical diving partial pressures are limited. More is not better.
Why does the blue colour of the sky have to do with oxygen?
N₂ and O₂ molecules in air scatter sunlight, and blue, with short wavelength, scatters much more than red (Rayleigh scattering). That is why the daytime sky is blue and the sunset turns red, when light passes through more air thickness. Without an atmosphere there is no colour: the sky seen from the Moon is black even in full daylight.
Why does iron rust and gold not?
Because rusting means donating electrons to oxygen, and each metal has a different 'generosity': iron, magnesium or sodium donate them easily and corrode; gold and platinum retain them and remain intact for millennia. The reactivity series orders metals by this tendency. Aluminium is a curious exception: it oxidises instantly, but its transparent oxide seals and protects it.
What was the Great Oxidation?
The planetary poisoning that created our air: about 2.4 billion years ago, photosynthetic cyanobacteria filled the atmosphere with oxygen, toxic to almost all anaerobic life of that time. It caused a mass extinction, oxidised the iron of the oceans (current mines) and opened the era of aerobic respiration, 18 times more efficient. We owe the air to a microbial catastrophe.
What is the difference between oxygen, ozone and peroxide?
They are the same chemical family with different numbers of oxygen atoms: O₂ that we breathe (double bond, stable molecule), O₃ or ozone (three atoms, aggressive oxidant that forms the stratospheric shield) and H₂O₂ or hydrogen peroxide (with a weak O–O bond that releases active oxygen when it breaks). From vital air to pharmacy disinfectant only the molecular architecture of oxygen changes.
Can the planet's oxygen run out?
No on a human scale: the atmospheric reserve is equivalent to thousands of years of respiration and photosynthesis, and burning all fossil fuels would consume less than 2%. The real danger is not the global amount but the local: eutrophication that leaves seas without oxygen, unventilated interiors or loss of forests. Taking care of oxygen is taking care of its cycles, not counting molecules.
Main sources consulted: PubChem, Royal Society of Chemistry, NIST, IUPAC. Last revision: November 2026.