Average mass of an atom in unified atomic mass units (1 u is 1/12 of the mass of a carbon-12 atom). Accounts for the natural mix of the element’s isotopes.
Number of the vertical column, 1 to 18 in IUPAC numbering. Elements in a group share the same outer-shell layout, so their properties are alike.
Number of the horizontal row. Usually equal to the number of electron shells in the atom; palladium is the exception, with four shells in period 5.
The subshell (s, p, d or f) that receives the last electrons in the usual filling order; it divides the table into four blocks. Real atoms sometimes differ: lanthanum starts the f-block but has no f electrons.
[He] 2s² 2p⁴
How electrons are arranged in subshells. The bracket stands for the previous noble gas; the superscript is the number of electrons in a subshell.
2, 6
Number of electrons in each shell, from the nucleus outwards.
-218.8 °C (54.36 K)
Temperature at which the solid turns liquid at standard pressure. Helium, carbon and arsenic do not melt at normal pressure: helium needs about 25 atm, carbon and arsenic turn straight into vapour.
-182.9 °C (90.2 K)
Temperature at which the liquid boils at standard pressure. For carbon and arsenic, the temperature at which the solid turns straight into vapour (sublimation).
0.001429 g/cm³
Mass of one cubic centimetre of the substance. For gases, at 0 °C and standard pressure.
3.44 Pauling
How strongly an atom pulls electrons towards itself in a chemical bond. On the Pauling scale it runs from 0.79 for caesium to 3.98 for fluorine.
13.618 eV
Energy needed to remove the most loosely bound electron from an atom. The lower it is, the more easily the element gives up an electron, which is typical of metals.
1.461 eV
Energy released when a neutral atom gains an electron. Highest for the halogens. Noble gases, nitrogen and some metals form no stable negative ion.
66 pm
Half the length of a single covalent bond between two atoms of the element. It shows the size of an atom inside a molecule or crystal: 76 pm for carbon, 132 pm for iron.
-2
Charges the atom would have in its compounds if all its bonds were ionic; the main ones are listed.
Oxygen (O) is a chemical element with atomic number 8. It sits in group 16 and period 2 of the periodic table, in the p-block. It is a nonmetal. At room temperature it is a gas. Its standard atomic weight is 15.999.
Origin of the name
In 1777 Antoine Lavoisier named the gas oxygène, from Greek oxýs ‘sharp, acid’ and -genēs ‘producing’: he believed every acid contained oxygen. He was wrong, since hydrogen chloride, for one, has none, but by then the name had stuck.
Where is oxygen found?
Oxygen is the third most abundant element in the universe, after hydrogen and helium, and the most common in the Earth’s crust, bound in silicates and other minerals. It makes up 21% of the air and nearly nine tenths of the mass of water. Almost all the oxygen in the air was released by photosynthesis in plants and microbes. Industry obtains it from liquefied air.
What is oxygen used for?
The largest use of oxygen is in steelmaking, where a jet of oxygen burns impurities out of molten iron. A flame of acetylene in oxygen welds and cuts metal. Hospitals give oxygen to patients who struggle to breathe; liquid oxygen, burned with liquid hydrogen, is a rocket propellant.
Interesting facts
Ozone, O₃, a form of oxygen high in the atmosphere, blocks most of the Sun’s ultraviolet light.
Carl Wilhelm Scheele made oxygen in 1771, but his book came out only in 1777, so Joseph Priestley, who published in 1774, was long credited with the discovery.
Liquid oxygen is attracted by a magnet and can even hang between its poles.
Pure oxygen is not good for us over long periods: breathing it for many hours can damage the lungs.