Commercial production and use
When required in tonnage quantities, oxygen is prepared by the fractional
distillation of liquid air. Of the main components of air, oxygen has the highest
boiling point and therefore is less volatile than nitrogen and argon. The process
takes advantage of the fact that when a compressed gas is allowed to expand, it
cools. Major steps in the operation include the following: (1) Air is filtered to
remove particulates; (2) moisture and carbon dioxide are removed by absorption in
alkali; (3) the air is compressed and the heat of compression removed by ordinary
cooling procedures; (4) the compressed and cooled air is passed into coils contained
in a chamber; (5) a portion of the compressed air (at about 200 atmospheres
pressure) is allowed to expand in the chamber, cooling the coils; (6) the expanded
gas is returned to the compressor with multiple subsequent expansion and
compression steps resulting finally in liquefaction of the compressed air at a
temperature of −196 °C; (7) the liquid air is allowed to warm to distill first the light
rare gases, then the nitrogen, leaving liquid oxygen. Multiple fractionations will
produce a product pure enough (99.5 percent) for most industrial purposes.
The steel industry is the largest consumer of pure oxygen in “blowing” high
carbon steel—that is, volatilizing carbon dioxide and other nonmetal impurities in a
more rapid and more easily controlled process than if air were used. The treatment
of sewage by oxygen holds promise for more efficient treatment of liquid effluents
than other chemical processes. Incineration of wastes in closed systems using pure
oxygen has become important. The so-called LOX of rocket oxidizer fuels is liquid
oxygen; the consumption of LOX depends upon the activity of space programs. Pure
oxygen is used in submarines and diving bells.
Commercial oxygen or oxygen-enriched air has replaced ordinary air in the
chemical industry for the manufacture of such oxidation-controlled chemicals as
acetylene, ethylene oxide, and methanol. Medical applications of oxygen include
use in oxygen tents, inhalators, and pediatric incubators. Oxygen-enriched gaseous
anesthetics ensure life support during general anesthesia. Oxygen is significant in a
number of industries that use kilns.
Chemical properties and reactions
The large values of the electronegativity and the electron affinity of oxygen are
typical of elements that show only nonmetallic behaviour. In all of its compounds,
oxygen assumes a negative oxidation state as is expected from the two half-filled
outer orbitals. When these orbitals are filled by electron transfer, the oxide ion O2−
is created. In peroxides (species containing the ion O22−) it is assumed that each
oxygen has a charge of −1. This property of accepting electrons by complete or
partial transfer defines an oxidizing agent. When such an agent reacts with an
electron-donating substance, its own oxidation state is lowered. The change
(lowering), from the zero to the −2 state in the case of oxygen, is called a reduction.
Oxygen may be thought of as the “original” oxidizing agent, the nomenclature used
to describe oxidation and reduction being based upon this behaviour typical of
oxygen.
As described in the section on
allotropy
, oxygen forms the diatomic species, O
2
,
under normal conditions and, as well, the triatomic species ozone, O
3
. There is some
evidence for a very unstable tetratomic species, O
4
. In the molecular diatomic form
there are two unpaired electrons that lie in antibonding orbitals. The paramagnetic
behaviour of oxygen confirms the presence of such electrons.
The intense reactivity of ozone is sometimes explained by suggesting that one of the
three oxygen atoms is in an “atomic” state; on reacting, this atom is dissociated from
the O
3
molecule, leaving molecular oxygen.
The molecular species, O
2
, is not especially reactive at normal (ambient)
temperatures and pressures. The atomic species, O, is far more reactive. The energy
of dissociation (O
2
→ 2O) is large at 117.2 kilocalories per mole.
Oxygen has an oxidation state of −2 in most of its compounds. It forms a large
range of covalently bonded compounds, among which are oxides of nonmetals, such
as water (H
2
O), sulfur dioxide (SO
2
), and carbon dioxide (CO
2
); organic compounds
such as alcohols, aldehydes, and carboxylic acids; common acids such as sulfuric
(H
2
SO
4
), carbonic (H
2
CO
3
), and nitric (HNO
3
); and corresponding salts, such
as
sodium
sulfate (Na
2
SO
4
), sodium carbonate (Na
2
CO
3
), and
sodium
nitrate
(NaNO
3
). Oxygen is present as the oxide ion, O
2-
, in the crystalline structure
of solid metallic oxides such as calcium oxide, CaO. Metallic superoxides, such as
potassium superoxide, KO
2
, contain the O
2
-
ion, whereas metallic peroxides, such as
barium peroxide, BaO
2
, contain the O
2
2-
ion.
