molecular oxygen

8nitrogenoxygenfluorine
-

O

S
General
Name, symbol, numberoxygen, O, 8
Chemical seriesnonmetals, chalcogens
Group, period, block162, p
Appearancecolorless (gas)
very pale blue (liquid)
Standard atomic weight15.9994(3) gmol−1
Electron configuration1s2 2s2 2p4
Electrons per shell2, 6
Physical properties
Phasegas
Density(0 C, 101.325 kPa)
1.429 g/L
Melting point54.36 K
(-218.79 °C, -361.82 °F)
Boiling point90.20 K
(-182.95 °C, -297.31 °F)
Critical point154.59 K, 5.043 MPa
Heat of fusion(O2) 0.444 kJmol−1
Heat of vaporization(O2) 6.82 kJmol−1
Heat capacity(25 C) (O2)
29.378 Jmol−1K−1
Vapor pressure
P/Pa1101001 k10 k100 k
at T/K   617390
Atomic properties
Crystal structurecubic
Oxidation states2, −1
(neutral oxide)
Electronegativity3.44 (Pauling scale)
Ionization energies
(more)
1st: 1313.9 kJmol−1
2nd: 3388.3 kJmol−1
3rd: 5300.5 kJmol−1
Atomic radius60 pm
Atomic radius (calc.)48 pm
Covalent radius73 pm
Van der Waals radius152 pm
Miscellaneous
Magnetic orderingparamagnetic
Thermal conductivity(300 K) 26.58 m Wm−1K−1
Speed of sound(gas, 27 C) 330 m/s
CAS registry number7782-44-7
Selected isotopes
Main article: Isotopes of oxygen
iso NA half-life DM DE (MeV) DP

16O99.76%O is stable with 8 neutrons
17O0.038%O is stable with 9 neutrons
18O0.21%O is stable with 10 neutrons
References
This box:     [ edit]


In science, oxygen (IPA: /ˈɒkˑsəˑdʒɪn/) is a chemical element with the chemical symbol O and atomic number 8. The word oxygen derives from two roots in Greek, οξύς (oxys) (acid, lit. sharp) and -γενής (-genēs) (producer, lit. begetter). It was recognized in 1777 by Antoine Lavoisier, who 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). Oxygen has a valency of 2. On Earth it is usually bonded to other elements covalently or ionically. Examples for common oxygen-containing compounds include water (H2O), sand (silica, SiO2), and rust (iron oxide, Fe2O3).

Diatomic oxygen (O2) is one of the two major components of air (20.95%). It is produced by plants during photosynthesis, and is necessary for aerobic respiration in animals. It is toxic to obligate anaerobic organisms and was a poisonous waste product for early life on Earth.

Triatomic oxygen (ozone, O3) forms through radiation in the upper layers of the atmosphere and acts as a shield against UV radiation.

Characteristics



At standard temperature and pressure, oxygen exists as a diatomic molecule with the formula O2, in which the two oxygen atoms are bonded to each other with the electron configuration of triplet oxygen. This bond has a bond order of two, and is thus often grossly simplified in description as a double bond.[2] Triplet oxygen is the ground state of the oxygen molecule. The electron configuration of the molecule has two unpaired electrons occupying two degenerate molecular orbitals. These orbitals are classified as antibonding, so the diatomic oxygen bond is weaker than the diatomic nitrogen bond, where all bonding molecular orbitals are filled. Though unpaired electrons are commonly associated with high reactivity in chemical compounds, triplet oxygen is relatively (and fortunately) nonreactive 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 O2 and solid O2 are clear substances with a light sky blue. The phenomena are not related; the color of the sky is due to Rayleigh scattering. 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 neighboring O2 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 O2 is usually obtained by the fractional distillation of liquid air or by the condensation out of air. It is a highly reactive substance and should be handled extremely carefully.

Oxygen is slightly soluble in water, but naturally occurring dissolved amounts are enough to support animal life (see below).

O2 has a bond length of 121 pm and a bond energy of 498 kJ/mol.[3]

Allotropes

The common allotrope of elemental oxygen on Earth, O2, is known as dioxygen. Elemental oxygen is most commonly encountered in this form, as 21% of Earth's atmosphere.

Ozone (O3), the less common triatomic allotrope of oxygen, is a poisonous gas with a distinct, sharp odor. It is thermodynamically unstable toward the more common dioxygen form. It is formed continuously in the upper atmosphere of the Earth by short-wave UV radiation, and also functions as a shield against UV radiation reaching the ground. Ozone has recently been found to be produced by the immune system as an antimicrobial (see below). Liquid and solid O3 (ozone) have a deeper blue color than ordinary oxygen, and they are unstable and explosive. Traces of it can be detected sometimes as a sharp smell coming from electromotors.

A newly discovered allotrope of oxygen, tetraoxygen (O4), is a deep red solid that is created by pressurizing O2 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 O2 or O3.[4][5] When tetraoxygen is subjected to a pressure of 96 GPa, it becomes metallic, similarly to hydrogen, and becomes more similar to the heavier chalcogens, such as tellurium and polonium, both of which show significant metallic character.

Applications

Uptake of oxygen from the air is the essential purpose of 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 augment the inspired oxygen partial pressure nearer to that found at sea-level.
Enlarge picture
A home oxygen concentrator in situ in an emphysema patient's house. The model shown is the DeVILBISS LT 4000.
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 ambient pressure of about one third 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 the issue of flexible spacesuits working at Earth sea-level pressures remains a technological challenge of aerospace technology.

Smelting of iron ore into steel consumes 55% of commercially produced oxygen.[5] In this process, oxygen is injected through a high-pressure lance into molten iron, which removes sulfur and carbon impurities. The reaction is exothermic, so the temperature increases to 1700 ° C.[5]

Another 25% of commercially produced oxygen is used by the chemical industry.[5] Ethylene is reacted with oxygen to create ethylene oxide, which in turn is converted into ethylene glycol; the primary feeder material used to manufacture a host of products, including antifreeze and polyester polymers (the precursors of many plastics and fabrics).[5]

Most of the remaining 20% of commercially produced oxygen is used in medical applications, metal cutting, as an oxidizer in rocket fuel, and in water treatment.[5] Oxygen is used in welding (such as the oxyacetylene torch).

Oxygen presents two spectrophotometric absorption bands peaking at 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.[6] 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 a pharmacological effect is doubtful, a metabolic boost being the most plausible explanation. Controlled tests of high oxygen mixtures in diving (see nitrox) and other activities, even at higher than normal pressures, demonstrated 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 temper its analgesic effect. A stable 50% gaseous mixture (Entonox) is commonly used in medicine today as an analgesic. However, the common basic anaesthetic mixture is 30% oxygen with 70% nitrous oxide; the pain-suppressing effects, obviously, are due to the nitrous oxide and not to oxygen.

History

Early experiments and Phlogiston theory

One of the first known experiments on the relationship between combustion and air was conducted by the 2nd century BCE Greek writer on mechanics Philo of Byzantium. In his work Pneumatica, Philo observed that inverting a vessel over a burning candle and surrounding the vessel's neck with water resulted in some water rising into the neck.[7] Philo incorrectly surmised that parts of the air in the vessel were converted into the classical element fire and thus were able to escape through pores in the glass. Many centuries later Leonardo da Vinci built on Philo's work by observing that a portion of air is consumed during combustion and respiration.[8]

Oxygen's discovery as a separate element was delayed by a philosophy of combustion and corrosion called the phlogiston theory. Established in 1667 by German alchemist J. J. Becher and modified by chemist Georg Ernst Stahl by 1731,[9] phlogiston theory stated that all combustible materials were made of two parts. One part, called phlogiston, was given off when the substance containing it was burned while the dephlogisticated part was thought to be its true form, its calx.[8] Highly combustible materials, such as coal, were made mostly of phlogiston while non-combustible substances, such as iron, contained very little. Air did not play a role in phlogiston theory and no initial quantitative experiments were conducted to test the idea; instead it was based on observations of what happened when something burned.[8]

In the late 16th century, Polish alchemist and philosopher Michał Sędziwój thought of the gas given off by warm niter (saltpeter) as "the elixir of life".

Robert Hooke, Ole Borch, Mikhail Lomonosov, and Pierre Bayen all also produced oxygen in experiments in the 17th century but none of them recognized it as an element.[10]

Discovery by Priestley and Scheele

Enlarge picture
Joseph Priestley is usually given priority in the discovery
An experiment conducted by British clergyman Joseph Priestley on August 1 1774 focused sunlight on mercuric oxide (HgO) inside a glass tube, which liberated a gas he named 'dephlogisticated air'.[11] He noted that candles burned brighter in the gas and that a mouse was more active and lived longer while breathing it. After breathing the gas himself, he wrote: "The feeling of it to my lungs was not sensibly different from that of common air, but I fancied that my breast felt peculiarly light and easy for some time afterwards."[10] Priestley published his findings in 1775 in a work titled Experiments and Observations on Different Kinds of Air.[8] Because he published first, Priestley is usually given priority in the discovery.

Enlarge picture
Carl Wilhelm Scheele beat Pristley to the discovery but published afterwards
Unknown to Priestley, Swedish pharmacist Carl Wilhelm Scheele had already produced oxygen by heating mercuric oxide and various nitrates some time around 1772.[8][11] Scheele wrote an account of this discovery in a manuscript he titled Treatise on Air and Fire, which he sent to his publisher in 1775. However, that document was not published until 1777.[12] Scheele called the gas 'fire air' because it was the only known supporter of combustion.

Noted French chemist Antoine Laurent Lavoisier later claimed to have independently discovered the new substance. However, Priestley visited Lavoisier in October 1774 and told him about his experiment and how he liberated the new gas. Scheele posted a letter to Lavoisier on September 30 1774 that described the discovery of the previously unknown substance, but Lavoisier never acknowledged receiving it (a copy of the letter was found in Scheele's belongings after his death).[12]

Lavoisier's contribution

Enlarge picture
Antoine Lavoisier discredited the Phlogiston theory
What Lavoisier did indisputably do was to conduct the first adequate quantitative experiments on oxidation and give the first correct explanation of how combustion works.[11] He used these and similar experiments, all started in 1774, to discredit the Phlogiston theory and to prove that the substance discovered by Priestley and Scheele was a chemical element.

In one experiment, Lavoisier observed that there was no overall increase in weight when tin and air were heated in a closed container.[11] He noted that air rushed in when he opened the container, which indicated that part of the trapped air had been consumed. He also noted that the tin had increased in weight and that increase was the same as the weight of the air that rushed back in. This and other experiments on combustion were documented in his book Sur la combustion en general, which was published in 1777.[11] In that work, he proved that air is a mixture of two gases; 'vital air', which is essential to combustion and respiration, and 'azote', which did not support either.

Lavoisier later renamed 'vital air' to oxygène after the Greek roots meaning "acid-former" while 'azote' was renamed nitrogen.[11] Oxygen entered the English language despite opposition by English scientists and the fact that Priestley had priority. This is partly due to a poem praising the gas titled "Oxygen" in the popular book The Botanic Garden (1791) by Erasmus Darwin, grandfather of Charles Darwin.[12]

Biological role

Enlarge picture
Delayed oxygen build-up in earth's atmosphere and oceans in reaction to the evolution of oxygenic photosynthesis: A) no oxygen produced by biosphere, B) oxygen produced, but absorbed in oceans and by seabed rock, C) oxygen starts to gas out of the oceans, but is absorbed by land surfaces and formation of ozone layer.
Molecular oxygen, O2, is essential for cellular respiration in all aerobic organisms. It is used as electron acceptor in the mitochondria to generate chemical energy in the form of adenosine triphosphate (ATP) during oxidative phosphorylation. During this reaction, oxygen is reduced to water. Conversely, free oxygen is produced in the biosphere through photolysis (light-driven oxidation and splitting) of water during photosynthesis in cyanobacteria, green algae and plants, thus closing the biological water-oxygen redox cycle.

Before the evolution of water oxidation in photosynthetic bacteria, oxygen was almost nonexistent in earth's atmosphere. Free oxygen first appeared in significant quantities 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). These organisms developed the mechanism of oxygen evolution between 3.5 and 2.7 billion years ago. At first, the produced oxygen dissolved in the oceans and reacted with iron. It started to "gas out" of the oxygen-saturated waters about 2.7 billion years ago as evident in the rusting of iron-rich terrestrial rocks starting around that time. The amount of oxygen in the atmosphere increased gradually at first and shot up rapidly around 2.2 to 1.7 billion years ago to about 10% of its present level.[13]

The development of an oxygen-rich atmosphere was one of the most important events in the history of life on earth. The presence of large amounts of dissolved and free oxygen in the oceans and atmosphere may have driven most of the anaerobic organisms then living to extinction during the oxygen catastrophe about 2.4 billion years ago. However, the high electronegativity of O2 creates a large potential energy drop for cellular respiration, thus enabling organisms using aerobic respiration to produce much more ATP than anaerobic organisms. This makes them so efficient that they have come to dominate earth's biosphere.[14] Photosynthesis and cellular respiration of oxygen allowed for the evolution of eukaryotic cells and ultimately complex multicellular organisms such as plants and animals.

The atmospheric abundance of free oxygen in later geological epochs and its gradual increase up to the present has been largely due to synthesis by photosynthetic organisms. Over the past 500 million years, oxygen levels fluctuated between 15 and 35% per volume. Towards the end of the Carboniferous era (coal age) about 300 million years ago, atmospheric oxygen levels reached a maximum of 35% by volume, allowing insects and amphibians with limiting respiratory systems to grow much larger than today's species. Today, oxygen is the second most common component of the earth's atmosphere (about 21% by volume) after nitrogen.

Occurrence



Oxygen is the third most abundant chemical element in the universe by mass, after hydrogen and helium (see chemical element). Some of this oxygen was produced during stellar nucleosynthesis as a step in the CNO-II branch of the CNO cycle. However oxygen is primarily produced in massive stars. In stars with at least four times the Sun's mass, 16O nuclei are produced during the Carbon burning process. 16O can also be produced in stars with at least 8 times the Sun's mass as a result of photodisintegration during the Neon burning process.[16]

Oxygen is the most common component of the Earth's crust (49% by mass),[17] the second most common component of the Earth as a whole (28% by mass), the most common component of the world's oceans (86% by mass), and the second most common component of the Earth's atmosphere (20.947% by volume), second to nitrogen.

Elemental oxygen occurs not only in the atmosphere, but also 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. [18]

See also , .

Production

Enlarge picture
Hoffman electrolysis apparatus used in electrolysis of water
Main article: Oxygen evolution
In nature, free oxygen is produced by the light-driven splitting of water during oxygenic photosynthesis in cyanobacteria, green algae and plants.[19] Algae and cyanobacteria in marine environments provide about 70% of the free oxygen produced on earth.[20] The remainder is produced by terrestrial plants, although almost all oxygen produced in tropical forests is consumed by organisms in those forests.[21]

Two major methods are employed to produce the 100 million tonnes of oxygen extracted from air for industrial uses annually.[12] The most common method is to fractionally distill liquefied air into its various components, with nitrogen distilling as a vapor while oxygen is left as a liquid.[12] The other major method of producing oxygen involves passing a stream of clean, dry air through a bed of zeolite molecular sieves, which absorb the nitrogen and leave a gas stream that is 90 to 93% oxygen.[12] Nitrogen is released from saturated zeolite by diverting air flow to another zeolite bed and reducing the chamber's air pressure. This allows for a continuous supply of gaseous oxygen to be pumped through a pipeline.

Oxygen can also be produced through electrolysis of water into oxygen and hydrogen. A similar method is the electrocatalytic oxygen evolution from oxides and oxoacids. Chemical catalysts can be used as well, such as in chemical oxygen generators or oxygen candles that are used as part of the life support equipment on spacecraft and submarines. Oxygen is increasingly obtained by non-cryogenic technologies such as pressure swing adsorption (PSA), vacuum-pressure swing adsorption (VPSA) [22], or vacuum swing adsorption (VSA) [1] technolgies. Air can be forced to dissolve through ceramic membranes based on zirconium oxide by either high pressure or an electric current to produce nearly pure oxygen.[5]

In large quantities, the price of liquid oxygen (2001) is approximately $0.21/kg. [24] Since the primary cost of production is the energy cost of liquefying the air, the production cost will change as energy cost varies.

Oxygen is often transported in bulk as a liquid in specially insulated tankers because one liter of liquefied oxygen is equivalent to 840 liters of the gas.[12] Oxygen is also stored and shipped in cylinders containing the compressed gas; a form that is useful in medical applications and Oxy-fuel welding and cutting.[12]

Compounds

Due to its electronegativity, oxygen forms chemical bonds with almost all other elements hence the original definition of oxidation. The only elements known to escape the possibility of oxidation are a few of the noble gases, and fluorine. However, many noble metals (common examples: gold, platinum) resist direct chemical combination with oxygen, and substances like gold oxide must be formed by an indirect route.

The most familiar oxygen compound is water. Other well-known examples include silica (found in sand, glass, rock, etc.), and the compounds of carbon and oxygen, such as carbon dioxide (CO2), alcohols (R-OH), carbonyls, (R-CO-H or R-CO-R), and carboxylic acids (R-COOH). Oxygenated radicals such as chlorates (ClO3), perchlorates (ClO4), chromates (CrO42−), dichromates (Cr2O72−), permanganates (MnO4), and nitrates (NO3) are strong oxidizing agents in and of themselves. Phosphorus is biologically important in its oxygenated form as the phosphate (PO43−) ion. Many metals bond with oxygen atoms, such as iron in iron(III) oxide (Fe2O3), commonly called rust.

Ozone (O3) is formed by electrostatic discharge in the presence of molecular oxygen. A double oxygen molecule (O2)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 O2+PtF6. It was discovered when Neil Bartlett was studying the properties of PtF6. He noticed a change in color when this compound was exposed to atmospheric air. Bartlett reasoned that xenon should be oxidized by PtF6. This led him to the discovery of xenon hexafluoroplatinate Xe+PtF6.

See also:

Isotopes

Main article: isotopes of oxygen
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. Nonetheless, 15O is used in positron emission tomography.

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 O2

Main article: oxygen toxicity


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.[25] 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 CO2 in the alveoli) is close to sea-level normal of 0.2 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 (O3), singlet oxygen, hydrogen peroxide, hydroxyl radicals and superoxide, are also highly toxic. Cells have developed various 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 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 Mn2+ 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.[26]

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. Oxygen itself is not the fuel, but as a reactant, concentrated oxygen may allow combustion to proceed dangerously rapidly. 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.)

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.

See also

References

1. ^ Note that the double bond depicted in the image is an oversimplification; see triplet oxygen
2. ^ Structure of Oxygen Molecule (triplet). Glasser Group, University of Missouri-Columbia. Retrieved on 2007-03-03.
3. ^ Chieh, Chung. Bond Lengths and Energies. University of Waterloo. Retrieved on 2007-03-03.
4. ^ Ball, Philip. "New form of oxygen found", news@nature.com, November 16, 2001. Retrieved on 2007-03-03. 
5. ^ F. Cacace, G. de Petris, A. Troiani, (2001). "Experimental Detection of Tetraoxygen". Angewandte Chemie International Edition 40 (21): 4062 - 4065. DOI:4062::AID-ANIE4062>3.0.CO;2-X 10.1002/1521-3773(20011105)40:21<4062::AID-ANIE4062>3.0.CO;2-X. 
6. ^ Zarco-Tejada, P.J., Miller, J.R.; Berger, M., Alonso, L., Cerovic, Z., Goulas, Y., Jacquemoud, S., Louis, J., Mohammed, G. Moya, I., Pedros, R., Moreno, J.F., Verhoef, W.. Progress on the development of an integrated canopy fluorescence model. Geoscience and Remote Sensing Symposium, 2003. IGARSS '03. Proceedings. 2003 IEEE International.
7. ^ Jastrow, Joseph (1936). Story of Human Error. Ayer Publishing, 171. ISBN 0836905687. 
8. ^ Cook (1969). The Encyclopedia of the Chemical Elements, "Oxygen", page 499
9. ^ Morris, Richard (2003). The last sorcerers: The path from alchemy to the periodic table (Hardback), Washington, D.C.: Joseph Henry Press. ISBN 0309089050. 
10. ^ Emsley (2001). Nature's Building Blocks, page 299
11. ^ Cook (1969). The Encyclopedia of the Chemical Elements, "Oxygen", page 500
12. ^ Emsley (2001). Nature's Building Blocks, page 300
13. ^ Campbell, Neil A.; Reece, Jane B. (2005). Biology, 7th Edition. San Francisco: Pearson - Benjamin Cummings, 522-523. ISBN 0-8053-7171-0. 
14. ^ Freeman, Scott (2005). Biological Science, 2nd Edition. Upper Saddle River, NJ: Pearson - Prentice Hall, 214, 586. ISBN 0-13-140941-7. 
15. ^ Data from the World Ocean Atlas 2001.
16. ^ Balachandran, S. C. (October 9-11, 1995). "Carbon and Oxygen Nucleosynthesis in the Galaxy: Problems and Prospects". Proceedings of the sixth (6th) annual October Astrophysics Conference: 188-195, College Park; Maryland: Astronomical Society of the Pacific. Retrieved on 2007-01-08. 
17. ^ Los Alamos National Laboratory – Oxygen
18. ^ From The Chemistry and Fertility of Sea Waters by H.W. Harvey, 1955, citing C.J.J. Fox, "On the coefficients of absorption of atmospheric gases in sea water", Publ. Circ. Cons. Explor. Mer, no. 41, 1907. Harvey however notes that according to later articles in Nature the values appear to be about 3% too high.
19. ^ Raven, Peter H.; Ray F. Evert, Susan E. Eichhorn (2005). Biology of Plants, 7th Edition. New York: W.H. Freeman and Company Publishers, 115-127. ISBN 0-7167-1007-2. 
20. ^ Fenical, William (September 1983). "Marine Plants: A Unique and Unexplored Resource", Plants: the potentials for extracting protein, medicines, and other useful chemicals (workshop proceedings). DIANE Publishing. ISBN 1428923977. 
21. ^ Broeker, W.S., 2006 "Breathing easy, Et tu, O2" Columbia University
22. ^ Non-Cryogenic Air Separation Processes 2003
23. ^ Emsley (2001). Nature's Building Blocks, page 301
24. ^ NASAFacts FS-2001-09-015-KSC, Space Shuttle Use of Propellants and Fluids, National Aeronautics and Space Administration, September 2001 (postscript file here
25. ^ Wade, Mark (2007). Space Suits. Encyclopedia Astronautica. Retrieved on 2006-08-10.
26. ^ Hoffmann, Roald (2004). "The Story of O". American Scientist 92 (1): 23. Retrieved on 2007-03-03. 

External links



3, 5, 4, 2
(strongly acidic oxide)
Electronegativity 3.04 (Pauling scale)
Ionization energies
(more) 1st: 1402.3 kJmol−1
2nd: 2856 kJmol−1
3rd: 4578.1 kJmol−1

Atomic radius 65 pm
Atomic radius (calc.
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100% F is stable with 10 neutrons
References
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6
(strongly acidic oxide)
Electronegativity 2.58 (Pauling scale)
Ionization energies
(more) 1st: 999.6 kJmol−1
2nd: 2252 kJmol−1
3rd: 3357 kJmol−1

Atomic radius 100 pm
Atomic radius (calc.
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An extended periodic table was suggested by Glenn T. Seaborg in 1969. It is a logical extension of the principles behind the standard periodic table to include possible undiscovered chemical elements.
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<onlyinclude> This is a list of chemical elements, sorted by name and color coded according to type of element.

Given is each element's element symbol, atomic number, atomic mass or most stable isotope, and group and period numbers on the periodic table.
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<onlyinclude> This is a list of chemical elements by symbol, including the current signification used to identify the chemical elements as recognized by the International Union of Pure and Applied Chemistry, as well as proposed and historical signs.
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A table of chemical elements ordered by atomic number and color coded according to type of element. Given is each element's name, element symbol, group and period, Chemical series, and atomic mass (or most stable isotope).
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A group, also known as a family, is a vertical column in the periodic table of the chemical elements. There are 18 groups in the standard periodic table.

The modern explanation of the pattern of the periodic table is that the elements in a group have similar
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Nonmetal is a term used in chemistry when classifying the chemical elements. On the basis of their general physical and chemical properties, every element in the periodic table can be termed either a metal or a non-metal.
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The chalcogens (with the "ch" pronounced with a hard "c" as in "chemistry") are the name for the periodic table group 16 (old-style: VIB or VIA) in the periodic table. It is sometimes known as the oxygen family.
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A group, also known as a family, is a vertical column in the periodic table of the chemical elements. There are 18 groups in the standard periodic table.

The modern explanation of the pattern of the periodic table is that the elements in a group have similar
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Periods:]] 1 2 3 4 5 6 7 8
Series Alkalis Alkaline earths Lanthanides Actinides Transition metals Poor metals Metalloids Nonmetals Halogens Noble gases
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A block of the periodic table of elements is a set of adjacent groups. The respective highest-energy electrons in each element in a block belong to the same atomic orbital type.
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The chalcogens (with the "ch" pronounced with a hard "c" as in "chemistry") are the name for the periodic table group 16 (old-style: VIB or VIA) in the periodic table. It is sometimes known as the oxygen family.
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A period 2 element is one of the chemical elements in the second row (or period) of the periodic table of the elements.

These are: Chemical elements in the second period
Group 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18
#
Name 3
Li 4
Be 5
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The p-block of the periodic table of the elements consists of the last six groups minus helium (which is located in the s-block). In the elemental form of the p-block elements, the highest energy electron occupies a p-orbital.
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Color or colour[1] (see spelling differences) is the visual perceptual property corresponding in humans to the categories called red, yellow, blue, black, etc.
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atomic mass (ma) is the mass of an atom at rest, most often expressed in unified atomic mass units.[1] The atomic mass may be considered to be the total mass of protons, neutrons and electrons in a single atom (when the atom is motionless).
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To help compare different orders of magnitude, the following list describes various mass levels between 10−36 kg and 1053 kg.

Factor (kg) Value Item
10−36 1.
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This is a list of chemical elements, sorted by relative atomic mass, or more precisely the standard atomic weights, (most stable isotope for artificial elements) and color coded according to type of element.
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electron configuration is the arrangement of electrons in an atom, molecule, or other physical structure (e.g., a crystal). Like other elementary particles, the electron is subject to the laws of quantum mechanics, and exhibits both particle-like and wave-like nature.
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Electron

Theoretical estimates of the electron density for the first few hydrogen atom electron orbitals shown as cross-sections with color-coded probability density
Composition: Elementary particle
Family: Fermion
Group: Lepton
Generation: First
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An electron shell, also known as a main energy level, is a group of atomic orbitals with the same value of the principal quantum number n. Electron shells are made up of one or more electron subshells, or sublevels
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In the physical sciences, a phase is a set of states of a macroscopic physical system that have relatively uniform chemical composition and physical properties (i.e. density, crystal structure, index of refraction, and so forth).
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Gas is one of the four major states of matter, consisting of freely moving atoms or molecules without a definite shape. Compared to the solid and liquid states of matter a gas has lower density and a lower viscosity.
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In physics, density is mass m per unit volume V—how heavy something is compared to its size. A small, heavy object, such as a rock or a lump of lead, is denser than a lighter object of the same size or a larger object of the same weight, such as pieces of
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The melting point of a crystalline solid is the temperature range at which it changes state from solid to liquid. Although the phrase would suggest a specific temperature and is commonly and incorrectly used as such in most textbooks and literature, most crystalline compounds
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The kelvin (symbol: K) is a unit increment of temperature and is one of the seven SI base units. The Kelvin scale is a thermodynamic (absolute) temperature scale where absolute zero — the coldest possible temperature — is zero kelvins
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Celsius is, or relates to, the Celsius temperature scale (previously known as the centigrade scale). The degree Celsius (symbol: °C) can refer to a specific temperature on the Celsius scale
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