Oxygen – facts and properties
- Symbol
- O
- Atomic number
- 8
- Atomic mass
- 15.999 u
- Period
- 2
- Group
- 16
- Category
- diatomic nonmetal
- Electron configuration
- [He] 2s² 2p⁴
- Electronegativity
- 3.44
- Atomic radius
- 48 pm
- Valence electrons
- 6
- State at standard conditions
- gas
- Melting point
- -218.79 °C
- Boiling point
- -183 °C
- Density
- 1.43 g/L
- Discovered
- 1774
Oxygen is the element without which neither fire nor animal life would exist. It has atomic number 8 and is the most abundant element in the Earth's crust, in the oceans and in the human body. In the air it is a free gas; in water, rock and living tissue it is chemically bound. Oxygen is a powerful oxidising agent and forms compounds with almost every other element.
Position in the periodic table
Occurrence and production
Oxygen is everywhere on Earth. Air is about 21 per cent oxygen by volume. Water contains even more by mass, since the oxygen atom accounts for most of the weight of each water molecule. The crust is nearly half oxygen by mass, bound in silicates, oxides and carbonates, the minerals that make up sand, clay, granite and limestone. Roughly two thirds of the mass of the human body is oxygen, mostly as water.
The free oxygen in the atmosphere is a product of life. Cyanobacteria and, later, plants release it during photosynthesis, when they use light energy to split water and build sugars from carbon dioxide. Before photosynthesis evolved, the atmosphere contained virtually no oxygen. Respiration, combustion, weathering and rusting consume oxygen again, and photosynthesis replaces it. Industrially, oxygen is produced mainly by separating air with the Linde process: air is cleaned, liquefied and distilled, and oxygen, the higher-boiling component, remains as a liquid. Smaller quantities come from oxygen concentrators, which use molecular sieves to remove nitrogen from air and are common in home medical care, and from the electrolysis of water. Carl Wilhelm Scheele and Joseph Priestley discovered oxygen independently; Priestley isolated the gas in 1774 by heating mercury oxide, and Antoine Lavoisier soon afterwards explained its role in combustion and gave it its name.
Properties and reactions
Oxygen has the electron configuration [He] 2s² 2p⁴ with six valence electrons, two short of a full shell. It belongs to period 2 and group 16, the chalcogens. Its electronegativity of 3.44 is second only to that of fluorine, so in almost all compounds oxygen has the oxidation state minus two; the main exceptions are peroxides such as hydrogen peroxide, where it is minus one, a few superoxides, and compounds with fluorine, where oxygen is positive.
The element forms diatomic molecules. The gas is colourless and odourless, while liquid oxygen is pale blue and is attracted by a magnet because the molecule has two unpaired electrons. The gas has a density of 1.43 g/L; it boils at -183 °C and freezes at -218.79 °C. A second form, ozone, consists of three atoms per molecule. It has a sharp smell, irritates the lungs at ground level and shields the Earth from ultraviolet light in the stratosphere. Oxygen is a strong oxidising agent. It combines with most elements to form oxides, often with the release of heat and light: magnesium burns with a dazzling white flame, sulfur with a blue one, and steel wool throws off sparks. Every ordinary fire is a reaction with oxygen, and the glowing splint test identifies the gas because a smouldering splint bursts into flame in pure oxygen. Rusting and cellular respiration are oxidations too, only slow ones. Oxygen dissolves slightly in water, which is what fish and other aquatic animals rely on.
Uses in everyday life and technology
Steelmaking consumes the largest share of industrial oxygen. In the basic oxygen converter, a jet of pure oxygen is blown onto molten pig iron to burn off excess carbon and other impurities, turning it into steel. Oxygen also enriches the flames in copper, lead and zinc smelting, in glass furnaces and in cement kilns, making them hotter and cleaner. In oxy-acetylene welding and cutting, the flame reaches temperatures that melt steel.
The chemical industry uses oxygen to make ethylene oxide, nitric acid and many intermediates. Sewage works bubble oxygen through wastewater so that bacteria break down organic matter faster, and fish farms oxygenate their tanks. In medicine, patients with lung disease receive oxygen through masks or from concentrators, and hospitals pipe it from central tanks to every ward. High-altitude climbers, pilots and astronauts carry oxygen with them, and divers blend it into their breathing gas. Rockets burn liquid hydrogen or kerosene with liquid oxygen, which they carry in enormous quantities. In everyday life you encounter oxygen as the active part of the air, as ozone in summer smog, and in hydrogen peroxide, which bleaches hair and disinfects wounds.
Molecules containing this element
- Oxygen O2
- Water H2O
- Hydrogen peroxide H2O2
Quick question
What is the atomic number of oxygen (O)?
FAQ
Can you breathe too much oxygen?
Yes. Pure oxygen breathed for long periods or under raised pressure damages the lungs and the nervous system. Divers must adjust the oxygen content of their breathing gas to their depth, and medical oxygen is carefully dosed. Breathing oxygen briefly at normal pressure is harmless.
Why doesn't oxygen itself burn?
Burning means reacting with oxygen. Oxygen is the partner in that reaction, not the fuel, so it cannot burn on its own. It does make other things burn far more fiercely: a glowing splint flares up in pure oxygen, and oil or grease can ignite spontaneously on contact with it. That is why oxygen cylinders and fittings must be kept free of oil.
Where does the oxygen in the air come from?
Almost all of it comes from photosynthesis. Algae and cyanobacteria in the oceans produce roughly half, and land plants the other half. Oxygen began to accumulate in the atmosphere more than two billion years ago. It was poisonous to many organisms of the time, but it made possible the energy-rich metabolism on which animals, including humans, depend.