At 46 percent of the mass, oxygen is the most plentiful element in Earth’s crust. The
proportion of oxygen by volume in the atmosphere is 21 percent and by weight in
seawater is 89 percent. In rocks, it is combined with metals and nonmetals in the
form of oxides that are acidic (such as those of sulfur, carbon, aluminum, and
phosphorus) or basic (such as those of calcium, magnesium, and iron) and as saltlike
compounds that may be regarded as formed from the acidic and basic oxides, as
sulfates, carbonates, silicates, aluminates, and phosphates. Plentiful as they are,
these solid compounds are not useful as sources of oxygen, because separation of
the element from its tight combinations with the metal atoms is too expensive.
During respiration, animals and some bacteria take oxygen from the atmosphere
and return to it carbon dioxide, whereas by photosynthesis, green plants assimilate
carbon dioxide in the presence of sunlight and evolve free oxygen. Almost all the
free oxygen in the atmosphere is due to photosynthesis. About 3 parts of oxygen by
volume dissolve in 100 parts of fresh water at 20 °C (68 °F), slightly less in
seawater. Dissolved oxygen is essential for the respiration of fish and other marine
life.
Natural oxygen is a mixture of three stable isotopes: oxygen-16 (99.759 percent),
oxygen-17 (0.037 percent), and oxygen-18 (0.204 percent). Several artificially
prepared radioactive isotopes are known. The longest-lived, oxygen-15 (124-second
half-life), has been used to study respiration in mammals.
Allotropy
Oxygen has two allotropic forms, diatomic (O2) and triatomic (O3, ozone). The
properties of the diatomic form suggest that six electrons bond the atoms and two
electrons remain unpaired, accounting for the paramagnetism of oxygen. The three
atoms in the ozone molecule do not lie along a straight line.
Ozone may be produced from oxygen according to the equation:
The process, as written, is endothermic (energy must be provided to make it
proceed); conversion of ozone back into diatomic oxygen is promoted by the
presence of transition metals or their oxides. Pure oxygen is partly transformed into
ozone by a silent electrical discharge; the reaction is also brought about by
absorption of ultraviolet light of wavelengths around 250 nanometres (nm, the
nanometre, equal to 10−9 metre); occurrence of this process in the upper atmosphere
removes radiation that would be harmful to life on the surface of the Earth. The
pungent odour of ozone is noticeable in confined areas in which there is sparking of
electrical equipment, as in generator rooms. Ozone is light blue; its density is 1.658
times that of air, and it has a boiling point of −112 °C (−170 °F) at atmospheric
pressure.
Ozone is a powerful oxidizing agent, capable of converting sulfur dioxide to sulfur
trioxide, sulfides to sulfates, iodides to iodine (providing an analytical method for its
estimation), and many organic compounds to oxygenated derivatives such as
aldehydes and acids. The conversion by ozone of hydrocarbons from automotive
exhaust gases to these acids and aldehydes contributes to the irritating nature of
smog. Commercially, ozone has been used as a chemical reagent, as a disinfectant,
in sewage treatment, water purification, and bleaching textiles.
Preparative methods
Production methods chosen for oxygen depend upon the quantity of the element
desired. Laboratory procedures include the following:
1.
Thermal decomposition of certain salts, such as potassium
chlorate or potassium nitrate:
The decomposition of potassium chlorate is catalyzed by oxides of transition metals;
manganese dioxide (pyrolusite, MnO2) is frequently used. The temperature
necessary to effect the evolution of oxygen is reduced from 400 °C to 250 °C by the
catalyst.
2.
Thermal decomposition of oxides of heavy metals:
Scheele and Priestley used mercury(II) oxide in their preparations of oxygen.
3.
Thermal decomposition of metal peroxides or of hydrogen peroxide:
An early commercial procedure for isolating oxygen from the atmosphere or
for manufacture of hydrogen peroxide depended on the formation of barium
peroxide from the oxide as shown in the equations.
4.
Electrolysis of water containing small proportions of salts or acids to allow
conduction of the electric current:
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