The periodic table/Oxygen
Discovery[edit | edit source]
Oxygen was discovered in 1774 by Joseph Priestley in Wiltshire, England, and independently by Carl Wilhelm Scheele in Uppsala, Sweden. The name is derived from the Greek, 'oxy genes' meaning acid forming.
The credit for discovering oxygen is shared by three chemists: an Englishman, a Swede, and a Frenchman. Joseph Priestley was the first to publish an account of oxygen, having made it in 1774 by focusing sunlight on to mercuric oxide (HgO), and collecting the gas which came off. Unknown to Priestly, Carl Wilhelm Scheele had produced oxygen in June 1771. He had written an account of his discovery but it was not published until 1777. Antoine Lavoisier also claimed to have discovered oxygen, and he proposed that the new gas be called oxy-gène, meaning acid-forming, because he thought it was the basis of all acids.
Quick Facts[edit | edit source]
Atomic Number: 8
Electron Configuration: [He] 2s2 2p4
CAS Number: 7782-44-7
Appearance: colourless odourless gas
Discovery in: 1774
Key Isotopes: 16O
Allotropes: O2, O3
Density: 1.429 g/L
Crystal Structure: cubic
Melting Point: -218.79 °C
Boiling Point: -182.962 °C
Uses[edit | edit source]
Industrially, oxygen is produced on a large scale from liquid air by liquefaction and fractional distillation. In the laboratory it can be prepared by the electrolysis of water or by adding manganese(IV) oxide as a catalyst to aqueous hydrogen peroxide. Oxygen is very reactive and capable of combining with most other elements. It is a component of thousands of organic compounds, and is essential for the aerobic respiration of all plants and animals and for almost all combustion. The greatest commercial use of gaseous oxygen is in the steel industry. Large quantities are also used in the manufacture of epoxyethane, nitric acid, hydrogen peroxide and chloroethene, the precursor to PVC and for oxy-acetylene welding and cutting of metals. A growing use is in the treatment of sewage and of effluent from industry.
Oxygen gas is fairly soluble in water, which makes aerobic life in rivers, lakes and oceans possible. Oxygen first appeared as a constituent of Earth's atmosphere around 2 billion years ago, when sufficient oxygen from the photosynthesis of blue-green algae had accumulated. In this process energy from the Sun splits water into oxygen, which passes into the atmosphere, and hydrogen, which joins with carbon dioxide to produce biomass. When living things need energy they take in oxygen so that it can react in their cells with the biomass absorbed by digestion of food; they then return that oxygen to the atmosphere in the form of carbon dioxide. We, for example, breathe in oxygen so that it can react with the fuel (food) in our bodies allowing the transfer of energy for our cells that keeps us alive, active and warm. We breathe out the carbon dioxide that forms.
Oxygen, as a gaseous element, forms 21% of the atmosphere by volume. The element and its compounds make up 49.2%, by mass of the Earth’s crust, about two-thirds of the human body and nine-tenths of water.
Atomic Data[edit | edit source]
Atomic radius: 1.520 Å
Covalent radius: 0.64 Å
Electron affinity: 140.926 kJ mol-1
First: 1313.942 kJ mol-1
Second: 3388.668 kJ mol-1
Third: 5300.466 kJ mol-1
Fourth: 7469.264 kJ mol-1
Fifth: 10989.574 kJ mol-1
Sixth: 13326.515 kJ mol-1
Seventh: 71330.586 kJ mol-1
Eighth: 84078.228 kJ mol-1
Supply Risk[edit | edit source]
Oxidation States and Isotopes[edit | edit source]
Common oxidation states: -1, -2
|Isotope||Atomic mass||Abundance (%)||Half life||Mode of decay|
Pressure and Temperature Data[edit | edit source]
Molar heat capacity: 29.378 J mol-1 K-1