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 . Oxygen is the
third most abundant element in the universe by mass after
hydrogen and
heliumEmsley 2001, p.297 and the
most abundant element by mass in the
Earth's crust. Diatomic oxygen gas constitutes 20.9% of the volume of
air.
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 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 began to accumulate in the atmosphere 2.5 billion years ago. Another form (
allotrope) of oxygen,
ozone (), 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. At even higher
low earth orbit altitudes atomic oxygen is a significant presence and a cause of
erosion for spacecraft.
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 in print first. The name oxygen was coined in 1777 by
Antoine Lavoisier, 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.
Characteristics
Structure
At
standard temperature and pressure, oxygen is a colorless, odorless gas with the
molecular formula , 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 or as a combination of one two-electron bond and two
three-electron bonds.
Triplet oxygen (not to be confused with
ozone, ) is the
ground state of the molecule. The electron configuration of the molecule has two unpaired electrons occupying two
degenerate molecular orbitals.An orbital is a concept from
quantum mechanics that models an electron as a
wave-like particle that has a spacial distribution about an atom or molecule. 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.
In normal triplet form, 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 molecules. 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.Oxygen's paramagnetism can be used analytically in paramagnetic oxygen gas analysers that determine the purity of gaseous oxygen. ()
Singlet oxygen, a name given to several higher-energy species of molecular 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 in the
troposphere by the photolysis of ozone by light of short wavelength, and 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.
Allotropes
The common
allotrope of elemental oxygen on Earth is called dioxygen, . It has a bond length of 121
pm and a bond energy of 498
kJ·mol-1. 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 are covered in the remainder of this article.
Trioxygen () is usually known as
ozone and is a very reactive allotrope of oxygen that is damaging to lung tissue. Ozone is produced in the
upper atmosphere when combines with atomic oxygen made by the splitting of by
ultraviolet (UV) radiation. 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. Near the Earth's surface, however, it is a
pollutant formed as a by-product of automobile exhaust. The
metastable molecule
tetraoxygen () was discovered in 2001, and was assumed to exist in one of the six phases of
solid oxygen. It was proven in 2006 that this phase, created by pressurizing to 20
GPa, is in fact a
rhombohedral cluster. molecular lattice in the phase of solid oxygen|journal=Nature|volume=443|pages=201–04|doi=10.1038/nature05174|first=Lars F. |last=Lundegaard| coauthors=Weck, Gunnar; McMahon, Malcolm I.; Desgreniers, Serge and Loubeyre, Paul|year=2006--> This cluster has the potential to be a much more powerful
oxidizer than either or and may therefore be used in
rocket fuel. A metallic phase was discovered in 1990 when solid oxygen is subjected to a pressure of above 96 GPa and it was shown in 1998 that at very low temperatures, this phase becomes
superconducting.
Physical properties
Oxygen is more
soluble in water than nitrogen; water contains approximately 1 molecule of for every 2 molecules of , 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). At 25 °C and of air, freshwater contains about 6.04
milliliter (mL) of oxygen per
liter, whereas
seawater contains about 4.95 mL per liter. 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). Both
liquid and
solid 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 is usually obtained by the
fractional distillation of liquefied air; 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.
Isotopes and stellar origin
Naturally occurring oxygen is composed of three stable
isotopes, 16O, 17O, and
18O, with 16O being the most abundant (99.762%
natural abundance).
Most 16O is
synthesized at the end of the
helium fusion process in
stars but some is made in the
neon burning process. 17O is primarily made by the burning of hydrogen into
helium during the
CNO cycle, making it a common isotope in the hydrogen burning zones of stars. Most 18O is produced when
14N (made abundant from CNO burning) captures a
4He nucleus, making 18O common in the helium-rich zones of stars.
Fourteen
radioisotopes have been characterized, the most stable being 15O with a
half-life of 122.24 seconds (s) and 14O with a half-life of 70.606 s. All of the remaining
radioactive isotopes have half-lives that are less than 27 s and the majority of these have half-lives that are less than 83 milliseconds. The most common
decay mode of the isotopes lighter than 16O is
β+ decay to yield nitrogen, and the most common mode for the isotopes heavier than 18O is
beta decay to yield
fluorine.
Occurrence
Oxygen is the most abundant chemical element, by mass, in our biosphere, air, sea and land.Oxygen is the third most abundant chemical element in the universe, after hydrogen and helium. About 0.9% of the
Sun's mass is oxygen. Oxygen constitutes 49.2% of the
Earth's crust by mass
and is the major component of the world's oceans (88.8% by mass). Oxygen gas is the second most common component of the
Earth's atmosphere, taking up 21.0% of its volume and 23.1% of its mass (some 1015 tonnes).
Emsley 2001, p.298Figures given are for values up to above the surface Earth is unusual among the planets of the
Solar System in having such a high concentration of oxygen gas in its atmosphere:
Mars (with 0.1% by volume) and
Venus have far lower concentrations. However, the surrounding these other planets is produced solely by ultraviolet radiation impacting oxygen-containing molecules such as carbon dioxide.
thumb|Cold water holds more dissolved .|alt=World map showing that the sea-surface oxygen is depleted around the equator and increases towards the poles.The unusually high concentration of oxygen gas on Earth is the result of the
oxygen cycle. This
biogeochemical cycle describes the movement of oxygen within and between its three main reservoirs on Earth: the atmosphere, the
biosphere, and the
lithosphere. The main driving factor of the oxygen cycle is
photosynthesis, which is responsible for modern Earth's atmosphere. Photosynthesis releases oxygen into the atmosphere, while
respiration and
decay remove it from the atmosphere. In the present equilibrium, production and consumption occur at the same rate of roughly 1/2000th of the entire atmospheric oxygen per year.
Free oxygen also occurs in solution in the world's water bodies. The increased solubility of at lower temperatures (see
Physical properties) has important implications for ocean life, as polar oceans support a much higher density of life due to their higher oxygen content.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.
Polluted water may have reduced amounts of in it, depleted by decaying algae and other biomaterials (see
eutrophication). Scientists assess this aspect of water quality by measuring the water's
biochemical oxygen demand, or the amount of needed to restore it to a normal concentration.
Emsley 2001, p.301
Biological role
Photosynthesis and respiration
In nature, free oxygen is produced by the
light-driven splitting of water during oxygenic
photosynthesis.
Green algae and
cyanobacteria in marine environments provide about 70% of the free oxygen produced on earth and the rest is produced by terrestrial plants.
A simplified overall formula for photosynthesis is:
- : 6 + 6 + photons → + 6 (or simply carbon dioxide + water + sunlight → glucose + dioxygen)
Photolytic
oxygen evolution occurs in the
thylakoid membranes of photosynthetic organisms and requires the energy of four
photons.Thylakoid membranes are part of
chloroplasts in algae and plants while they simply are one of many membrane structures in cyanobacteria. In fact, chloroplasts are thought to have evolved from
cyanobacteria that were once symbiotic partners with the progenerators of plants and algae. Many steps are involved, but the result is the formation of a
proton gradient across the thylakoid membrane, which is used to synthesize
ATP via
photophosphorylation.
Raven 2005, 115–27 The remaining after oxidation of the water molecule is released into the atmosphere.Water oxidation is catalyzed by a
manganese-containing
enzyme complex known as the
oxygen evolving complex (OEC) or water-splitting complex found associated with the lumenal side of thylakoid membranes. Manganese is an important
cofactor, and
calcium and
chloride are also required for the reaction to occur.(Raven 2005)
Molecular dioxygen, , is essential for cellular respiration in all
aerobic organisms. Oxygen is used in
mitochondria to help generate
adenosine triphosphate (ATP) during
oxidative phosphorylation. The reaction for aerobic respiration is essentially the reverse of photosynthesis and is simplified as:
- : + 6 → 6 + 6 + 2880 kJ·mol-1
In
vertebrates, is
diffused through membranes in the lungs and into
red blood cells.
Hemoglobin binds , changing its color from bluish red to bright red. is released from another part of hemoglobin (see
Bohr effect) Other animals use
hemocyanin (
mollusc and some
arthropods) or
hemerythrin (
spiders and
lobsters). A liter of blood can dissolve 200 cm3 of .
Reactive oxygen species, such as
superoxide ion () and
hydrogen peroxide (), are dangerous by-products of oxygen use in organisms. Parts of the
immune system of higher organisms, however, create peroxide, superoxide, and singlet oxygen to destroy invading microbes. Reactive oxygen species also play an important role in the
hypersensitive response of plants against pathogen attack.
An adult human in rest
inhales 1.8 to 2.4 grams of oxygen per minute.
"For humans, the normal volume is 6-8 liters per minute." This amounts to more than 6 billion tonnes of oxygen inhaled by humanity per year.(1.8 grams/min/person)×(60 min/h)×(24 h/day)×(365 days/year)×(6.6 billion people)/1,000,000 g/t=6.24 billion tonnes
Build-up in the atmosphere
Free oxygen gas was almost nonexistent in
Earth's atmosphere before photosynthetic
archaea and
bacteria evolved. Free oxygen first appeared in significant quantities during the
Paleoproterozoic era (between 2.5 and 1.6 billion years ago). At first, the oxygen combined with dissolved
iron in the oceans to form
banded iron formations. Free oxygen started to gas out of the oceans 2.7 billion years ago, reaching 10% of its present level around 1.7 billion years ago.
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,
cellular respiration using enables
aerobic organisms to produce much more ATP than anaerobic organisms, helping the former to dominate Earth's
biosphere. Photosynthesis and cellular respiration of allowed for the evolution of
eukaryotic cells and ultimately complex multicellular organisms such as plants and animals.
Since the beginning of the
Cambrian era 540 million years ago, levels have fluctuated between 15% and 30% by volume. Towards the end of the
Carboniferous era (about 300 million years ago) atmospheric levels reached a maximum of 35% by volume, which may have contributed to the large size of insects and amphibians at this time. Human activities, including the burning of 7 billion
tonnes of
fossil fuels each year have had very little effect on the amount of free oxygen in the atmosphere. At the current rate of photosynthesis it would take about 2,000 years to regenerate the entire in the present atmosphere.
History
Early experiments
One of the first known experiments on the relationship between
combustion and air was conducted by the second 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.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.
Cook & Lauer 1968, p.499.
In the late 17th century,
Robert Boyle proved that air is necessary for combustion. English chemist
John Mayow refined this work by showing that fire requires only a part of air that he called spiritus nitroaereus or just nitroaereus.In one experiment he found that placing either a mouse or a lit candle in a closed container over water caused the water to rise and replace one-fourteenth of the air's volume before extinguishing the subjects.From this he surmised that nitroaereus is consumed in both
respiration and combustion.
Mayow observed that
antimony increased in weight when heated, and inferred that the nitroaereus must have combined with it. He also thought that the lungs separate nitroaereus from air and pass it into the blood and that animal heat and muscle movement result from the reaction of nitroaereus with certain substances in the body. Accounts of these and other experiments and ideas were published in 1668 in his work Tractatus duo in the tract "De respiratione".
Phlogiston theory
Discovery
Oxygen was first discovered by
Swedish pharmacist
Carl Wilhelm Scheele. He had produced oxygen gas by heating mercuric oxide and various
nitrates by about 1772. Scheele called the gas 'fire air' because it was the only known supporter of combustion, and 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.
Emsley 2001, p.300
In the meantime, on August 1, 1774, an experiment conducted by the
British clergyman
Joseph Priestley focused sunlight on
mercuric oxide (HgO) inside a glass tube, which liberated a gas he named 'dephlogisticated air'.
Cook & Lauer 1968, p.500 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." Priestley published his findings in 1775 in a paper titled "An Account of Further Discoveries in Air" which was included in the second volume of his book titled
Experiments and Observations on Different Kinds of Air. Because he published his findings first, Priestley is usually given priority in the discovery.
The noted French chemist
Antoine Laurent Lavoisier later claimed to have discovered the new substance independently. However, Priestley visited Lavoisier in October 1774 and told him about his experiment and how he liberated the new gas. Scheele also posted a letter to Lavoisier on September 30, 1774 that described his own 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).
Lavoisier's contribution
What Lavoisier did indisputably do (although this was disputed at the time) was to conduct the first adequate quantitative experiments on
oxidation and give the first correct explanation of how combustion works. 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. 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 général, which was published in 1777. In that work, he proved that air is a mixture of two gases; 'vital air', which is essential to combustion and respiration, and azote (Gk. "lifeless"), which did not support either. Azote later became
nitrogen in English, although it has kept the name in French and several other European languages.
Lavoisier renamed 'vital air' to oxygène in 1777 from the
Greek roots (oxys) (
acid, literally "sharp," from the taste of acids) and -γενής (-genēs) (producer, literally begetter), because he mistakenly believed that oxygen was a constituent of all acids. Chemists eventually determined that Lavoisier was wrong in this regard, but by that time it was too late, the name had taken. Actually, the gas that could more appropriately have been given the description, "acid producer," is hydrogen.
Oxygen entered the English language despite opposition by English scientists and the fact that the Englishman Priestley had first isolated the gas and written about it. 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.
Later history
John Dalton's original
atomic hypothesis assumed that all elements were monoatomic and that the atoms in compounds would normally have the simplest atomic ratios with respect to one another. For example, Dalton assumed that water's formula was HO, giving the
atomic mass of oxygen as 8 times that of hydrogen, instead of the modern value of about 16. In 1805,
Joseph Louis Gay-Lussac and
Alexander von Humboldt showed that water is formed of two volumes of hydrogen and one volume of oxygen; and by 1811
Amedeo Avogadro had arrived at the correct interpretation of water's composition, based on what is now called
Avogadro's law and the assumption of diatomic elemental molecules.However, these results were mostly ignored until 1860. Part of this rejection was due to the belief that atoms of one element would have no
chemical affinity towards atoms of the same element, and part was due to apparent exceptions to Avogadro's law that were not explained until later in terms of dissociating molecules.
By the late 19th century scientists realized that air could be liquefied, and its components isolated, by compressing and cooling it. Using a
cascade method, Swiss chemist and physicist
Raoul Pierre Pictet evaporated liquid
sulfur dioxide in order to liquefy carbon dioxide, which in turn was evaporated to cool oxygen gas enough to liquefy it. He sent a telegram on December 22, 1877 to the
French Academy of Sciences in Paris announcing his discovery of
liquid oxygen. Just two days later, French physicist
Louis Paul Cailletet announced his own method of liquefying molecular oxygen. Only a few drops of the liquid were produced in either case so no meaningful analysis could be conducted. Oxygen was liquified in stable state for the first time on March 29, 1877 by Polish scientists from
Jagiellonian University,
Zygmunt Wróblewski and
Karol Olszewski.
Poland - Culture, Science and Media. Condensation of oxygen and nitrogen. Retrieved on 2008-10-04.
In 1891 Scottish chemist
James Dewar was able to produce enough liquid oxygen to study.
Emsley 2001, p.303 The first commercially viable process for producing liquid oxygen was independently developed in 1895 by German engineer
Carl von Linde and British engineer William Hampson. Both men lowered the temperature of air until it liquefied and then
distilled the component gases by boiling them off one at a time and capturing them. Later, in 1901, oxyacetylene
welding was demonstrated for the first time by burning a mixture of
acetylene and compressed . This method of welding and cutting metal later became common.
In 1923 the American scientist
Robert H. Goddard became the first person to develop a
rocket engine; the engine used
gasoline for fuel and liquid oxygen as the
oxidizer. Goddard successfully flew a small liquid-fueled rocket 56 m at 97 km/h on March 16, 1926 in
Auburn, Massachusetts, USA.
Industrial production
Two major methods are employed to produce 100 million tonnes of extracted from air for industrial uses annually. 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.
The other major method of producing gas involves passing a stream of clean, dry air through one bed of a pair of identical
zeolite molecular sieves, which absorbs the nitrogen and delivers a gas stream that is 90% to 93% . Simultaneously, nitrogen gas is released from the other nitrogen-saturated zeolite bed, by reducing the chamber operating pressure and diverting part of the oxygen gas from the producer bed through it, in the reverse direction of flow. After a set cycle time the operation of the two beds is interchanged, thereby allowing for a continuous supply of gaseous oxygen to be pumped through a pipeline. This is known as
pressure swing adsorption. Oxygen gas is increasingly obtained by these non-
cryogenic technologies (see also the related
vacuum swing adsorption).
Oxygen gas can also be produced through
electrolysis of water into molecular oxygen and hydrogen. DC electricity must be used: if AC is used, the gases in each limb consist of hydrogen and oxygen in the explosive ratio 2:1. Contrary to popular belief, the 2:1 ratio observed in the DC electrolysis of acidified water does not prove that the empirical formula of water is H2O unless certain assumptions are made about the molecular formulae of hydrogen and oxygen themselves.
A similar method is the electrocatalytic 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 submarines, and are still part of standard equipment on commercial airliners in case of depressurization emergencies. Another air separation technology involves forcing air to dissolve through
ceramic membranes based on
zirconium dioxide by either high pressure or an electric current, to produce nearly pure gas.
In large quantities, the price of liquid oxygen in 2001 was approximately $0.21/kg. Since the primary cost of production is the energy cost of liquefying the air, the production cost will change as energy cost varies.
For reasons of economy, oxygen is often transported in bulk as a liquid in specially insulated tankers, since one
litre of liquefied oxygen is equivalent to 840 liters of gaseous oxygen at atmospheric pressure and 20 °
C. Such tankers are used to refill bulk liquid oxygen storage containers, which stand outside hospitals and other institutions with a need for large volumes of pure oxygen gas. Liquid oxygen is passed through
heat exchangers, which convert the cryogenic liquid into gas before it enters the building. Oxygen is also stored and shipped in smaller
cylinders containing the compressed gas; a form that is useful in certain portable medical applications and
oxy-fuel welding and cutting.
Applications
Medical
Uptake of from the air is the essential purpose of
respiration, so oxygen supplementation is used in
medicine.
Oxygen therapy is used to treat
emphysema,
pneumonia, some heart disorders, and any
disease that impairs the body's ability to take up and use gaseous oxygen.
Cook & Lauer 1968, p.510 Treatments are flexible enough to be used in hospitals, the patient's home, or increasingly by portable devices.
Oxygen tents were once commonly used in oxygen supplementation, but have since been replaced mostly by the use of
oxygen masks or
nasal cannulas.
Hyperbaric (high-pressure) medicine uses special
oxygen chamber to increase the
partial pressure of around the patient and, when needed, the medical staff.
Carbon monoxide poisoning,
gas gangrene, and
decompression sickness (the 'bends') are sometimes treated using these devices. Increased concentration in the lungs helps to displace
carbon monoxide from the heme group of
hemoglobin. Oxygen gas is poisonous to the
anaerobic bacteria that cause gas gangrene, so increasing its partial pressure helps kill them. Decompression sickness occurs in divers who decompress too quickly after a dive, resulting in bubbles of inert gas, mostly nitrogen and helium, forming in their blood. Increasing the pressure of as soon as possible is part of the treatment.
Oxygen is also used medically for patients who require
mechanical ventilation, often at concentrations above 21% found in ambient air./NEEDS TO BE EXPANDED -->
Life support and recreational use
A notable application of as a low-pressure
breathing gas is in modern
space suits, which surround their occupant's body with pressurized air. These devices use nearly pure oxygen at about one third normal pressure, resulting in a normal blood partial pressure of . This trade-off of higher oxygen concentration for lower pressure is needed to maintain flexible spacesuits.
Scuba diver and
submariners also rely on artificially delivered , but most often use normal pressure, and/or mixtures of oxygen and air. Pure or nearly pure use in diving at higher-than-sea-level pressures is usually limited to rebreather, decompression, or emergency treatment use at relatively shallow depths (~6 meters depth, or less). Deeper diving requires significant dilution of with other gases, such as nitrogen or helium, to help prevent
oxygen toxicity.
People who climb mountains or fly in non-pressurized
fixed-wing aircraft sometimes have supplemental supplies.The reason is that increasing the proportion of oxygen in the breathing gas at low pressure acts to augment the inspired partial pressure nearer to that found at sea-level. Passengers traveling in (pressurized) commercial airplanes have an emergency supply of automatically supplied to them in case of cabin depressurization. Sudden cabin pressure loss activates
chemical oxygen generators above each seat, causing
oxygen masks to drop and forcing iron filings into the
sodium chlorate inside the canister. A steady stream of oxygen gas is produced by the
exothermic reaction. However, even this may pose a danger if inappropriately triggered: a
ValuJet airplane crashed after use-date-expired canisters, which were being shipped in the cargo hold, activated and caused fire. The canisters were mis-labeled as empty, and carried against
dangerous goods regulations.)
Oxygen, as a supposed mild
euphoric, has a history of recreational use in
oxygen bars and in
sports. Oxygen bars are establishments, found in
Japan,
California, and
Las Vegas, Nevada since the late 1990s that offer higher than normal exposure for a fee. Professional athletes, especially in
American football, also sometimes go off field between plays to wear oxygen masks in order to get a supposed "boost" in performance. However, the reality of a pharmacological effect is doubtful; a
placebo or psychological boost being the most plausible explanation. Available studies support a performance boost from enriched mixtures only if they are breathed during actual aerobic exercise. Other recreational uses include
pyrotechnic applications, such as
George Goble's five-second ignition of
barbecue grills.
Industrial
Smelting of
iron ore into
steel consumes 55% of commercially produced oxygen. In this process, is injected through a high-pressure lance into molten iron, which removes
sulfur impurities and excess
carbon as the respective oxides, and . The reactions are
exothermic, so the temperature increases to 1,700 °
C.
Another 25% of commercially produced oxygen is used by the chemical industry.
Ethylene is reacted with 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).
Most of the remaining 20% of commercially produced oxygen is used in medical applications,
metal cutting and welding, as an oxidizer in
rocket fuel, and in
water treatment. Oxygen is used in
oxyacetylene welding burning
acetylene with to produce a very hot flame. In this process, metal up to 60
cm thick is first heated with a small oxy-acetylene flame and then quickly cut by a large stream of .
Cook & Lauer 1968, p.508
Rocket propulsion requires a fuel and an oxidizer. Larger
rockets use liquid oxygen as their oxidizer, which is mixed and ignited with the fuel for propulsion.
Scientific
Paleoclimatologists measure the ratio of
oxygen-18 and oxygen-16 in the
shells and
skeletons of marine organisms to determine what the climate was like millions of years ago (see
oxygen isotope ratio cycle).
Seawater molecules that contain the lighter
isotope, oxygen-16, evaporate at a slightly faster rate than water molecules containing the 12% heavier oxygen-18; this disparity increases at lower temperatures.
Emsley 2001, p.304 During periods of lower global temperatures, snow and rain from that evaporated water tends to be higher in oxygen-16, and the seawater left behind tends to be higher in oxygen-18. Marine organisms then incorporate more oxygen-18 into their skeletons and shells than they would in a warmer climate. Paleoclimatologists also directly measure this ratio in the water molecules of
ice core samples that are up to several hundreds of thousands of years old.
Planetary geologists have measured different abundances of oxygen isotopes in samples from the
Earth, the
Moon,
Mars, and
meteorites, but were long unable to obtain reference values for the isotope ratios in the
Sun, believed to be the same as those of the
primordial solar nebula. However, analysis of a
silicon wafer exposed to the
solar wind in space and returned by the crashed
Genesis spacecraft has shown that the Sun has a higher proportion of oxygen-16 than does the Earth. The measurement implies that an unknown process depleted oxygen-16 from the Sun's
disk of protoplanetary material prior to the coalescence of dust grains that formed the Earth.
Oxygen presents two spectrophotometric
absorption bands peaking at the wavelengths 687 and 760
nm. Some
remote sensing scientists have proposed using the measurement of the radiance coming from vegetation canopies in those bands to characterize plant health status from a
satellite platform. This approach exploits the fact that in those bands it is possible to discriminate the vegetation's
reflectance from its
fluorescence, which is much weaker. The measurement is technically difficult owing to the low
signal-to-noise ratio and the physical structure of vegetation; but it has been proposed as a possible method of monitoring the
carbon cycle from satellites on a global scale.
Compounds
The
oxidation state of oxygen is −2 in almost all known compounds of oxygen. The oxidation state −1 is found in a few compounds such as
peroxides.,p. 28 Compounds containing oxygen in other oxidation states are very uncommon: −1/2 (
superoxides), −1/3 (
ozonides), 0 (
elemental,
hypofluorous acid), +1/2 (
dioxygenyl), +1 (
dioxygen difluoride), and +2 (
oxygen difluoride).
Oxides and other inorganic compounds
Water () is the oxide of
hydrogen and the most familiar oxygen compound. Hydrogen atoms are
covalently bonded to oxygen in a water molecule but also have an additional attraction (about 23.3 kJ·mol−1 per hydrogen atom) to an adjacent oxygen atom in a separate molecule. These
hydrogen bonds between water molecules hold them approximately 15% closer than what would be expected in a simple liquid with just
Van der Waals forces. Also, since oxygen has a higher electronegativity than hydrogen, the charge difference makes it a
polar molecule. The interactions between the different
dipoles of each molecule cause a net attraction force.
Due to its
electronegativity, oxygen forms
chemical bonds with almost all other elements at elevated temperatures to give corresponding
oxides. However, some elements readily form oxides at
standard conditions for temperature and pressure; the
rusting of
iron is an example. The surface of metals like
aluminium and
titanium are oxidized in the presence of air and become coated with a thin film of oxide that
passivates the metal and slows further
corrosion. Some of the
transition metal oxides are found in nature as
non-stoichiometric compounds, with a slightly less metal than the
chemical formula would show. For example, the natural occurring
FeO (
wüstite) is actually written as , where x is usually around 0.05.
Oxygen as a compound is present in the atmosphere in trace quantities in the form of
carbon dioxide (). The
earth's crustal rock is composed in large part of oxides of
silicon (
silica , found in
granite and
sand),
aluminium (
aluminium oxide , in
bauxite and
corundum), iron (
iron oxide , in
hematite and
rust) and other
metals.
The rest of the Earth's crust is also made of oxygen compounds, in particular
calcium carbonate (in
limestone) and
silicates (in
feldspars). Water-
soluble silicates in the form of , , and are used as
detergents and
adhesives.
Cook & Lauer 1968, p.507
Oxygen also acts as a ligand for transition metals, forming metal– bonds with the
iridium atom in
Vaska's complex, with the
platinum in
platinum hexafluoride,
Cook & Lauer 1968, p.505 and with the iron center of the
heme group of
hemoglobin.
Organic compounds and biomolecules
Among the most important classes of organic compounds that contain oxygen are (where "R" is an organic group):
alcohols (R-OH);
ethers (R-O-R);
ketones (R-CO-R);
aldehydes (R-CO-H);
carboxylic acids (R-COOH);
esters (R-COO-R);
acid anhydrides (R-CO-O-CO-R); and
amides (). There are many important organic
solvents that contain oxygen, including:
acetone,
methanol,
ethanol,
isopropanol,
furan,
THF,
diethyl ether,
dioxane,
ethyl acetate,
DMF,
DMSO,
acetic acid, and
formic acid.
Acetone () and
phenol () are used as feeder materials in the synthesis of many different substances. Other important organic compounds that contain oxygen are:
glycerol,
formaldehyde,
glutaraldehyde,
citric acid,
acetic anhydride, and
acetamide.
Epoxides are
ethers in which the oxygen atom is part of a ring of three atoms.
Oxygen reacts spontaneously with many
organic compounds at or below room temperature in a process called
autoxidation.
Cook & Lauer 1968, p.506 Most of the
organic compounds that contain oxygen are not made by direct action of . Organic compounds important in industry and commerce that are made by direct oxidation of a precursor include
ethylene oxide and
peracetic acid.
The element is found in almost all
biomolecules that are important to (or generated by) life. Only a few common complex biomolecules, such as
squalene and the
carotenes, contain no oxygen. Of the organic compounds with biological relevance,
carbohydrates contain the largest proportion by mass of oxygen. All
fats,
fatty acids,
amino acids, and
proteins contain oxygen (due to the presence of
carbonyl groups in these acids and their
ester residues). Oxygen also occurs in
phosphate () groups in the biologically important energy-carrying molecules
ATP and
ADP, in the backbone and the
purines (except
adenine) and
pyrimidines of
RNA and
DNA, and in bones as
calcium phosphate and
hydroxylapatite.
Safety and Precautions
Toxicity
Eyes - visual field loss, near)sightedness, cataract formation, bleeding, fibrosis; Head - seizures; Muscles - twitching; Respiratory system - jerky breathing, irritation, coughing, pain, shortness of breath, tracheobronchitis, acute respiratory distress syndrome.Oxygen gas () can be
toxic at elevated
partial pressures, leading to
convulsions and other health problems.Since 's partial pressure is the fraction of times the total pressure, elevated partial pressures can occur either from high fraction in breathing gas or from high breathing gas pressure, or a combination of both.
Cook & Lauer 1968, p.511 Oxygen toxicity usually begins to occur at partial pressures more than 50 kilo
pascal (kPa), or 2.5 times the normal sea-level partial pressure of about 21 kPa. Therefore, air supplied through
oxygen masks in medical applications is typically composed of 30% by volume (about 30 kPa at standard pressure). At one time,
premature babies were placed in incubators containing -rich air, but this practice was discontinued after some babies were blinded by it.
Breathing pure in space applications, such as in some modern
space suits, or in early spacecraft such as
Apollo, causes no damage due to the low total pressures used. In the case of spacesuits, the partial pressure in the breathing gas is, in general, about 30 kPa (1.4 times normal), and the resulting partial pressure in the astronaut's arterial blood is only marginally more than normal sea-level partial pressure (see
arterial blood gas).
Oxygen toxicity to the lungs and
central nervous system can also occur in deep
scuba diving and
surface supplied diving. Prolonged breathing of an air mixture with an partial pressure more than 60 kPa can eventually lead to permanent
pulmonary fibrosis. Exposure to a partial pressures greater than 160 kPa may lead to convulsions (normally fatal for divers). Acute oxygen toxicity can occur by breathing an air mixture with 21% at 66 m or more of depth while the same thing can occur by breathing 100% at only 6 m.
Combustion and other hazards
Highly concentrated sources of oxygen promote rapid combustion.
Fire and
explosion hazards exist when concentrated oxidants and
fuels are brought into close proximity; however, an ignition event, such as heat or a spark, is needed to trigger combustion. Oxygen itself is not the fuel, but the oxidant. Combustion hazards also apply to compounds of oxygen with a high oxidative potential, such as
peroxides,
chlorates,
nitrates,
perchlorates, and
dichromates because they can donate oxygen to a fire.
Concentrated will allow combustion to proceed rapidly and energetically.
Steel pipes and storage vessels used to store and transmit both gaseous and
liquid oxygen will act as a fuel; and therefore the design and manufacture of systems requires special training to ensure that ignition sources are minimized. The fire that killed the
Apollo 1 crew on a test launch pad spread so rapidly because the capsule was pressurized with pure but at slightly more than atmospheric pressure, instead of the normal pressure that would be used in a mission.No single ignition source of the fire was conclusively identified, although some evidence points to arc from an electrical spark). (Report of Apollo 204 Review Board NASA Historical Reference Collection, NASA History Office, NASA HQ, Washington, DC)
Liquid oxygen spills, if allowed to soak into organic matter, such as
wood,
petrochemicals, and
asphalt can cause these materials to
detonate unpredictably on subsequent mechanical impact. As with other
cryogenic liquids, on contact with the human body it can cause burns to the skin and the eyes.
See also
Notes and citations
References
Further reading
External links
Chemical elementsNonmetalsChalcogensBreathing gasesOxygenBiology and pharmacology of chemical elements
This entry is from Wikipedia,the leading user-contributed encyclopedia. It may not have been reviewed by professional editors (See full disclaimer)
