Oxygen 1

Oxygen

Oxygen

Appearance

colourless gas, pale blue liquid

Spectral lines of Oxygen

General properties

Name, symbol, number oxygen, O, 8

Pronunciation English pronunciation: /ˈɒksɨdʒɨn/, OK-si-jin

Element category nonmetal, chalcogens

Group, period, block 16, 2, p

Standard atomic weight 15.9994(3) g·mol−1

Electron configuration 1s2 2s2 2p4

Electrons per shell 2, 6 (Image)

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 (O2) 0.444 kJ·mol−1

Heat of vaporization (O2) 6.82 kJ·mol−1

Specific heat capacity (25 °C) (O2)29.378 J·mol−1·K−1

Vapor pressure

P/Pa 1 10 100 1 k 10 k 100 k

atT/K

61 73 90

Atomic properties

Oxidation states 2, 1, −1, −2(neutral oxide)

Electronegativity 3.44 (Pauling scale)

Oxygen 2

Ionization energies(more)

1st: 1313.9 kJ·mol−1

2nd: 3388.3 kJ·mol−1

3rd: 5300.5 kJ·mol−1

Covalent radius 66±2 pm

Van der Waals radius 152 pm

Miscellanea

Crystal structure cubic

Magnetic ordering paramagnetic

Thermal conductivity (300 K) 26.58x10-3  W·m−1·K−1

Speed of sound (gas, 27 °C) 330 m/s

CAS registry number 7782-44-7

Most stable isotopes

Main article: Isotopes of oxygen

iso NA half-life DM DE(MeV)

DP

16O 99.76% 16O is stable with 8 neutron

17O 0.039% 17O is stable with 9 neutron

18O 0.201% 18O is stable with 10 neutron

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 memberof the chalcogen group on the periodic table, and is a highly reactive nonmetallic period 2 element that readily formscompounds (notably oxides) with almost all other elements. At standard temperature and pressure two atoms of theelement bind to form dioxygen, a colorless, odorless, tasteless diatomic gas with the formula O2. Oxygen is the thirdmost abundant element in the universe by mass after hydrogen and helium[1] and the most abundant element by massin 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, containoxygen, as do the major inorganic compounds that comprise animal shells, teeth, and bone. Oxygen in the form of O2is produced from water by cyanobacteria, algae and plants during photosynthesis and is used in cellular respirationfor all complex life. Oxygen is toxic to obligately anaerobic organisms, which were the dominant form of early lifeon 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 pollutantnear the surface where it is a by-product of smog. At even higher low earth orbit altitudes atomic oxygen is asignificant presence and a cause of erosion for spacecraft.[5]

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,[6] 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

Oxygen 3

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

StructureAt standard temperature and pressure, oxygen is a colorless, odorless gas with the molecular formula O2, in whichthe two oxygen atoms are chemically bonded to each other with a spin triplet electron configuration. This bond has abond order of two, and is often simplified in description as a double bond[7] or as a combination of one two-electronbond and two three-electron bonds.[8]

Triplet oxygen (not to be confused with ozone, O3) is the ground state of the O2 molecule.[9] The electronconfiguration of the molecule has two unpaired electrons occupying two degenerate molecular orbitals.[10] Theseorbitals are classified as antibonding (weakening the bond order from three to two), so the diatomic oxygen bond isweaker than the diatomic nitrogen triple bond in which all bonding molecular orbitals are filled, but someantibonding orbitals are not.[9]

In normal triplet form, O2 molecules are paramagnetic—they form a magnet in the presence of a magneticfield—because of the spin magnetic moments of the unpaired electrons in the molecule, and the negative exchangeenergy between neighboring O2 molecules.[11] Liquid oxygen is attracted to a magnet to a sufficient extent that, inlaboratory demonstrations, a bridge of liquid oxygen may be supported against its own weight between the poles of apowerful magnet.[12] [13]

Singlet oxygen, a name given to several higher-energy species of molecular O2 in which all the electron spins arepaired, is much more reactive towards common organic molecules. In nature, singlet oxygen is commonly formedfrom water during photosynthesis, using the energy of sunlight.[14] It is also produced in the troposphere by thephotolysis of ozone by light of short wavelength,[15] and by the immune system as a source of active oxygen.[16]

Carotenoids in photosynthetic organisms (and possibly also in animals) play a major role in absorbing energy fromsinglet oxygen and converting it to the unexcited ground state before it can cause harm to tissues.[17]

Allotropes

Ozone is a rare gas on Earth foundmostly in the stratosphere

The common allotrope of elemental oxygen on Earth is called dioxygen, O2. Ithas a bond length of 121 pm and a bond energy of 498 kJ·mol-1.[18] This is theform that is used by complex forms of life, such as animals, in cellularrespiration (see Biological role) and is the form that is a major part of the Earth'satmosphere (see Occurrence). Other aspects of O2 are covered in the remainderof this article.

Trioxygen (O3) is usually known as ozone and is a very reactive allotrope ofoxygen that is damaging to lung tissue.[19] Ozone is produced in the upperatmosphere when O2 combines with atomic oxygen made by the splitting of O2by ultraviolet (UV) radiation.[6] Since ozone absorbs strongly in the UV regionof the spectrum, the ozone layer of the upper atmosphere functions as aprotective radiation shield for the planet.[6] Near the Earth's surface, however, itis a pollutant formed as a by-product of automobile exhaust.[19] The metastablemolecule tetraoxygen (O4) was discovered in 2001,[20] [21] and was assumed toexist in one of the six phases of solid oxygen. It was proven in 2006 that this

Oxygen 4

phase, created by pressurizing O2 to 20 GPa, is in fact a rhombohedral O8 cluster.[22] This cluster has the potential tobe a much more powerful oxidizer than either O2 or O3 and may therefore be used in rocket fuel.[21] [20] A metallicphase was discovered in 1990 when solid oxygen is subjected to a pressure of above 96 GPa[23] and it was shown in1998 that at very low temperatures, this phase becomes superconducting.[24]

Physical propertiesOxygen is more soluble in water than nitrogen; water contains approximately 1 molecule of O2 for every 2 moleculesof N2, compared to an atmospheric ratio of approximately 1:4. The solubility of oxygen in water istemperature-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 standard atmosphere (101.3 kPa) of air, freshwater contains about 6.04 milliliters (mL) of oxygen perliter, whereas seawater contains about 4.95 mL per liter.[27] At 5 °C the solubility increases to 9.0 mL (50% morethan 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] Bothliquid and solid O2 are clear substances with a light sky-blue color caused by absorption in the red (in contrast withthe blue color of the sky, which is due to Rayleigh scattering of blue light). High-purity liquid O2 is usually obtainedby 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 combustiblematerials.[30]

Isotopes and stellar origin

Late in a massive star's life, 16O concentrates inthe O-shell, 17O in the H-shell and 18O in the

He-shell

Naturally occurring oxygen is composed of three stable isotopes, 16O,17O, and 18O, with 16O being the most abundant (99.762% naturalabundance).[31]

Most 16O is synthesized at the end of the helium fusion process in starsbut some is made in the neon burning process.[32] 17O is primarilymade by the burning of hydrogen into helium during the CNO cycle,making it a common isotope in the hydrogen burning zones of stars.[32]

Most 18O is produced when 14N (made abundant from CNO burning)captures a 4He nucleus, making 18O common in the helium-rich zonesof stars.[32]

Fourteen radioisotopes have been characterized, the most stable being15O with a half-life of 122.24 seconds (s) and 14O with a half-life of70.606 s.[31] All of the remaining radioactive isotopes have half-livesthat are less than 27 s and the majority of these have half-lives that areless than 83 milliseconds.[31] The most common decay mode of the isotopes lighter than 16O is β+ decay[33] [34] [35]

to yield nitrogen, and the most common mode for the isotopes heavier than 18O is beta decay to yield fluorine.[31]

OccurrenceOxygen 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.[1] About 0.9% of the Sun's mass is oxygen.[3] Oxygen constitutes 49.2% of the Earth's crust by mass[2] and is the major component of the world's oceans (88.8% by mass).[3] 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).[3] [36] [37] 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% O2 by volume) and Venus have far lower concentrations. However, the O2 surrounding these other planets is produced solely by

Oxygen 5

ultraviolet radiation impacting oxygen-containing molecules such as carbon dioxide.

Cold water holds more dissolved O2.

The unusually high concentration of oxygen gas on Earth is the resultof the oxygen cycle. This biogeochemical cycle describes themovement of oxygen within and between its three main reservoirs onEarth: the atmosphere, the biosphere, and the lithosphere. The maindriving factor of the oxygen cycle is photosynthesis, which isresponsible for modern Earth's atmosphere. Photosynthesis releasesoxygen into the atmosphere, while respiration and decay remove itfrom the atmosphere. In the present equilibrium, production andconsumption occur at the same rate of roughly 1/2000th of the entireatmospheric oxygen per year.

Free oxygen also occurs in solution in the world's water bodies. The increased solubility of O2 at lower temperatures(see Physical properties) has important implications for ocean life, as polar oceans support a much higher density oflife due to their higher oxygen content.[38] Polluted water may have reduced amounts of O2 in it, depleted bydecaying algae and other biomaterials (see eutrophication). Scientists assess this aspect of water quality bymeasuring the water's biochemical oxygen demand, or the amount of O2 needed to restore it to a normalconcentration.[39]

Biological role

Photosynthesis and respiration

Photosynthesis splits water to liberate O2 andfixes CO2 into sugar

In nature, free oxygen is produced by the light-driven splitting of waterduring oxygenic photosynthesis. Green algae and cyanobacteria inmarine environments provide about 70% of the free oxygen producedon earth and the rest is produced by terrestrial plants.[40]

A simplified overall formula for photosynthesis is:[41]

6 CO2 + 6 H2O + photons → C6H12O6 + 6 O2 (or simply

carbon dioxide + water + sunlight → glucose + dioxygen)

Photolytic oxygen evolution occurs in the thylakoid membranes ofphotosynthetic organisms and requires the energy of four photons.[42]

Many steps are involved, but the result is the formation of a protongradient across the thylakoid membrane, which is used to synthesizeATP via photophosphorylation.[43] The O2 remaining after oxidation ofthe water molecule is released into the atmosphere.[44]

Molecular dioxygen, O2, is essential for cellular respiration in allaerobic organisms. Oxygen is used in mitochondria to help generateadenosine triphosphate (ATP) during oxidative phosphorylation. Thereaction for aerobic respiration is essentially the reverse of photosynthesis and is simplified as:

C6H12O6 + 6 O2 → 6 CO2 + 6 H2O + 2880 kJ·mol-1

In vertebrates, O2 is diffused through membranes in the lungs and into red blood cells. Hemoglobin binds O2,changing its color from bluish red to bright red.[19] [45] Other animals use hemocyanin (molluscs and somearthropods) or hemerythrin (spiders and lobsters).[36] A liter of blood can dissolve 200 cm3 of O2.[36]

Reactive oxygen species, such as superoxide ion (O−2) and hydrogen peroxide (H2O2), are dangerous by-products of

oxygen use in organisms.[36] Parts of the immune system of higher organisms, however, create peroxide, superoxide,

Oxygen 6

and singlet oxygen to destroy invading microbes. Reactive oxygen species also play an important role in thehypersensitive response of plants against pathogen attack.[43]

An adult human in rest inhales 1.8 to 2.4 grams of oxygen per minute.[46] This amounts to more than 6 billion tonnesof oxygen inhaled by humanity per year.[47]

Build-up in the atmosphere

O2 build-up in Earth's atmosphere: 1) no O2 produced; 2) O2 produced, butabsorbed in oceans & seabed rock; 3) O2 starts to gas out of the oceans, but is

absorbed by land surfaces and formation of ozone layer; 4–5) O2 sinks filled andthe gas accumulates

Free oxygen gas was almost nonexistent inEarth's atmosphere before photosyntheticarchaea and bacteria evolved. Free oxygenfirst appeared in significant quantitiesduring the Paleoproterozoic era (between2.5 and 1.6 billion years ago). At first, theoxygen combined with dissolved iron in theoceans to form banded iron formations. Freeoxygen started to gas out of the oceans2.7 billion years ago, reaching 10% of itspresent level around 1.7 billion yearsago.[48]

The presence of large amounts of dissolvedand free oxygen in the oceans and atmosphere may have driven most of the anaerobic organisms then living toextinction during the oxygen catastrophe about 2.4 billion years ago. However, cellular respiration using O2 enablesaerobic organisms to produce much more ATP than anaerobic organisms, helping the former to dominate Earth'sbiosphere.[49] Photosynthesis and cellular respiration of O2 allowed for the evolution of eukaryotic cells andultimately complex multicellular organisms such as plants and animals.

Since the beginning of the Cambrian era 540 million years ago, O2 levels have fluctuated between 15% and 30% byvolume.[50] Towards the end of the Carboniferous era (about 300 million years ago) atmospheric O2 levels reached amaximum of 35% by volume,[50] which may have contributed to the large size of insects and amphibians at thistime.[51] Human activities, including the burning of 7 billion tonnes of fossil fuels each year have had very littleeffect on the amount of free oxygen in the atmosphere.[11] At the current rate of photosynthesis it would take about2,000 years to regenerate the entire O2 in the present atmosphere.[52]

Oxygen 7

History

Early experiments

Philo's experiment inspired laterinvestigators

One of the first known experiments on the relationship between combustion andair was conducted by the second century BCE Greek writer on mechanics, Philoof Byzantium. In his work Pneumatica, Philo observed that inverting a vesselover a burning candle and surrounding the vessel's neck with water resulted insome water rising into the neck.[53] Philo incorrectly surmised that parts of the airin the vessel were converted into the classical element fire and thus were able toescape through pores in the glass. Many centuries later Leonardo da Vinci builton Philo's work by observing that a portion of air is consumed during combustionand respiration.[54]

In the late 17th century, Robert Boyle proved that air is necessary forcombustion. English chemist John Mayow refined this work by showing that firerequires only a part of air that he called spiritus nitroaereus or justnitroaereus.[55] In one experiment he found that placing either a mouse or a litcandle in a closed container over water caused the water to rise and replaceone-fourteenth of the air's volume before extinguishing the subjects.[56] From thishe surmised that nitroaereus is consumed in both respiration and combustion.

Mayow observed that antimony increased in weight when heated, and inferredthat the nitroaereus must have combined with it.[55] He also thought that thelungs separate nitroaereus from air and pass it into the blood and that animal heatand muscle movement result from the reaction of nitroaereus with certain substances in the body.[55] Accounts ofthese and other experiments and ideas were published in 1668 in his work Tractatus duo in the tract "Derespiratione".[56]

Phlogiston theory

Stahl helped develop and popularizethe phlogiston theory.

Robert Hooke, Ole Borch, Mikhail Lomonosov, and Pierre Bayen all producedoxygen in experiments in the 17th and the 18th century but none of themrecognized it as an element.[25] This may have been in part due to the prevalenceof the philosophy of combustion and corrosion called the phlogiston theory,which was then the favored explanation of those processes.

Established in 1667 by the German alchemist J. J. Becher, and modified by thechemist Georg Ernst Stahl by 1731,[57] phlogiston theory stated that allcombustible materials were made of two parts. One part, called phlogiston, wasgiven off when the substance containing it was burned, while thedephlogisticated part was thought to be its true form, or calx.[54]

Highly combustible materials that leave little residue, such as wood or coal, werethought to be made mostly of phlogiston; whereas non-combustible substancesthat corrode, such as iron, contained very little. Air did not play a role inphlogiston theory, nor were any initial quantitative experiments conducted to test

the idea; instead, it was based on observations of what happens when something burns, that most common objectsappear to become lighter and seem to lose something in the process.[54] The fact that a substance like wood actually

Oxygen 8

gains overall weight in burning was hidden by the buoyancy of the gaseous combustion products. Indeed one of thefirst clues that the phlogiston theory was incorrect was that metals, too, gain weight in rusting (when they weresupposedly losing phlogiston).

Carl Wilhelm Scheele beat Priestleyto the discovery but published

afterwards.

Discovery

Oxygen was first discovered by Swedish pharmacist Carl Wilhelm Scheele. Hehad produced oxygen gas by heating mercuric oxide and various nitrates byabout 1772.[3] [54] Scheele called the gas 'fire air' because it was the only knownsupporter of combustion, and wrote an account of this discovery in a manuscripthe titled Treatise on Air and Fire, which he sent to his publisher in 1775.However, that document was not published until 1777.[58]

Joseph Priestley is usually givenpriority in the discovery

In the meantime, on August 1, 1774, an experiment conducted by the Britishclergyman Joseph Priestley focused sunlight on mercuric oxide (HgO) inside aglass tube, which liberated a gas he named 'dephlogisticated air'.[3] He noted thatcandles burned brighter in the gas and that a mouse was more active and livedlonger while breathing it. After breathing the gas himself, he wrote: "The feelingof it to my lungs was not sensibly different from that of common air, but Ifancied that my breast felt peculiarly light and easy for some timeafterwards."[25] Priestley published his findings in 1775 in a paper titled "AnAccount of Further Discoveries in Air" which was included in the second volumeof his book titled Experiments and Observations on Different Kinds of Air.[54]

[59] Because he published his findings first, Priestley is usually given priority inthe discovery.

The noted French chemist Antoine Laurent Lavoisier later claimed to havediscovered the new substance independently. However, Priestley visited Lavoisier in October 1774 and told himabout 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 acknowledgedreceiving it (a copy of the letter was found in Scheele's belongings after his death).[58]

Oxygen 9

Lavoisier's contributionWhat Lavoisier did indisputably do (although this was disputed at the time) was to conduct the first adequatequantitative experiments on oxidation and give the first correct explanation of how combustion works.[3] He usedthese and similar experiments, all started in 1774, to discredit the phlogiston theory and to prove that the substancediscovered by Priestley and Scheele was a chemical element.

Antoine Lavoisier discredited thePhlogiston theory

In one experiment, Lavoisier observed that there was no overall increase inweight when tin and air were heated in a closed container.[3] He noted that airrushed in when he opened the container, which indicated that part of the trappedair had been consumed. He also noted that the tin had increased in weight andthat increase was the same as the weight of the air that rushed back in. This andother experiments on combustion were documented in his book Sur lacombustion en général, which was published in 1777.[3] In that work, he provedthat air is a mixture of two gases; 'vital air', which is essential to combustion andrespiration, and azote (Gk. ἄζωτον "lifeless"), which did not support either.Azote later became nitrogen in English, although it has kept the name in Frenchand several other European languages.[3]

Lavoisier renamed 'vital air' to oxygène in 1777 from the Greek roots ὀξύς (oxys)(acid, literally "sharp," from the taste of acids) and -γενής (-genēs) (producer, literally begetter), because hemistakenly believed that oxygen was a constituent of all acids.[6] Chemists eventually determined that Lavoisier waswrong in this regard, but by that time it was too late, the name had taken. Actually, the gas that could moreappropriately have been given the description, "acid producer," is hydrogen.

Oxygen entered the English language despite opposition by English scientists and the fact that the EnglishmanPriestley 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.[58]

Later history

Robert H. Goddard and a liquid oxygen-gasolinerocket

John Dalton's original atomic hypothesis assumed that all elementswere monoatomic and that the atoms in compounds would normallyhave the simplest atomic ratios with respect to one another. Forexample, Dalton assumed that water's formula was HO, giving theatomic mass of oxygen as 8 times that of hydrogen, instead of themodern value of about 16.[60] In 1805, Joseph Louis Gay-Lussac andAlexander von Humboldt showed that water is formed of two volumesof hydrogen and one volume of oxygen; and by 1811 AmedeoAvogadro had arrived at the correct interpretation of water'scomposition, based on what is now called Avogadro's law and theassumption of diatomic elemental molecules.[61] [62]

By the late 19th century scientists realized that air could be liquefied,and its components isolated, by compressing and cooling it. Using acascade method, Swiss chemist and physicist Raoul Pierre Pictetevaporated 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.[63] Just two days later, French physicist Louis Paul Cailletet announced his own method of liquefying molecular oxygen.[63] Only a few drops of the liquid were produced in either case so no meaningful analysis could

Oxygen 10

be conducted. Oxygen was liquified in stable state for the first time on March 29, 1877 by Polish scientists fromJagiellonian University, Zygmunt Wróblewski and Karol Olszewski.[64]

In 1891 Scottish chemist James Dewar was able to produce enough liquid oxygen to study.[11] The firstcommercially viable process for producing liquid oxygen was independently developed in 1895 by German engineerCarl von Linde and British engineer William Hampson. Both men lowered the temperature of air until it liquefiedand then distilled the component gases by boiling them off one at a time and capturing them.[65] Later, in 1901,oxyacetylene welding was demonstrated for the first time by burning a mixture of acetylene and compressed O2.This method of welding and cutting metal later became common.[65]

In 1923 the American scientist Robert H. Goddard became the first person to develop a rocket engine; the engineused gasoline for fuel and liquid oxygen as the oxidizer. Goddard successfully flew a small liquid-fueled rocket 56 mat 97 km/h on March 16, 1926 in Auburn, Massachusetts, USA.[65] [66]

Industrial productionTwo major methods are employed to produce 100 million tonnes of O2 extracted from air for industrial usesannually.[58] The most common method is to fractionally distill liquefied air into its various components, withnitrogen N2 distilling as a vapor while oxygen O2 is left as a liquid.[58]

Hoffman electrolysis apparatus used inelectrolysis of water.

The other major method of producing O2 gas involves passing a streamof clean, dry air through one bed of a pair of identical zeolite molecularsieves, which absorbs the nitrogen and delivers a gas stream that is90% to 93% O2.[58] Simultaneously, nitrogen gas is released from theother nitrogen-saturated zeolite bed, by reducing the chamber operatingpressure and diverting part of the oxygen gas from the producer bedthrough it, in the reverse direction of flow. After a set cycle time theoperation of the two beds is interchanged, thereby allowing for acontinuous supply of gaseous oxygen to be pumped through a pipeline.This is known as pressure swing adsorption. Oxygen gas isincreasingly obtained by these non-cryogenic technologies (see alsothe related vacuum swing adsorption).[67]

Oxygen gas can also be produced through electrolysis of water intomolecular oxygen and hydrogen. DC electricity must be used: if AC isused, the gases in each limb consist of hydrogen and oxygen in theexplosive ratio 2:1. Contrary to popular belief, the 2:1 ratio observed inthe DC electrolysis of acidified water does not prove that the empiricalformula of water is H2O unless certain assumptions are made about themolecular formulae of hydrogen and oxygen themselves.

A similar method is the electrocatalytic O2 evolution from oxides and oxoacids. Chemical catalysts can be used aswell, such as in chemical oxygen generators or oxygen candles that are used as part of the life-support equipment onsubmarines, 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 zirconiumdioxide by either high pressure or an electric current, to produce nearly pure O2 gas.[39]

In large quantities, the price of liquid oxygen in 2001 was approximately $0.21/kg.[68] Since the primary cost ofproduction 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.[58] Such tankers are used to refill bulk liquid oxygen storage containers, which stand outside hospitals and other institutions with a

Oxygen 11

need for large volumes of pure oxygen gas. Liquid oxygen is passed through heat exchangers, which convert thecryogenic liquid into gas before it enters the building. Oxygen is also stored and shipped in smaller cylinderscontaining the compressed gas; a form that is useful in certain portable medical applications and oxy-fuel weldingand cutting.[58]

Applications

Medical

An oxygen concentrator in anemphysema patient's house

Uptake of O2 from the air is the essential purpose of respiration, so oxygensupplementation is used in medicine. Oxygen therapy is used to treatemphysema, pneumonia, some heart disorders, and any disease that impairs thebody's ability to take up and use gaseous oxygen.[69] Treatments are flexibleenough to be used in hospitals, the patient's home, or increasingly by portabledevices. Oxygen tents were once commonly used in oxygen supplementation, buthave since been replaced mostly by the use of oxygen masks or nasalcannulas.[70]

Hyperbaric (high-pressure) medicine uses special oxygen chambers to increasethe partial pressure of O2 around the patient and, when needed, the medicalstaff.[71] Carbon monoxide poisoning, gas gangrene, and decompression sickness(the 'bends') are sometimes treated using these devices.[72] Increased O2concentration in the lungs helps to displace carbon monoxide from the hemegroup of hemoglobin.[73] [74] Oxygen gas is poisonous to the anaerobic bacteriathat cause gas gangrene, so increasing its partial pressure helps kill them.[75] [76] Decompression sickness occurs indivers 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 O2 as soon as possible is part of the treatment.[69] [77] [78]

Oxygen is also used medically for patients who require mechanical ventilation, often at concentrations above 21%found in ambient air.

Life support and recreational use

Low pressure pure O2 is used in space suits

A notable application of O2 as a low-pressure breathing gas is inmodern space suits, which surround their occupant's body withpressurized air. These devices use nearly pure oxygen at about onethird normal pressure, resulting in a normal blood partial pressure ofO2.[79] [80] This trade-off of higher oxygen concentration for lowerpressure is needed to maintain flexible spacesuits.

Scuba divers and submariners also rely on artificially delivered O2, butmost often use normal pressure, and/or mixtures of oxygen and air.Pure or nearly pure O2 use in diving at higher-than-sea-level pressuresis usually limited to rebreather, decompression, or emergencytreatment use at relatively shallow depths (~6 meters depth, or less).[81]

[82] Deeper diving requires significant dilution of O2 with other gases,such as nitrogen or helium, to help prevent oxygen toxicity.[81]

People who climb mountains or fly in non-pressurized fixed-wing aircraft sometimes have supplemental O2 supplies.[83] Passengers traveling in (pressurized) commercial airplanes have an emergency supply of O2

Oxygen 12

automatically supplied to them in case of cabin depressurization. Sudden cabin pressure loss activates chemicaloxygen generators above each seat, causing oxygen masks to drop and forcing iron filings into the sodium chlorateinside the canister.[39] A steady stream of oxygen gas is produced by the exothermic reaction. However, even thismay pose a danger if inappropriately triggered: a ValuJet airplane crashed after use-date-expired O2 canisters, whichwere being shipped in the cargo hold, activated and caused fire. The canisters were mis-labeled as empty, and carriedagainst dangerous goods regulations.[84]

Oxygen, as a supposed mild euphoric, has a history of recreational use in oxygen bars and in sports. Oxygen bars areestablishments, found in Japan, California, and Las Vegas, Nevada since the late 1990s that offer higher than normalO2 exposure for a fee.[85] Professional athletes, especially in American football, also sometimes go off field betweenplays to wear oxygen masks in order to get a supposed "boost" in performance. However, the reality of apharmacological effect is doubtful; a placebo or psychological boost being the most plausible explanation.[85]

Available studies support a performance boost from enriched O2 mixtures only if they are breathed during actualaerobic exercise.[86] Other recreational uses include pyrotechnic applications, such as George Goble's five-secondignition of barbecue grills.[87]

Industrial

Most commercially produced O2 is used to smeltiron into steel.

Smelting of iron ore into steel consumes 55% of commerciallyproduced oxygen.[39] In this process, O2 is injected through ahigh-pressure lance into molten iron, which removes sulfur impuritiesand excess carbon as the respective oxides, SO2 and CO2. Thereactions are exothermic, so the temperature increases to 1,700 °C.[39]

Another 25% of commercially produced oxygen is used by thechemical industry.[39] Ethylene is reacted with O2 to create ethyleneoxide, which, in turn, is converted into ethylene glycol; the primaryfeeder material used to manufacture a host of products, includingantifreeze and polyester polymers (the precursors of many plastics andfabrics).[39]

Most of the remaining 20% of commercially produced oxygen is used in medical applications, metal cutting andwelding, as an oxidizer in rocket fuel, and in water treatment.[39] Oxygen is used in oxyacetylene welding burningacetylene with O2 to produce a very hot flame. In this process, metal up to 60 cm thick is first heated with a smalloxy-acetylene flame and then quickly cut by a large stream of O2.[88] Rocket propulsion requires a fuel and anoxidizer. Larger rockets use liquid oxygen as their oxidizer, which is mixed and ignited with the fuel for propulsion.

Oxygen 13

Scientific

500 million years of climate change vs 18O

Paleoclimatologists measure the ratio ofoxygen-18 and oxygen-16 in the shells andskeletons of marine organisms to determinewhat the climate was like millions of yearsago (see oxygen isotope ratio cycle).Seawater molecules that contain the lighterisotope, oxygen-16, evaporate at a slightlyfaster rate than water molecules containingthe 12% heavier oxygen-18; this disparityincreases at lower temperatures.[89] Duringperiods of lower global temperatures, snowand rain from that evaporated water tends tobe higher in oxygen-16, and the seawaterleft behind tends to be higher in oxygen-18.Marine organisms then incorporate moreoxygen-18 into their skeletons and shellsthan they would in a warmer climate.[89] Paleoclimatologists also directly measure this ratio in the water moleculesof 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 bethe same as those of the primordial solar nebula. However, analysis of a silicon wafer exposed to the solar wind inspace and returned by the crashed Genesis spacecraft has shown that the Sun has a higher proportion of oxygen-16than does the Earth. The measurement implies that an unknown process depleted oxygen-16 from the Sun's disk ofprotoplanetary material prior to the coalescence of dust grains that formed the Earth.[90]

Oxygen presents two spectrophotometric absorption bands peaking at the wavelengths 687 and 760 nm. Someremote sensing scientists have proposed using the measurement of the radiance coming from vegetation canopies inthose bands to characterize plant health status from a satellite platform.[91] This approach exploits the fact that inthose 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 ofvegetation; but it has been proposed as a possible method of monitoring the carbon cycle from satellites on a globalscale.

Oxygen 14

Compounds

Water (H2O) is the most familiaroxygen compound.

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.[92]

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 (H2O) is the oxide of hydrogen and the most familiar oxygen compound.Hydrogen atoms are covalently bonded to oxygen in a water molecule but alsohave an additional attraction (about 23.3 kJ·mol−1 per hydrogen atom) to anadjacent oxygen atom in a separate molecule.[93] These hydrogen bonds betweenwater molecules hold them approximately 15% closer than what would beexpected in a simple liquid with just van der Waals forces.[94] [95]

Oxides, such as iron oxide or rust form whenoxygen combines with other elements.

Due to its electronegativity, oxygen forms chemical bonds with almostall other elements at elevated temperatures to give correspondingoxides. However, some elements readily form oxides at standardconditions for temperature and pressure; the rusting of iron is anexample. The surface of metals like aluminium and titanium areoxidized in the presence of air and become coated with a thin film ofoxide that passivates the metal and slows further corrosion. Some ofthe transition metal oxides are found in nature as non-stoichiometriccompounds, with a slightly less metal than the chemical formula wouldshow. For example, the natural occurring FeO (wüstite) is actuallywritten as Fe1 − xO, where x is usually around 0.05.[96]

Oxygen as a compound is present in the atmosphere in trace quantities in the form of carbon dioxide (CO2). Theearth's crustal rock is composed in large part of oxides of silicon (silica SiO2, found in granite and sand), aluminium(aluminium oxide Al2O3, in bauxite and corundum), iron (iron(III) oxide Fe2O3, in hematite and rust) and othermetals.

The rest of the Earth's crust is also made of oxygen compounds, in particular calcium carbonate (in limestone) andsilicates (in feldspars). Water-soluble silicates in the form of Na4SiO4, Na2SiO3, and Na2Si2O5 are used asdetergents and adhesives.[97]

Oxygen also acts as a ligand for transition metals, forming metal–O2 bonds with the iridium atom in Vaska'scomplex,[98] with the platinum in PtF6,[99] and with the iron center of the heme group of hemoglobin.

Oxygen 15

Organic compounds and biomolecules

Acetone is an important feeder material in thechemical industry.      Oxygen      Carbon

     Hydrogen |alt=A ball structure of a molecule.Its backbone is a zig-zag chain of three carbon

atoms connected in the center to an oxygen atomand on the end to 6 hydrogens.

Oxygen represents more than 40% of themolecular mass of the ATP molecule.

Among the most important classes of organic compounds that containoxygen 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); andamides (R-C(O)-NR2). There are many important organic solvents thatcontain oxygen, including: acetone, methanol, ethanol, isopropanol,furan, THF, diethyl ether, dioxane, ethyl acetate, DMF, DMSO, aceticacid, and formic acid. Acetone ((CH3)2CO) and phenol (C6H5OH) areused 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, andacetamide. Epoxides are ethers in which the oxygen atom is part of aring of three atoms.

Oxygen reacts spontaneously with many organic compounds at orbelow room temperature in a process called autoxidation.[100] Most ofthe organic compounds that contain oxygen are not made by directaction of O2. Organic compounds important in industry and commercethat are made by direct oxidation of a precursor include ethylene oxideand peracetic acid.[97]

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 organiccompounds with biological relevance, carbohydrates contain thelargest proportion by mass of oxygen. All fats, fatty acids, amino acids,and proteins contain oxygen (due to the presence of carbonyl groups inthese acids and their ester residues). Oxygen also occurs in phosphate (PO3−

4) groups in the biologically importantenergy-carrying molecules ATP and ADP, in the backbone and the purines (except adenine) and pyrimidines ofRNA and DNA, and in bones as calcium phosphate and hydroxylapatite.

Oxygen 16

Safety and precautions

Toxicity

Main symptoms of oxygen toxicity.[101]

Oxygen toxicity occurs when lungs take in ahigher than normal O2 partial pressure, which can

occur in deep scuba diving.

Oxygen gas (O2) can be toxic at elevatedpartial pressures, leading to convulsions andother health problems.[81] [102] [103] Oxygentoxicity usually begins to occur at partialpressures more than 50 kilopascals (kPa), or2.5 times the normal sea-level O2 partialpressure of about 21 kPa. Therefore, airsupplied through oxygen masks in medicalapplications is typically composed of 30%O2 by volume (about 30 kPa at standardpressure).[25] At one time, premature babieswere placed in incubators containingO2-rich air, but this practice wasdiscontinued after some babies were blindedby it.[25] [104]

Breathing pure O2 in space applications,such as in some modern space suits, or inearly spacecraft such as Apollo, causes nodamage due to the low total pressuresused.[79] [105] In the case of spacesuits, theO2 partial pressure in the breathing gas is, ingeneral, about 30 kPa (1.4 times normal),and the resulting O2 partial pressure in theastronaut's arterial blood is only marginallymore than normal sea-level O2 partialpressure (see arterial blood gas).

Oxygen toxicity to the lungs and centralnervous system can also occur in deep scubadiving and surface supplied diving.[25] [81]

Prolonged breathing of an air mixture withan O2 partial pressure more than 60 kPa caneventually lead to permanent pulmonaryfibrosis.[106] Exposure to a O2 partialpressures greater than 160 kPa may lead to

convulsions (normally fatal for divers). Acute oxygen toxicity can occur by breathing an air mixture with 21% O2 at66 m or more of depth while the same thing can occur by breathing 100% O2 at only 6 m.[106] [107] [108] [109]

Oxygen 17

Combustion and other hazards

Pure O2 at higher than normal pressure and aspark led to a fire and the loss of the Apollo 1

crew.

Highly concentrated sources of oxygen promote rapid combustion. Fireand explosion hazards exist when concentrated oxidants and fuels arebrought into close proximity; however, an ignition event, such as heator a spark, is needed to trigger combustion.[110] Oxygen itself is not thefuel, but the oxidant. Combustion hazards also apply to compounds ofoxygen with a high oxidative potential, such as peroxides, chlorates,nitrates, perchlorates, and dichromates because they can donate oxygento a fire.

Concentrated O2 will allow combustion to proceed rapidly andenergetically.[110] Steel pipes and storage vessels used to store andtransmit both gaseous and liquid oxygen will act as a fuel; andtherefore the design and manufacture of O2 systems requires specialtraining to ensure that ignition sources are minimized.[110] The fire that killed the Apollo 1 crew in a launch pad testspread so rapidly because the capsule was pressurized with pure O2 but at slightly more than atmospheric pressure,instead of the 1⁄3 normal pressure that would be used in a mission.[111] [112]

Liquid oxygen spills, if allowed to soak into organic matter, such as wood, petrochemicals, and asphalt can causethese materials to detonate unpredictably on subsequent mechanical impact.[110] As with other cryogenic liquids, oncontact with the human body it can cause burns to the skin and the eyes.

See also• Oxygen compounds• Hypoxia, a lack of oxygen• Hypoxia (environmental) for O2 depletion in aquatic ecology• Optode for a method of measuring O2 concentration in solution• Oxygen catastrophe: the sudden rise in oxygen in the atmosphere around 2.4 billion years ago• Oxygen isotope ratio cycle• Oxygen plant• Oxygen sensor• Winkler test for dissolved oxygen• Limiting oxygen concentration

References• Cook, Gerhard A.; Lauer, Carol M. (1968). "Oxygen". in Clifford A. Hampel. The Encyclopedia of the Chemical

Elements. New York: Reinhold Book Corporation. pp. 499–512. LCCN 68-29938.• Emsley, John (2001). "Oxygen". Nature's Building Blocks: An A-Z Guide to the Elements. Oxford, England, UK:

Oxford University Press. pp. 297–304. ISBN 0198503407.• Raven, Peter H.; Ray F. Evert, Susan E. Eichhorn (2005). Biology of Plants, 7th Edition. New York: W.H.

Freeman and Company Publishers. pp. 115–27. ISBN 0-7167-1007-2.

Oxygen 18

Further reading• Walker, J. (1980). "The oxygen cycle". in Hutzinger O.. Handbook of Environmental Chemistry. Volume 1. Part

A: The natural environment and the biogeochemical cycles. Berlin; Heidelberg; New York: Springer-Verlag.pp. 258. ISBN 0387096884.

External links• Oxidizing Agents > Oxygen [113]

• Oxygen (O2) Properties, Uses, Applications [114]

• Roald Hoffmann article on "The Story of O" [115]

• WebElements.com – Oxygen [116]

• Chemistry in its element podcast [117] (MP3) from the Royal Society of Chemistry's Chemistry World: Oxygen[118]

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Oxygen 19

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Article Sources and ContributorsOxygen  Source: http://en.wikipedia.org/w/index.php?oldid=358500807  Contributors: 0612, 123qwe, 1266asdsdjapg, 1297, 84user, Abarenbo, Acalamari, Acroterion, Acs4b, Adambiswanger1,Adashiel, Addshore, Adrian, AdultSwim, Aeros320, AgainErick, Ahoerstemeier, AidepikiW kcuF, Aitias, Alex.muller, AlexG, Alexeymorgunov, Alexf, AlexiusHoratius, Algont, Alias Flood,Alison, All Is One, Alsandro, Amphetamine Analogue, Ancheta Wis, AndonicO, Andre Engels, Andres, Andrewlp1991, Andros 1337, AngelOfSadness, AngryParsley, Angusmclellan, AnnStouter, Anna512, AnonMoos, Antandrus, Anthony Appleyard, Arcadian, Archimerged, Ardric47, Arjun01, Arkuat, ArnoldReinhold, Arnon Chaffin, Art LaPella, Atemperman, AuburnPilot,Axlq, AzaToth, BANZ111, BHS Sux, Bachrach44, Badocter, BalazsH, Ballsonyourwalls, Balthazarduju, BanyanTree, Bart133, Bbatsell, Bbi5291, Bboy14, Beetstra, Beland, Benbest,Bender235, Benjah-bmm27, Benjiboi, Bensderbest, Bevo, Bggoldie, Bhadani, BigCow, Bigbuck, Billcurtis, Blackjack3, BlueEarth, BlueMoonlet, Bluezy, Bobak, Bobo192, Bomac, Bongwarrior,Bonjour amis, Bornfury, Borovy3488, Bradkittenbrink, Bryan Derksen, Buchanan-Hermit, Buckthebronco, Burntsauce, Burzmali, Buzzgrav08, C777, CJLL Wright, CWii, CYD, Cactus.man,Caesura, Caltas, CambridgeBayWeather, Can't sleep, clown will eat me, CanadianCaesar, Candlewicke, CanisRufus, Captain-n00dle, Carlo.milanesi, Carloseduardo, Carnildo, Casliber, Catslash,CattleGirl, Causesobad, Cd12holden, Cdc, CelticJobber, Cfw Master, Chameleon, CharlesGillingham, Charleythegodfather, Chcknpie04, Chilisauce2727, Chlämens, Cholmes75, Chowbok, ChrisDybala, Chris the speller, Chrislk02, Christian List, Christopherlin, Ck lostsword, Clivegrey, Cmapm, Cnaude, Colbuckshot, ColdFeet, Cometstyles, CommonsDelinker, Conn, Kit, Conversionscript, Corpx, Cosmium, Cryptic C62, Crystallina, Ctbolt, Curps, Cybercobra, D, DJ Clayworth, DRosenbach, DVD R W, Dac107, Damieng, Dan56, Dana boomer, DancingPenguin, Daniel Case,DanielCD, Danny, Dantheman531, Danyg, Dark Mage, Darrien, Dauno, David Latapie, Davidj1991, Davumaya, Dawn Bard, Ddday-z, DeadEyeArrow, Deglr6328, Deli nk, Delta G, Demoscn,DerHexer, Derek Ross, Derek.cashman, Deryck Chan, Devl2666, Digitalme, Dillard421, Dirac66, DirectEdge, Discospinster, Dmz5, Doct.proloy, Dolive21, DomCleal, Dominus, DonSiano,Donarreiskoffer, Dragonmaster84, Dravick, Dreadstar, Drini, Droll, Drphilharmonic, Dsyzdek, Dwmyers, Dycedarg, Dysepsion, EEMIV, EL Willy, Ecophreek, EdC, Eddideigel, Edgar181,Edsanville, Edward, Egil, Eilatybartfast, El C, Eldin raigmore, Eleassar, Eliz81, Elkman, Emperorbma, Enemyunknown, Eng02019, Enigmaman, Eric Forste, Eric Kvaalen, Eric119, Erik Zachte,Etaoin, Ethel Aardvark, Eu.stefan, Euyyn, Evercat, Everyking, Evil Monkey, Ewen, Ex nihil, Exarion, Eye.earth, FTGHSmith, Fabartus, Faithlessthewonderboy, Femto, Finell, Firestarterrulz,FisherQueen, Foobaz, FrancoGG, Francs2000, Freakofnurture, Fredrik, Frencheigh, FreplySpang, Frymaster, Funky Monkey, Fvw, G. 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Image Sources, Licenses and Contributors 23

File:Clabecq JPG01.jpg  Source: http://en.wikipedia.org/w/index.php?title=File:Clabecq_JPG01.jpg  License: Creative Commons Attribution-Sharealike 3.0  Contributors: user: Jean-PolGRANDMONTFile:Phanerozoic Climate Change.png  Source: http://en.wikipedia.org/w/index.php?title=File:Phanerozoic_Climate_Change.png  License: unknown  Contributors: Royer, Dana L., Robert A.Berner, Isabel P. Montañez, Neil J. Tabor, and David J. BeerlingFile:Stilles Mineralwasser.jpg  Source: http://en.wikipedia.org/w/index.php?title=File:Stilles_Mineralwasser.jpg  License: GNU Free Documentation License  Contributors: Walter J. Pilsak,Waldsassen, GermanyFile:Rust screw.jpg  Source: http://en.wikipedia.org/w/index.php?title=File:Rust_screw.jpg  License: Creative Commons Attribution 2.0  Contributors: User:Paulnasca. Original uploader wasPaulnasca at en.wikipediaFile:Acetone-3D-vdW.png  Source: http://en.wikipedia.org/w/index.php?title=File:Acetone-3D-vdW.png  License: Public Domain  Contributors: Ben MillsFile:ATP structure.svg  Source: http://en.wikipedia.org/w/index.php?title=File:ATP_structure.svg  License: Public Domain  Contributors: w:User:MysidUser:MysidFile:Symptoms of oxygen toxicity.png  Source: http://en.wikipedia.org/w/index.php?title=File:Symptoms_of_oxygen_toxicity.png  License: Public Domain  Contributors: Mikael HäggströmFile:Scuba-diving.jpg  Source: http://en.wikipedia.org/w/index.php?title=File:Scuba-diving.jpg  License: Creative Commons Attribution-Sharealike 2.5  Contributors: BLueFiSH.as, Civertan,Diwas, Fschoenm, Man vyi, Wikipeder, 3 anonymous editsFile:Apollo 1 fire.jpg  Source: http://en.wikipedia.org/w/index.php?title=File:Apollo_1_fire.jpg  License: Public Domain  Contributors: NASA

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