Selasa, 20 Januari 2009

oxygen

Oxygen (from the Greek roots ὀξύς (oxys) (acid, literally "sharp," from the taste of acids) and -γενής (-genēs) (producer, literally begetter) is the element with atomic number 8 and represented by the symbol O. It is a member of the chalcogen group on the periodic table, and is a highly reactive nonmetallic period 2 element that readily forms compounds (notably oxides) with almost all other elements. At standard temperature and pressure two atoms of the element bind to form dioxygen, a colorless, odorless, tasteless diatomic gas with the formula O2. Oxygen is the third most abundant element in the universe by mass after hydrogen and helium[1] and the most abundant element by mass in the Earth's crust.[2] Diatomic oxygen gas constitutes 20.9% of the volume of air.[3]

All major classes of structural molecules in living organisms, such as proteins, carbohydrates, and fats, contain oxygen, as do the major inorganic compounds that comprise animal shells, teeth, and bone. Oxygen in the form of O2 is produced from water by cyanobacteria, algae and plants during photosynthesis and is used in cellular respiration for all complex life. Oxygen is toxic to obligately anaerobic organisms, which were the dominant form of early life on Earth until O2 began to accumulate in the atmosphere 2.5 billion years ago.[4] Another form (allotrope) of oxygen, ozone (O3), helps protect the biosphere from ultraviolet radiation with the high-altitude ozone layer, but is a pollutant near the surface where it is a by-product of smog.

Oxygen was independently discovered by Carl Wilhelm Scheele, in Uppsala, in 1773 or earlier, and Joseph Priestley in Wiltshire, in 1774, but Priestley is often given priority because his publication came out of print first. The name oxygen was coined in 1777 by Antoine Lavoisier,[5] whose experiments with oxygen helped to discredit the then-popular phlogiston theory of combustion and corrosion. Oxygen is produced industrially by fractional distillation of liquefied air, use of zeolites to remove carbon dioxide and nitrogen from air, electrolysis of water and other means. Uses of oxygen include the production of steel, plastics and textiles; rocket propellant; oxygen therapy; and life support in aircraft, submarines, spaceflight and diving.Structure

At standard temperature and pressure, oxygen is a colorless, odorless gas with the molecular formula O2, in which the two oxygen atoms are chemically bonded to each other with a spin triplet electron configuration. This bond has a bond order of two, and is often simplified in description as a double bond[6] or as a combination of one two-electron bond and two three-electron bonds.[7]

Triplet oxygen is the ground state of the O2 molecule.[8] The electron configuration of the molecule has two unpaired electrons occupying two degenerate molecular orbitals.[9] These orbitals are classified as antibonding (weakening the bond order from three to two), so the diatomic oxygen bond is weaker than the diatomic nitrogen triple bond in which all bonding molecular orbitals are filled, but some antibonding orbitals are not.[8]

In normal triplet form, O2 molecules are paramagnetic—they form a magnet in the presence of a magnetic field—because of the spin magnetic moments of the unpaired electrons in the molecule, and the negative exchange energy between neighboring O2 molecules.[10] Liquid oxygen is attracted to a magnet to a sufficient extent that, in laboratory demonstrations, a bridge of liquid oxygen may be supported against its own weight between the poles of a powerful magnet.[11][12]

Singlet oxygen, a name given to several higher-energy species of molecular O2 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.[13] It is also produced in the troposphere by the photolysis of ozone by light of short wavelength,[14] and by the immune system as a source of active oxygen.[15] 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.[16]

Allotropes

Main article: Allotropes of oxygen

Ozone is a rare gas on Earth found mostly in the stratosphere.

The common allotrope of elemental oxygen on Earth is called dioxygen, O2. It has a bond length of 121 pm and a bond energy of 498 kJ·mol-1.[17] This is the form that is used by complex forms of life, such as animals, in cellular respiration (see Biological role) and is the form that is a major part of the Earth's atmosphere (see Occurrence). Other aspects of O2 are covered in the remainder of this article.

Trioxygen (O3) is usually known as ozone and is a very reactive allotrope of oxygen that is damaging to lung tissue.[18] Ozone is produced in the upper atmosphere when O2 combines with atomic oxygen made by the splitting of O2 by ultraviolet (UV) radiation.[5] Since ozone absorbs strongly in the UV region of the spectrum, the ozone layer of the upper atmosphere functions as a protective radiation shield for the planet.[5] Near the Earth's surface, however, it is a pollutant formed as a by-product of automobile exhaust.[19]

The metastable molecule tetraoxygen (O4) was discovered in 2001,[20][21] and was assumed to exist in one of the six phases of solid oxygen. It was proven in 2006 in that phase, created by pressurizing O2 to 20 GPa, is in fact a rhombohedral O8 cluster.[22] This cluster has the potential to be a much more powerful oxidizer than either O2 or O3 and may therefore be used in rocket fuel.[20][21] A metallic phase was discovered in 1990 when solid oxygen is subjected to a pressure of above 96 GPa[23] and it was shown in 1998 that at very low temperatures, this phase becomes superconducting.[24]

Physical properties

See also: Liquid oxygen and solid oxygen

Oxygen is more soluble in water than nitrogen; water contains approximately 1 molecule of O2 for every 2 molecules of N2, compared to an atmospheric ratio of approximately 1:4. The solubility of oxygen in water is temperature-dependent, and about twice as much (14.6 mg·L−1) dissolves at 0 °C than at 20 °C (7.6 mg·L−1).[25][26] At 25 °C and 1 atm of air, freshwater contains about 6.04 milliliters (mL) of oxygen per liter, whereas seawater contains about 4.95 mL per liter.[27] At 5 °C the solubility increases to 9.0 mL (50% more than at 25 °C) per liter for water and 7.2 mL (45% more) per liter for sea water.

Oxygen condenses at 90.20 K (−182.95 °C, −297.31 °F), and freezes at 54.36 K (−218.79 °C, −361.82 °F).[28] Both liquid and solid O2 are clear substances with a light sky-blue color caused by absorption in the red (in contrast with the blue color of the sky, which is due to Rayleigh scattering of blue light). High-purity liquid O2 is usually obtained by the fractional distillation of liquefied air;[29] Liquid oxygen may also be produced by condensation out of air, using liquid nitrogen as a coolant. It is a highly-reactive substance and must be segregated from combustible materials.