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Oxygen

2007 Schools Wikipedia Selection. Related subjects: Chemical elements


                 8              nitrogen ← oxygen → fluorine
                 -
                ↑
                O
                ↓
                S

                                  Periodic Table - Extended Periodic Table

                                                                   General
                                         Name, Symbol, Number oxygen, O, 8
                                     Chemical series Nonmetals, chalcogens
                                             Group, Period, Block 16, 2, p
                                                Appearance colorless (gas)
                                                   very pale blue (liquid)
                                             Atomic mass 15.9994 (3) g/mol
                                     Electron configuration 1s^2 2s^2 2p^4
                                                  Electrons per shell 2, 6
                                                       Physical properties
                                                                 Phase gas
                                              Density (0 °C, 101.325 kPa)
                                                                 1.429 g/L
                                                    Melting point 54.36  K
                                              (-218.79 ° C, -361.82 ° F)
                                                     Boiling point 90.20 K
                                              (-182.95 ° C, -297.31 ° F)
                                        Critical point 154.59 K, 5.043 MPa
                                  Heat of fusion (O[2]) 0.444 kJ·mol^−1
                             Heat of vaporization (O[2]) 6.82 kJ·mol^−1
                                             Heat capacity (25 °C) (O[2])
                                                29.378 J·mol^−1·K^−1

   CAPTION: Vapor pressure

                                             P/Pa  1 10 100 1 k 10 k 100 k
                                            at T/K          61   73   90

                                                         Atomic properties
                                                   Crystal structure cubic
                                               Oxidation states −2, −1
                                                           (neutral oxide)
                                    Electronegativity 3.44 (Pauling scale)
                                                       Ionization energies
                                          ( more) 1st: 1313.9 kJ·mol^−1
                                                  2nd: 3388.3 kJ·mol^−1
                                                  3rd: 5300.5 kJ·mol^−1
                                                       Atomic radius 60 pm
                                               Atomic radius (calc.) 48 pm
                                                     Covalent radius 73 pm
                                               Van der Waals radius 152 pm
                                                             Miscellaneous
                                            Magnetic ordering paramagnetic
                     Thermal conductivity (300 K) 26.58 mW·m^−1·K^−1
                                      Speed of sound (gas, 27 °C) 330 m/s
                                             CAS registry number 7782-44-7
                                                         Selected isotopes

                  CAPTION: Main article: Isotopes of oxygen

                                  iso    NA   half-life DM DE ( MeV) DP
                                  ^16O 99.76% O is stable with 8 neutrons
                                  ^17O 0.038% O is stable with 9 neutrons
                                  ^18O 0.21%  O is stable with 10 neutrons

                                                                References

   Oxygen ( IPA: /ˈɔksidʒən/) is a chemical element with the chemical
   symbol O and atomic number 8. On Earth, it is usually covalently or
   ionically bonded to other elements.

   Unbound oxygen (also called molecular dioxygen, O[2], a diatomic
   molecule) first appeared in significant quantities on Earth during the
   Paleoproterozoic era (between 2.5 billion years ago and 1.6 billion
   years ago) as a product of the metabolic action of early anaerobes (
   archaea and bacteria). The presence of large amounts of free oxygen may
   have driven most of the organisms then living to extinction. The
   atmospheric abundance of free oxygen in later geological epochs and up
   to the present has been largely driven by photosynthetic organisms;
   roughly three quarters of the free element being produced by algae in
   the oceans, and one quarter from terrestrial plants.

Characteristics

   By an interesting coincidence of nature, liquid oxygen has a sky blue
   color. It is important to note however, that the two phenomena are
   totally unrelated (the blue of sky is due to Rayleigh scattering and
   would be present even if there were no oxygen in air).
   Enlarge
   By an interesting coincidence of nature, liquid oxygen has a sky blue
   colour. It is important to note however, that the two phenomena are
   totally unrelated (the blue of sky is due to Rayleigh scattering and
   would be present even if there were no oxygen in air).

   Oxygen is a major component of air, produced by plants during
   photosynthesis, and is necessary for aerobic respiration in animals.
   The word oxygen derives from two roots in Greek, οξυς (oxys) (acid,
   sharp) and -γενης (-genēs) (born of). In the early 18th century,
   Antoine Lavoisier coined the name oxygen from the Greek roots mentioned
   above because he erroneously thought that it was a constituent of all
   acids. (The definition of acid has since been revised).

   At standard temperature and pressure, oxygen exists as a diatomic
   molecule with the formula O[2], in which the two oxygen atoms are
   doubly bonded to each other. In its most stable form, oxygen exists as
   a diradical ( triplet oxygen) with two unpaired electrons in molecular
   orbitals of the O[2] molecule. Though unpaired electrons are commonly
   associated with high reactivity in chemical compounds, triplet oxygen
   is relatively (and fortunately) unreactive by comparison with most
   radicals.

   Singlet oxygen, a name given to several higher energy species of
   molecular oxygen in which all the electron spins are paired, is much
   more reactive towards common organic molecules. In nature, singlet
   oxygen is commonly formed from water during photosynthesis, using the
   energy of sunlight. It is also produced by the immune system as a
   source of active oxygen. Carotenoids in photosynthetic organisms and
   possibly also in animals, play a major role in absorbing energy from
   singlet oxygen and converting it to the unexcited ground state, before
   it can cause harm to tissues.

   Liquid O[2] and solid O[2] are clear substances with a light sky-blue
   colour. In normal triplet form they are paramagnetic due to the spin
   magnetic moments of the unpaired electrons in the molecule, and the
   negative exchange energy between neighbouring O[2] molecules. Liquid
   oxygen is attracted to a magnet to a sufficient extent that a bridge of
   liquid oxygen may be supported against its own weight between the poles
   of a powerful magnet, in laboratory demonstrations. Liquid O[2] is
   usually obtained by the fractional distillation of liquid air.

   Oxygen is slightly soluble in water, but naturally occuring disolved
   amounts support all ocean animal life (see below).

   O[2] has a bond length of 121 pm and a bond energy of 498 kJ/mol.

Allotropes

   Dioxygen, O2, is a gas and consists of 2 oxygen atoms. Oxygen is most
   commonly encountered in this form, as it makes up 21% of the
   atmosphere.
   Enlarge
   Dioxygen, O[2], is a gas and consists of 2 oxygen atoms. Oxygen is most
   commonly encountered in this form, as it makes up 21% of the
   atmosphere.
   Ozone, O3, is a gas and consists of 3 oxygen atoms
   Enlarge
   Ozone, O[3], is a gas and consists of 3 oxygen atoms

   Ozone, the triatomic allotrope of oxygen, is a poisonous gas with a
   sharp odour. It functions in the upper atmosphere of the Earth as a
   shield against UV radiation, and has recently been found to be produced
   by the immune system as an antimicrobial (see below). Liquid and solid
   O[3] (ozone) have a deeper blue colour than ordinary oxygen, and they
   are unstable and explosive.

   A recently discovered allotrope of oxygen, tetraoxygen (O[4]), is a
   deep red solid that is created by pressurizing O[2] to the order of 20
   GPa. Its properties are being studied for use in rocket fuels and
   similar applications, as it is a much more powerful oxidizer than
   either O[2] or O[3].

Applications

   Oxygen is essential to respiration, so oxygen supplementation has found
   use in medicine (as oxygen therapy). People who climb mountains or fly
   in non-pressurized aeroplanes sometimes have supplemental oxygen
   supplies; the reason is that increasing the proportion of oxygen in the
   breathing gas at low pressure acts to increase the inspired oxygen
   partial pressure nearer to that found at sea-level. A notable
   application of oxygen as a very low-pressure breathing gas, is in
   modern spacesuits, where use of nearly pure oxygen at a total pressure
   of about 1/3rd normal, results in normal blood partial pressures of
   oxygen. This trade-off of breathing gas content and needed pressure is
   important for space applications, because flexible spacesuits working
   at Earth sea-level pressures remains a technological challenge beyond
   today's capabilities.

   Oxygen is used in welding (such as the oxyacetylene torch), and in the
   making of steel and methanol. Liquid oxygen finds use as a classic
   oxidizer in rocket propulsion.

   Oxygen presents two absorption bands centered in the wavelengths 687
   and 760 nanometers. Some scientists have proposed to use the
   measurement of the radiance coming from vegetation canopies in those
   oxygen bands to characterize plant health status from a satellite
   platform. This is because in those bands, it is possible to
   discriminate the vegetation's reflectance from the vegetation's
   fluorescence, which is much weaker. The measurement presents several
   technical difficulties due to the low signal to noise ratio and due to
   the vegetation's architecture, but it has been proposed as a
   possibility to monitor the carbon cycle from satellites on a global
   scale.

   Oxygen, as a supposed mild euphoric, has a history of recreational use
   (see oxygen bar), however the reality of this effect is doubtful.
   Controlled tests of high oxygen mixtures in diving (see nitrox) and
   other activities, even at higher than normal pressures, show no
   particular effects on humans other than promotion of an increased
   tolerance to aerobic exercise.

   In the 19th century, oxygen was often mixed with nitrous oxide to
   promote an analgesic effect; a stable 50% gaseous mixture ( Entonox) is
   commonly used in medicine today as an analgesic, and 30% oxygen with
   70% nitrous oxide is the common basic anaesthetic mixture. These
   effects, however, are due to the nitrous oxide.

Scientific history

   Oxygen was first described by Michał Sędziwój, a Polish alchemist and
   philosopher in the late 16th century. Sędziwój thought of the gas given
   off by warm nitre (saltpeter) as "the elixir of life".

   Oxygen was more quantitatively discovered by the Swedish pharmacist
   Carl Wilhelm Scheele some time before 1773, but the discovery was not
   published until after the independent discovery by Joseph Priestley on
   August 1, 1774, who called the gas dephlogisticated air (see phlogiston
   theory). Priestley published discoveries in 1775 and Scheele in 1777;
   consequently Priestley is usually given the credit. Both Scheele and
   Priestley produced oxygen by heating mercuric oxide.

   Scheele called the gas 'fire air' because it was the only known
   supporter of combustion. It was later called 'vital air' because it was
   and is vital for the existence of animal life.

   The gas was named by Antoine Laurent Lavoisier, after Priestley's
   publication in 1775, from Greek roots meaning " acid-former". As noted,
   the name reflects the then-common incorrect belief that acids contain
   oxygen.

Occurrence

   Annual mean sea surface dissolved oxygen for the World Ocean. Note more
   oxygen in cold water near the poles.
   Enlarge
   Annual mean sea surface dissolved oxygen for the World Ocean. Note more
   oxygen in cold water near the poles.

   Oxygen is the most common component of the Earth's crust (49% by mass),
   the second most common component of the Earth as a whole (28.2% by
   mass), and the second most common component of the Earth's atmosphere
   (20.947% by volume), second to nitrogen.

   Oxygen occurs as solution in the world's water bodies. At 25° C under 1
   atm of air, a litre of water will dissolve about 6.04 cc (8.63 mg,
   0.270 mmol) of oxygen, whereas sea water will dissolve about 4.9 cc
   (7.0 mg, 0.22 mmol). At 0° C the solubilities increase to 10.29 cc
   (14.7 mg, 0.460 mmol) for water and 8.0 cc (11.4 mg, 0.36 mmol) for sea
   water. This difference has important implications for ocean life, as
   polar oceans support a much higher density of life due to their oxygen
   content.

Compounds

   The most familiar of oxygen compounds is water.

   Due to its electronegativity, oxygen forms chemical bonds with almost
   all other elements hence the origin of the original definition of
   oxidation. The only elements known to escape the possibility of
   oxidation are a few of the noble gases, and fluorine. Other than water
   (H[2]O), well known examples include compounds of carbon and oxygen,
   such as carbon dioxide (CO[2]), alcohols (R-OH), carbonyls, (R-CO-H or
   R-CO-R)), and carboxylic acids (R-COOH). Oxygenated radicals such as
   chlorates (ClO[3]^−), perchlorates (ClO[4]^−), chromates (CrO[4]^2−),
   dichromates (Cr[2]O[7]^2−), permanganates (MnO[4]^−), and nitrates
   (NO[3]^−) are strong oxidizing agents in and of themselves. Many metals
   such as iron bond with oxygen atoms, iron(III) oxide (Fe[2]O[3]). Ozone
   (O[3]) is formed by electrostatic discharge in the presence of
   molecular oxygen. A double oxygen molecule (O[2])[2] is known and is
   found as a minor component of liquid oxygen. Epoxides are ethers in
   which the oxygen atom is part of a ring of three atoms.

   One unexpected oxygen compound is dioxygen hexafluoroplatinate
   O[2]^+PtF[6]^−. It was discovered when Neil Bartlett was studying the
   properties of PtF[6]. He noticed a change in colour when this compound
   was exposed to atmospheric air. Bartlett reasoned that xenon should be
   oxidized by PtF[6]. This led him to the discovery of xenon
   hexafluoroplatinate Xe^+PtF[6]^−.

Isotopes

   Oxygen has seventeen known isotopes with atomic masses ranging from
   12.03 u to 28.06 u. Three are stable, ^16O, ^17O, and ^18O, of which
   ^16O is the most abundant (over 99.7%). The radioisotopes all have
   half-lives of less than three minutes.

   An atomic weight of 16 was assigned to oxygen prior to the definition
   of the unified atomic mass unit based upon ^12C. Since physicists
   referred to ^16O only, while chemists meant the naturally abundant
   mixture of isotopes, this led to slightly different atomic weight
   scales.

Precautions

Toxicity of O[2]

   Oxygen can be toxic at elevated partial pressures. Since oxygen partial
   pressure is the fraction of oxygen times the total pressure, elevated
   partial pressures can occur either from high oxygen fraction in
   breathing gas, or from high breathing gas pressure, or a combination of
   both. Oxygen toxicity usually begins to occur at partial pressures more
   than 0.5 atmospheres, or 2.5 times the normal sea-level oxygen partial
   pressure of about 0.2 atmospheres or bars. This means that at sea-level
   pressures, mixtures containing less than 50% oxygen are essentially
   non-toxic. However in medical applications (such as in ventilation gas
   mixtures in hospital applications) mixtures containing more than 50%
   oxygen can be expected to show lung toxicity, causing slow damage to
   the lungs over periods of days, with the rate of damage rising rapidly
   from mixtures between 50% and 100% oxygen. On the other hand, breathing
   100% oxygen in space applications (such as in some modern spacesuits,
   or in early spacecraft such as the Apollo spacecraft), causes no damage
   due to the low total pressures (30% to 33% sea-level) used . In the
   case of spacesuits, oxygen partial pressure in the breathing gas is
   typically about 0.30 bar (1.4 times normal), and oxygen partial
   pressure in the astronaut's blood (due to downward adjustments due to
   water vapor and CO[2] in the alveoli) is close to sea-level normal of
   0.14 bar.

   In deep scuba diving and surface supplied diving and when using
   equipment which can provide high partial pressures of oxygen, such as
   rebreathers, oxygen toxicity to the lungs can occur, just as in medical
   applications. Due to the higher total pressures in these applications,
   the fraction of oxygen which produces lung damage may be considerably
   less than 50%. More importantly, under pressures higher than normal
   sea-level, a far more serious form of oxygen toxicity in the central
   nervous system may lead to generalized seizures. This form of oxygen
   toxicity usually occurs after several hours exposure to oxygen partial
   pressures over about 1.4 atmospheres (bars) (i.e. 7 times normal), with
   the time decreasing for higher pressures above this, and with great
   variation from person to person. At over three bars of oxygen partial
   pressure (15 times normal), seizures typically occur within minutes.

Toxicity and antibacterial use of other chemical oxygen forms

   Certain derivatives of oxygen, such as ozone (O[3]), singlet oxygen,
   hydrogen peroxide, hydroxyl radicals and superoxide, are also highly
   toxic. The body has developed mechanisms to protect against all of
   these toxic compounds. For instance, the naturally-occurring
   glutathione can act as an antioxidant, as can bilirubin which is
   normally a breakdown product of hemoglobin. To protect against the
   destructive nature of peroxides, nearly every organism on earth has
   developed some form of the enzyme catalase, which very quickly
   disproportionates hydrogen peroxide into water and dioxygen. Another
   nearly universally present enzyme in living organisms (except for a few
   species of bacteria which use Mn^2+ ions directly for the job) is
   superoxide dismutase. This family of enzymes disproportionates
   superoxide to oxygen and peroxide, which is then in turn dealt with, by
   catalase.

   Immune systems of higher organisms have long made use of reactive forms
   of oxygen which they produce. Not only do antibodies catalyze
   production of peroxide from oxygen, it is now known that immune cells
   produce peroxide, superoxide, and singlet oxygen in the course of an
   immune response. Recently, singlet oxygen has been found to be a source
   of biologically-produced ozone: this reaction proceeds through an
   unusual compound dihydrogen trioxide, also known as trioxidane, (HOOOH)
   which is an antibody-catalyzed product of singlet oxygen and water.
   This compound in turn disproportionates to ozone and peroxide,
   providing two powerful antibacterials. The body's range of defense
   against all of these active oxidizing agents is hardly surprising,
   then, given their "deliberate" employment as antimicrobial agents in
   the immune response.

   Oxygen derivatives are prone to form free radicals, especially in
   metabolic processes. Because they can cause severe damage to cells and
   their DNA before they are dealt with, they form part of many theories
   of carcinogenesis and aging.

Combustion hazard

   Highly concentrated sources of oxygen promote rapid combustion and
   therefore are fire and explosion hazards in the presence of fuels. The
   fire that killed the Apollo 1 crew on a test launchpad spread so
   rapidly because the capsule was pressurized with pure oxygen as would
   be usual in an actual flight, but to maintain positive pressure in the
   capsule, this was at slightly more than atmospheric pressure instead of
   the ⅓ normal pressure that would be used in flight. (See partial
   pressure.)

   Similar hazards also apply to compounds of oxygen with a high oxidative
   potential, such as high concentration peroxides, chlorates,
   perchlorates, and dichromates; they also can often cause chemical
   burns.

   Retrieved from " http://en.wikipedia.org/wiki/Oxygen"
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