Oxygen was isolated byMichael Sendivogius before 1604, but it is commonly believed that the element was discovered independently byCarl Wilhelm Scheele, inUppsala, in 1773 or earlier, andJoseph Priestley inWiltshire, in 1774. Priority is often given for Priestley because his work was published first. Priestley, however, called oxygen "dephlogisticated air", and did not recognize it as a chemical element. In 1777,Antoine Lavoisier first recognized oxygen as a chemical element and correctly characterized the role it plays in combustion.
The modern concept of the element oxygen developed over five centuries and included many related discoveries and unsuccessful theories. Multiple people made different contributions to the concept: no one person discovered oxygen.[8][9]
Early experiments
One of the first known experiments on the relationship betweencombustion and air was conducted by the 2nd-century BCE Greek writer on mechanics,Philo of Byzantium. In his workPneumatica, 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.[10] Philo incorrectly surmised that parts of the air in the vessel were converted into theclassical element fire and thus were able to escape through pores in the glass. Many centuries laterIbn al-Nafis, writing in 1250 CE, correctly described oxygenation of blood in the circulatory system;Michael Servetus rediscovered this concept in 1553 but his books were systematically destroyed.[8] A scientifically based and influential description was published byWilliam Harvey in 1628.[11]
Leonardo da Vinci observed that a portion of air is consumed during combustion andrespiration.[12]Polish alchemist, philosopher, and physicianMichael Sendivogius (Michał Sędziwój), writing in 1604,[13] described a substance contained in air, referring to it ascibus vitae ('food of life');[8] this substance is identical with oxygen.[14] During his experiments, performed between 1598 and 1604, Sendivogius properly recognized that the substance is equivalent to the gaseous byproduct released by thethermal decomposition ofpotassium nitrate. However, this important connection was not understood by contemporary scientists like Robert Boyle.[8][14]
Unaware of Sendivogius's work,John Mayow wrote about a portion of air that provided heat in a fire and the human body. This work was ignored because it failed to align with the prevailing phlogiston theory of air and fire. Mayow observed thatantimony increased in weight when heated, and inferred that thenitroaereus must have combined with it.[15] He also thought that the lungs separatenitroaereus from air and pass it into the blood and that animal heat and muscle movement result from the reaction ofnitroaereus with certain substances in the body.[15] Accounts of these and other experiments and ideas were published in 1668 in his workTractatus duo in the tract "De respiratione".[16]
AfterRobert Boyle proved that air is necessary for combustion in the late 17th century, English chemistJohn Mayow (1641–1679) refined this work by showing that fire requires only a part of air that he calledspiritus nitroaereus.[15] 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.[16] From this, he surmised thatnitroaereus is consumed in both respiration and combustion.[17]
Robert Hooke,Ole Borch,Mikhail Lomonosov, andPierre Bayen all produced oxygen in experiments in the 17th and the 18th century but none of them recognized it as achemical element.[18] This may have been in part due to the prevalence of the philosophy of combustion andcorrosion called thephlogiston theory, which was then the favored explanation of those processes.[19]
Established in 1667 by the German alchemistJ. J. Becher, and modified by the chemistGeorg Ernst Stahl by 1731,[20] phlogiston theory stated that all combustible materials were made of two parts. One part, called phlogiston, was given off when the substance containing it was burned, while the dephlogisticated part was thought to be its true form, orcalx.[12]
Highly combustible materials that leave littleresidue, such as wood or coal, were thought to be made mostly of phlogiston; non-combustible substances that corrode, such as iron, contained very little. Air did not play a role in phlogiston 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 objects appear to become lighter and seem to lose something in the process.[12]
Scientific era
Among several contemporaries who had made discoveries independently from one another,Joseph Priestley was the first to publish his findings on oxygen.
Swedish pharmacistCarl Wilhelm Scheele produced and described some properties of oxygen sometime around 1770–1775, but did not publish his work until a few years later[21] because he was unable to interpret his work in the framework of the phlogiston theory.[8] Scheele had produced oxygen gas by heatingmercuric oxide (HgO) and variousnitrates in 1771–1772.[22][23][12] After reading about Priestley's work in 1775, Scheele published in 1777, calling the gas "fire air" because it was then the only knownagent to support combustion.[24]
In the meantime, on August 1, 1774, an experiment conducted by the British clergymanJoseph Priestley focused sunlight on mercuric oxide contained in a glass tube, which liberated a gas he named "dephlogisticated air".[23] He noted that candles burned brighter in the gas and that a mouse was more active and lived longer whilebreathing it. After breathing the gas himself, Priestley wrote: "The feeling of it to my lungs was not sensibly different from that ofcommon air, but I fancied that my breast felt peculiarly light and easy for some time afterwards."[18] 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 titledExperiments and Observations on Different Kinds of Air.[12][25]
The French chemistAntoine Lavoisier later claimed to have discovered the new substance independently. Priestley visited Lavoisier in October 1774 and told him about his experiment and how he liberated the new gas. Scheele had also dispatched a letter to Lavoisier on September 30, 1774, which described his 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).[24]
Lavoisier conducted the first adequate quantitative experiments onoxidation and gave 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 achemical element.[23]
In one experiment, Lavoisier observed that there was no overall increase in weight whentin and air were heated in a closed container.[23] 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 bookSur la combustion en général, which was published in 1777.[23] In that work, he proved that air is a mixture of two gases; 'vital air', which is essential to combustion and respiration, andazote (from Greekἄζωτον 'lifeless'), which did not support either.Azote later becamenitrogen in English, although it has kept the earlier name in French and several other European languages.[23]
Etymology
Lavoisier renamed "vital air" tooxygène in 1777 from theGreek rootsoxys (ὀξύς; "acid", literally 'sharp', from the taste of acids) and-genēs (-γενής; "producer", literally 'begetter'), because he mistakenly believed that oxygen was a constituent of all acids.[26] Chemists (such as SirHumphry Davy in 1812) eventually determined that Lavoisier was wrong in this regard (e.g.Hydrogen chloride (HCl) is a strong acid that does not contain oxygen), but by then the name was too well established.[3]: 793
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 bookThe Botanic Garden (1791) byErasmus Darwin, grandfather ofCharles Darwin.[24]
Later history
John Dalton's originalatomic hypothesis presumed that all elements were monatomic 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, leading to the conclusion that theatomic mass of oxygen was 8 times that of hydrogen, instead of the modern value of about 16.[27] In 1805,Joseph Louis Gay-Lussac andAlexander von Humboldt showed that water is formed of two volumes of hydrogen and one volume of oxygen; and by 1811Amedeo Avogadro had arrived at the correct interpretation of water's composition, based on what is now calledAvogadro's law and the diatomic elemental molecules in those gases.[28]
In 1879 the French brothers Quentin and ArthurBrin discovered acommercially viable reaction to create oxygen. They realized that the known reversible reaction2BaO(s) +O2(g) ↔2BaO2(s) was deactivated by the formation of barium carbonate from carbon dioxide in the air; treating air to remove the carbon dioxide allowed the reaction be reversed indefinitely. Their company used the process between 1886 and 1906 when more economicalfractional distillation began to be used.[29]
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 physicistRaoul Pierre Pictetevaporated liquidsulfur 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 theFrench Academy of Sciences in Paris announcing his discovery ofliquid oxygen.[30] Just two days later, French physicistLouis Paul Cailletet announced his own method of liquefying molecular oxygen.[30] Only a few drops of the liquid were produced in each case and no meaningful analysis could be conducted. Oxygen was liquefied in a stable state for the first time on March 29, 1883, by Polish scientists fromJagiellonian University,Zygmunt Wróblewski andKarol Olszewski.[31]
In 1891 Scottish chemistJames Dewar was able to produce enough liquid oxygen for study.[32] The first commercially viable process for producing liquid oxygen was independently developed in 1895 by German engineerCarl von Linde and British engineerWilliam Hampson. Both men lowered the temperature of air until it liquefied and thendistilled the component gases by boiling them off one at a time and capturing them separately.[33] Later, in 1901, oxyacetylenewelding was demonstrated for the first time by burning a mixture ofacetylene and compressedO 2. This method of welding and cutting metal later became common.[33]
In 1923, the American scientistRobert H. Goddard became the first person to develop arocket engine that burned liquid fuel; the engine usedgasoline for fuel and liquid oxygen as theoxidizer. Goddard successfully flew a small liquid-fueled rocket 56 m at 97 km/h on March 16, 1926, inAuburn, Massachusetts, US.[33][34]
Characteristics
Properties and molecular structure
Orbital diagram, after Barrett (2002),[35] showing the participating atomic orbitals from each oxygen atom, the molecular orbitals that result from their overlap, and theaufbau filling of the orbitals with the 12 electrons, 6 from each O atom, beginning from the lowest-energy orbitals, and resulting in covalent double-bond character from filled orbitals (and cancellation of the contributions of the pairs of σ and σ* and π and π* orbital pairs).
Asdioxygen, two oxygen atoms arechemically bound to each other. The bond can be variously described based on level of theory, but is reasonably and simply described as a covalentdouble bond that results from the filling ofmolecular orbitals formed from theatomic orbitals of the individual oxygen atoms, the filling of which results in abond order of two. More specifically, the double bond is the result of sequential, low-to-high energy, orAufbau, filling of orbitals, and the resulting cancellation of contributions from the 2s electrons, after sequential filling of the low σ and σ* orbitals; σ overlap of the two atomic 2p orbitals that lie along the O–O molecular axis and π overlap of two pairs of atomic 2p orbitals perpendicular to the O–O molecular axis, and then cancellation of contributions from the remaining two 2p electrons after their partial filling of the π* orbitals.[35]
This combination of cancellations and σ and π overlaps results in dioxygen's double-bond character and reactivity, and a triplet electronicground state. Anelectron configuration with two unpaired electrons, as is found in dioxygen orbitals (see the filled π* orbitals in the diagram) that are of equal energy—i.e.,degenerate—is a configuration termed aspin triplet state. Hence, the ground state of theO 2 molecule is referred to astriplet oxygen.[37] The highest-energy, partially filled orbitals areantibonding, and so their filling weakens the bond order from three to two. Because of its unpaired electrons, triplet oxygen reacts only slowly with most organic molecules, which have paired electron spins; this prevents spontaneous combustion.[38]
Liquid oxygen, temporarily suspended in a magnet owing to its paramagnetism
In the triplet form,O 2 molecules areparamagnetic. That is, they impart magnetic character to oxygen when it is in the presence of a magnetic field, because of thespinmagnetic moments of the unpaired electrons in the molecule, and the negativeexchange energy between neighboringO 2 molecules.[32] Liquid oxygen is somagnetic that, in laboratory demonstrations, a bridge of liquid oxygen may be supported against its own weight between the poles of a powerful magnet.[39] Oxygen's paramagnetism can be used analytically in paramagnetic oxygen gas analysers that determine gaseous oxygen concentration, especially in industrial process control and medicine.[40][41]
Singlet oxygen is a name given to several higher-energy species of molecularO 2 in which all the electron spins are paired. It is much more reactive with commonorganic molecules than is normal (triplet) molecular oxygen. In nature, singlet oxygen is commonly formed from water during photosynthesis, using the energy of sunlight.[42] It is also produced in thetroposphere by the photolysis of ozone by light of short wavelength[43] and by theimmune system as a source of active oxygen.[44]Carotenoids in photosynthetic organisms (and possibly 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.[45]
The commonallotrope of elemental oxygen on Earth is calleddioxygen,O 2, the major part of the Earth's atmospheric oxygen (seeOccurrence). O2 has a bond length of 121 pm and a bond energy of 498 kJ/mol.[46] Trioxygen (O 3) is usually known asozone and is a very reactive allotrope of oxygen that is damaging to lung tissue.[47] Ozone is produced in theupper atmosphere whenO 2 combines with atomic oxygen made by the splitting ofO 2 byultraviolet (UV) radiation.[26] Since ozone absorbs strongly in the UV region of thespectrum, theozone layer of the upper atmosphere functions as a protective radiation shield for the planet.[26] Near the Earth's surface, it is apollutant formed as a by-product ofautomobile exhaust.[47] Atlow earth orbit altitudes, sufficient atomic oxygen is present to causecorrosion of spacecraft.[48]
Themetastable moleculetetraoxygen (O 4) was discovered in 2001,[49][50] and was assumed to exist in one of the six phases ofsolid oxygen. In 2006 this phase, created by pressurizingO 2 to 20 GPa, was shown to form arhombohedralO 8cluster.[51] This cluster has the potential to be a much more powerfuloxidizer than eitherO 2 orO 3 and may therefore be used inrocket fuel.[49][50] A metallic phase was discovered in 1990 when solid oxygen is subjected to a pressure of above 96 GPa[52] and it was shown in 1998 that at very low temperatures, this phase becomessuperconducting.[53]
Oxygendissolves more readily in water than nitrogen does. Water in equilibrium with air contains approximately 1 molecule of dissolvedO 2 for every 2 molecules ofN 2 (1:2), compared with 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) dissolves at 0 °C (32 °F) than at 20 °C (68 °F) (7.6 mg/L).[18][54]
At 25 °C (77 °F) and 1standard atmosphere (101.325 kPa) of air, freshwater can dissolve about 6.04 milliliters (mL) of oxygen perliter, andseawater contains about 4.95 mL per liter. At 5 °C (41 °F) the solubility increases to 9.0 mL (50% more than at 25 °C (77 °F)) per liter for freshwater and 7.2 mL (45% more) per liter for sea water.[55]
Oxygen gas dissolved in water at sea-level (milliliters per liter)
5 °C (41 °F)
25 °C (77 °F)
Freshwater
9.00
6.04
Seawater
7.20
4.95
Oxygen condenses at 90.20 K (−182.95 °C, −297.31 °F) and freezes at 54.36 K (−218.79 °C, −361.82 °F).[56] Bothliquid andsolidO 2 are clear substances with a lightsky-blue color caused by absorption in the red (in contrast with the blue color of the sky, which is due toRayleigh scattering of blue light). High-purity liquidO 2 is usually obtained by thefractional distillation of liquefied air.[57] Liquid oxygen may also be condensed from air usingliquid nitrogen as a coolant.[58] Liquid oxygen is a highly reactive substance and must be segregated from combustible materials.[58]
The spectroscopy of molecular oxygen is associated with the atmospheric processes ofaurora andairglow.[59] The absorption in theHerzberg continuum andSchumann–Runge bands in the ultraviolet produces atomic oxygen that is important in the chemistry of the middle atmosphere.[60] Excited-state singlet molecular oxygen is responsible for red chemiluminescence in solution.[61]
Table of thermal and physical properties of oxygen (O2) at atmospheric pressure:[62][63]
16O is the one of the dominant fusion products in massivestars. It issynthesized at the end of thehelium fusion process with some synthesis in theneon burning process.[65] Both17O and18O require seed nuclei.17O is primarily made by the burning of hydrogen intohelium during theCNO cycle, making it a common isotope in the hydrogen burning zones of stars.[65] Most18O is produced when14N (made abundant from CNO burning) captures a4He nucleus, making18O common in the helium-rich zones ofevolved, massive stars.[65]
Fifteenradioisotopes have been characterized, ranging from11O to28O.[6][66] The most stable are15O with ahalf-life of 122.24 seconds and14O with a half-life of 70.606 seconds.[64] All of the remainingradioactive isotopes have half-lives that are less than 27 seconds and the majority of these have half-lives that are less than 83 milliseconds.[64] The most commondecay mode of the isotopes lighter than16O isβ+ decay[67][68][69] to yield nitrogen, and the most common mode for the isotopes heavier than18O isbeta decay to yieldfluorine.[64]
Oxygen is the third most abundant chemical element in the universe, after hydrogen and helium.[71] About 0.9% of theSun's mass is oxygen.[23] Oxygen constitutes 49.2% of theEarth's crust by mass[72] as part of oxide compounds such assilicon dioxide and is the most abundant element by mass in theEarth's crust. It is also the major component of the world's oceans (88.8% by mass).[23] Oxygen gas is the second most common component of theEarth's atmosphere, taking up 20.8% of its volume and 23.1% of its mass (some 1015 tonnes).[23][73][a]
Earth is unusual among the planets of theSolar System in having such a high concentration of oxygen gas in its atmosphere:Mars (with 0.1%O 2 by volume) andVenus have much less. TheO 2 surrounding those planets is produced solely by the action of ultraviolet radiation on oxygen-containing molecules such as carbon dioxide.[74]The unusually high concentration of oxygen gas on Earth is the result of theoxygen cycle. Thisbiogeochemical cycle describes the movement of oxygen within and between its three main reservoirs on Earth: the atmosphere, the biosphere, and thelithosphere. The main driving factor of the oxygen cycle isphotosynthesis, which is responsible for modern Earth's atmosphere. Photosynthesis releases oxygen into the atmosphere, whilerespiration,decay, and combustion remove it from the atmosphere. In the present equilibrium, production and consumption occur at the same rate.[3]: 602
Oxygen levels in the atmosphere are trending slightly downward globally, possibly because of fossil-fuel burning.[75]
Cold water holds more dissolvedO 2.
Free oxygen also occurs in solution in the world's water bodies. The increased solubility ofO 2 at lower temperatures (seePhysical properties) has important implications for ocean life, as polar oceans support a much higher density of life due to their higher oxygen content.[76] Scientists assess this aspect of water quality by measuring the water'sbiochemical oxygen demand, or the amount ofO 2 needed to restore it to a normal concentration.[77]Significant deoxygenation has been observed in tropical oceans. Warming oceans waters are expected to lose oxygen over the next century and into the future for a thousand years; the possible consequences include minimal oxygen zones which are unable to supportmacrofauna.[78]
Paleoclimatologists measure the ratio of oxygen-18 and oxygen-16 in theshells andskeletons of marine organisms to determine the climate millions of years ago (seeoxygen isotope ratio cycle).Seawater molecules that contain the lighterisotope, oxygen-16, evaporate at a slightly faster rate than water molecules containing the 12% heavier oxygen-18, and this disparity increases at lower temperatures.[79] 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.[79] Paleoclimatologists also directly measure this ratio in the water molecules ofice core samples as old as hundreds of thousands of years.[80][81]
Planetary geologists have measured the relative quantities of oxygen isotopes in samples from theEarth, theMoon,Mars, andmeteorites, but were long unable to obtain reference values for the isotope ratios in theSun, believed to be the same as those of theprimordial solar nebula. Analysis of asilicon wafer exposed to thesolar wind in space and returned by the crashedGenesis 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'sdisk of protoplanetary material prior to the coalescence of dust grains that formed the Earth.[82][83]
Oxygen presents two spectrophotometricabsorption bands peaking at the wavelengths 687 and 760 nm. Someremote sensing scientists have proposed using the measurement of the radiance coming from vegetation canopies in those bands to characterize plant health status from asatellite platform.[84] This approach exploits the fact that in those bands it is possible to discriminate the vegetation'sreflectance from itsfluorescence, which is much weaker. The measurement is technically difficult owing to the lowsignal-to-noise ratio and the physical structure of vegetation; but it has been proposed as a possible method of monitoring thecarbon cycle from satellites on a global scale.[85]
Photosynthesis splits water to liberateO 2 and fixesCO 2 into sugar in what is called aCalvin cycle.
In nature,free oxygen is produced as abyproduct oflight-driven splitting of water duringchlorophyllicphotosynthesis. According to some estimates, marinephotoautotrophs such asred/green algae andcyanobacteria provide about 70% of the free oxygen produced on Earth, and the rest is produced in terrestrial environments by plants.[86] Other estimates of the oceanic contribution to atmospheric oxygen are higher, while some estimates are lower, suggesting oceans produce ~45% of Earth's atmospheric oxygen each year.[87]
A simplified overall formula for photosynthesis is[88]
6 CO2 + 6H 2O + photons →C 6H 12O 6 + 6O 2
or simply
carbon dioxide + water + sunlight →glucose + dioxygen
Photolyticoxygen evolution occurs in thethylakoid membranes of photosynthetic organisms and requires the energy of fourphotons.[b] Many steps are involved, but the result is the formation of aproton gradient across the thylakoid membrane, which is used to synthesizeadenosine triphosphate (ATP) viaphotophosphorylation.[89] TheO 2 remaining (after production of the water molecule) is released into the atmosphere.[c]
Oxygen is used inmitochondria ofeukaryotes to generate ATP duringoxidative phosphorylation. The reaction for aerobic respiration is essentially the reverse of photosynthesis and is simplified as
Until the discovery ofanaerobicmetazoa,[92] oxygen was thought to be a requirement for all complex life.[93]
Reactive oxygen species, such assuperoxide ion (O− 2) andhydrogen peroxide (H 2O 2), are reactive by-products of oxygen use in organisms.[73] Parts of theimmune system of higher organisms create peroxide, superoxide, and singlet oxygen to destroy invading microbes. Reactive oxygen species also play an important role in thehypersensitive response of plants against pathogen attack.[89] Oxygen is damaging toobligately anaerobic organisms, which were the dominant form ofearly life on Earth untilO 2 began to accumulate in the atmosphere about 2.5 billion years ago during theGreat Oxygenation Event, about a billion years after the first appearance of these organisms.[94][95]
An adult human at rest inhales 1.8 to 2.4 grams of oxygen per minute.[96] This amounts to more than 6 billion tonnes of oxygen inhaled by humanity per year.[d]
Living organisms
Partial pressures of oxygen in the human body (PO2)
The free oxygenpartial pressure in the body of a living vertebrate organism is highest in therespiratory system, and decreases along anyarterial system, peripheral tissues, andvenous system, respectively. Partial pressure is the pressure that oxygen would have if it alone occupied the volume.[99]
O 2 build-up in Earth's atmosphere: 1) noO 2 produced; 2)O 2 produced, but absorbed in oceans & seabed rock; 3)O 2 starts to gas out of the oceans, but is absorbed by land surfaces and formation of ozone layer; 4–5)O 2 sinks filled and the gas accumulates[100]
Oxygen is both a result of biological activity and a key enabler.Photosynthesis produces oxygen while plants and animals usingaerobic respiration consume it. Consequently theevolution of life is closely related to the concentration of available oxygen. Understanding the relationship between oxygen and evolution would aid in seeking evidence ofextraterrestrial life inexoplanet data.[103]: 252 Oxygen concentration plays a key role in the geochemical composition of sedimentary rocks, making oxygen concentration important for geology and sedimentary rocks important for understanding oxygen concentration over geologic time.[104] The increase in oxygen concentrations had wide-ranging and significant impacts on Earth'sgeochemistry andbiosphere. However, detailed connections between oxygen and evolution remain elusive.[105]
Variations in atmospheric oxygen concentration may have shaped past climates. When oxygen declined, atmospheric density dropped, which in turn increased surface evaporation, causing precipitation increases and warmer temperatures.[106]
Every year one hundred million tonnes ofO 2 are extracted from air for industrial uses.[24] The most common method of extraction isfractional distillation of liquefied air, withN 2distilling as a vapor whileO 2 is left as a liquid.[24] The other primary method of producingO 2 is passing a stream of clean, dry air through one bed of a pair of identicalzeolite molecular sieves, which absorbs the nitrogen and delivers a gas stream that is 90% to 93%O 2.[24] 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 aspressure swing adsorption. Oxygen gas is increasingly obtained by these non-cryogenic technologies (see also the relatedvacuum swing adsorption).[113]
An experiment setup for preparation of oxygen in academic laboratories
In academic laboratories, oxygen can be prepared by heating together potassium chlorate mixed with a small proportion of manganese dioxide.[114]
Oxygen gas can also be produced throughelectrolysis 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. A similar method is the electrocatalyticO 2 evolution from oxides andoxoacids. Chemical catalysts can be used as well, such as inchemical 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 method is forcing air to dissolve throughceramic membranes based onzirconium dioxide by either high pressure or an electric current, to produce nearly pureO 2 gas.[77]
Storage
Oxygen andMAPP gas compressed-gas cylinders with regulators
Oxygen storage methods include high-pressureoxygen tanks, cryogenics and chemical compounds. For reasons of economy, oxygen is often transported in bulk as a liquid in specially insulated tankers, since oneliter of liquefied oxygen is equivalent to 840 liters of gaseous oxygen at atmospheric pressure and 20 °C (68 °F).[24] Such tankers are used to refill bulk liquid-oxygen storage containers, which stand outside hospitals and other institutions that need large volumes of pure oxygen gas. Liquid oxygen is passed throughheat 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 andoxy-fuel welding and cutting.[24]
Uptake ofO 2 from the air is the essential purpose ofrespiration, so oxygen supplementation is used inmedicine. Treatment not only increases oxygen levels in the patient's blood, but has the secondary effect of decreasing resistance to blood flow in many types of diseased lungs, easing work load on the heart.Oxygen therapy is used to treatemphysema,pneumonia, some heart disorders (congestive heart failure), some disorders that cause increasedpulmonary artery pressure, and anydisease that impairs the body's ability to take up and use gaseous oxygen.[115]
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 ofoxygen masks ornasal cannulas.[116]
Hyperbaric (high-pressure) medicine uses specialoxygen chambers to increase thepartial pressure ofO 2 around the patient and, when needed, the medical staff.[117]Carbon monoxide poisoning,gas gangrene, anddecompression sickness (the 'bends') are sometimes addressed with this therapy.[118] IncreasedO 2 concentration in the lungs helps to displacecarbon monoxide from the heme group ofhemoglobin.[119][120] Oxygen gas is poisonous to theanaerobic bacteria that cause gas gangrene, so increasing its partial pressure helps kill them.[121][122] Decompression sickness occurs in divers who decompress too quickly after a dive, resulting in bubbles of inert gas, mostly nitrogen and helium, forming in the blood. Increasing the pressure ofO 2 as soon as possible helps to redissolve the bubbles back into the blood so that these excess gasses can be exhaled naturally through the lungs.[115][123][124] Normobaric oxygen administration at the highest available concentration is frequently used as first aid for any diving injury that may involve inert gas bubble formation in the tissues. There is epidemiological support for its use from a statistical study of cases recorded in a long term database.[125][126][127]
In modernspace suits, which surround their occupant's body, oxygen gas is used as a low-pressurebreathing gas. These devices use nearly pure oxygen at about one-third normal pressure, resulting in a normal blood partial pressure ofO 2. This trade-off of higher oxygen concentration for lower pressure is needed to maintain suit flexibility.[128][129]
Scuba andsurface-suppliedunderwater divers andsubmarines also rely on artificially deliveredO 2. Submarines, submersibles andatmospheric diving suits usually operate at normal atmospheric pressure. Breathing air is scrubbed of carbon dioxide by chemical extraction and oxygen is replaced to maintain a constant partial pressure.Ambient pressure divers breathe air or gas mixtures with an oxygen fraction suited to the operating depth. Pure or nearly pureO 2 use in diving at pressures higher than atmospheric is usually limited torebreathers, ordecompression at relatively shallow depths (~6 meters depth, or less),[130][131] ormedical treatment in recompression chambers at pressures up to 2.8 bar, where acute oxygen toxicity can be managed without the risk of drowning. Deeper diving requires significant dilution ofO 2 with other gases, such as nitrogen or helium, to preventoxygen toxicity.[130]
People who climb mountains or fly in non-pressurizedfixed-wing aircraft sometimes have supplementalO 2 supplies.[f] Pressurized commercial airplanes have an emergency supply ofO 2 automatically supplied to the passengers in case of cabin depressurization. Sudden cabin pressure loss activateschemical oxygen generators above each seat, causingoxygen masks to drop. Pulling on the masks "to start the flow of oxygen" as cabin safety instructions dictate, forces iron filings into thesodium chlorate inside the canister.[77] A steady stream of oxygen gas is then produced by theexothermic reaction.[3]
Oxygen, as a mildeuphoric, has a history of recreational use inoxygen bars and insports. Oxygen bars are establishments found in the United States since the late 1990s that offer higher than normalO 2 exposure for a minimal fee.[132] Professional athletes, especially inAmerican football, sometimes go off-field between plays to don oxygen masks to boost performance. The pharmacological effect is doubted; aplacebo effect is a more likely explanation.[132] Available studies support a performance boost from oxygen enriched mixtures only if it is inhaledduringaerobic exercise.[133]
Other recreational uses that do not involve breathing includepyrotechnic applications, such asGeorge Goble's five-second ignition ofbarbecue grills.[134]
Smelting ofiron ore intosteel consumes 55% of commercially produced oxygen.[77] In this process,O 2 is injected through a high-pressure lance into molten iron, which removessulfur impurities and excesscarbon as the respective oxides,SO 2 andCO 2. The reactions areexothermic, so the temperature increases to 1,700 °C.[77]
Another 25% of commercially produced oxygen is used by the chemical industry.[77]Ethylene is reacted withO 2 to createethylene oxide, which, in turn, is converted intoethylene glycol; the primary feeder material used to manufacture a host of products, includingantifreeze andpolyester polymers (the precursors of manyplastics andfabrics).[77]
Most of the remaining 20% of commercially produced oxygen is used in medical applications,metal cutting and welding, as an oxidizer inrocket fuel, and inwater treatment.[77] Oxygen is used inoxyacetylene welding, burningacetylene withO 2 to produce a very hot flame. In this process, metal up to 60 cm (24 in) thick is first heated with a small oxy-acetylene flame and then quickly cut by a large stream ofO 2.[135]
Water (H 2O) is an oxide ofhydrogen and the most familiar oxygen compound. Hydrogen atoms arecovalently bonded to oxygen in a water molecule but also have an additional attraction (about 23.3 kJ/mol per hydrogen atom) to an adjacent oxygen atom in a separate molecule.[137] Thesehydrogen bonds between water molecules hold them approximately 15% closer than what would be expected in a simple liquid with justvan der Waals forces.[138][g]
Oxides, such asiron oxide orrust, form when oxygen combines with other elements.
TheNFPA 704 standard rates compressed oxygen gas as nonhazardous to health, nonflammable and nonreactive, but an oxidizer. Refrigerated liquid oxygen (LOX) is given a health hazard rating of 3 (for increased risk ofhyperoxia from condensed vapors, and for hazards common to cryogenic liquids such as frostbite), and all other ratings are the same as the compressed gas form.[146]
Oxygen gas (O 2) can betoxic at elevatedpartial pressures, leading toconvulsions and other health problems.[130][h][148] Oxygen toxicity usually begins to occur at partial pressures more than 50 kilopascals (kPa), equal to about 50% oxygen composition at standard pressure or 2.5 times the normal sea-levelO 2 partial pressure of about 21 kPa. This is not a problem except for patients onmechanical ventilators, since gas supplied throughoxygen masks in medical applications is typically composed of only 30–50%O 2 by volume (about 30 kPa at standard pressure).[18]
At one time,premature babies were placed in incubators containingO 2-rich air, but this practice was discontinued after some babies were blinded by the oxygen content being too high.[18]
Breathing pureO 2 in space applications, such as in some modern space suits, or in early spacecraft such asApollo, causes no damage due to the low total pressures used.[128][149] In the case of spacesuits, theO 2 partial pressure in the breathing gas is, in general, about 30 kPa (1.4 times normal), and the resultingO 2 partial pressure in the astronaut's arterial blood is only marginally more than normal sea-levelO 2 partial pressure.[150]
Oxygen toxicity to the lungs andcentral nervous system can also occur in deepscuba diving andsurface-supplied diving.[18][130] Prolonged breathing of an air mixture with anO 2 partial pressure more than 60 kPa can eventually lead to permanentpulmonary fibrosis.[151] Exposure to anO 2 partial pressure greater than 160 kPa (about 1.6 atm) may lead to convulsions (normally fatal for divers). Acute oxygen toxicity (causing seizures, its most feared effect for divers) can occur by breathing an air mixture with 21%O 2 at 66 m (217 ft) or more of depth; the same thing can occur by breathing 100%O 2 at only 6 m (20 ft).[151][152][153][154]
Combustion and other hazards
The interior of theApollo 1 Command Module. PureO 2 at higher than normal pressure and a spark led to a fire and the loss of the Apollo 1 crew.
Unless non-flammable containers are used or all sources of ignition are eliminated, oxygen rich environments are extremely hazardous. Many materials including most metals burn faster in oxygen rich environments and ignite at lower temperatures.[155] ConcentratedO 2 will allow combustion to proceed rapidly and energetically.[38] Steel pipes and storage vessels used to store and transmit both gaseous andliquid oxygen will act as a fuel; and therefore the design and manufacture ofO 2 systems requires special training to ensure that ignition sources are minimized.[38]
The 1967 fire that killed theApollo 1 crew in a launch pad test spread so rapidly because the capsule was pressurized with pureO 2 at slightly more than atmospheric pressure to force the doors against their seals. Once in space the exterior pressure would be minimal and the oxygen pressure would have been reduced to 5psi to reduce fire hazard. The ignition source for the fire was traced to an electrical surge under the seat of one of the astronauts, probably due to chafed wires.[156]
Liquid oxygen spills, if allowed to soak into organic matter, such aswood,petrochemicals, andasphalt can cause these materials todetonate unpredictably on subsequent mechanical impact.[38]
^Figures given are for values up to 80 km (50 mi) above the surface
^Thylakoid membranes are part ofchloroplasts in algae and plants while they simply are one of many membrane structures in cyanobacteria. In fact, chloroplasts are thought to have evolved fromcyanobacteria that were once symbiotic partners with the progenitors of plants and algae.
^Water oxidation is catalyzed by amanganese-containingenzyme complex known as theoxygen evolving complex (OEC) or water-splitting complex found associated with the lumenal side of thylakoid membranes. Manganese is an importantcofactor, andcalcium andchloride are also required for the reaction to occur. (Raven 2005)
^abcdDerived from mmHg values using 0.133322 kPa/mmHg
^The reason is that increasing the proportion of oxygen in the breathing gas at low pressure acts to augment the inspiredO 2 partial pressure nearer to that found at sea-level.
^Also, since oxygen has a higher electronegativity than hydrogen, the charge difference makes it apolar molecule. The interactions between the differentdipoles of each molecule cause a net attraction force.
^SinceO 2's partial pressure is the fraction ofO 2 times the total pressure, elevated partial pressures can occur either from highO 2 fraction in breathing gas or from high breathing gas pressure, or a combination of both.
^Arblaster, John W. (2018).Selected Values of the Crystallographic Properties of Elements. Materials Park, Ohio: ASM International.ISBN978-1-62708-155-9.
^Weast, Robert (1984).CRC, Handbook of Chemistry and Physics. Boca Raton, Florida: Chemical Rubber Company Publishing. pp. E110.ISBN0-8493-0464-4.
^De Lapide Philosophorum Tractatus duodecim e naturae fonte et manuali experientia depromti (Twelve Treatises on the Philosopher's Stone drawn from the source of nature and manual experience; 1604)
^abJack Barrett, 2002, "Atomic Structure and Periodicity", (Basic concepts in chemistry, Vol. 9 of Tutorial chemistry texts), Cambridge, UK: Royal Society of Chemistry, p. 153,ISBN0854046577. SeeGoogle Books.Archived May 30, 2020, at theWayback Machine accessed January 31, 2015.
^"Oxygen Facts". Science Kids. February 6, 2015.Archived from the original on May 7, 2020. RetrievedNovember 14, 2015.
^abcdWerley, Barry L., ed. (1991).ASTM Technical Professional training.Fire Hazards in Oxygen Systems. Philadelphia:ASTM International Subcommittee G-4.05.
^Desgreniers, S.; Vohra, Y. K.; Ruoff, A. L. (1990). "Optical response of very high density solid oxygen to 132 GPa".J. Phys. Chem.94 (3):1117–22.Bibcode:1990JPhCh..94.1117D.doi:10.1021/j100366a020.
^Kearns, David R. (1971). "Physical and chemical properties of singlet molecular oxygen".Chemical Reviews.71 (4):395–427.doi:10.1021/cr60272a004.
^Holman, Jack P. (2002).Heat transfer (9th ed.). New York, NY: McGraw-Hill Companies, Inc. pp. 600–606.ISBN9780072406559.OCLC46959719.
^Incropera 1 Dewitt 2 Bergman 3 Lavigne 4, Frank P. 1 David P. 2 Theodore L. 3 Adrienne S. 4 (2007).Fundamentals of heat and mass transfer (6th ed.). Hoboken, NJ: John Wiley and Sons, Inc. pp. 941–950.ISBN9780471457282.OCLC62532755.{{cite book}}: CS1 maint: numeric names: authors list (link)
^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 notes that according to later articles inNature, the values appear to be about 3% too high.
^Miller, J. R.; Berger, M.; Alonso, L.; Cerovic, Z.; et al. (2003).Progress on the development of an integrated canopy fluorescence model.Geoscience and Remote Sensing Symposium, 2003. IGARSS '03. Proceedings. 2003 IEEE International. Vol. 1. pp. 601–603.CiteSeerX10.1.1.473.9500.doi:10.1109/IGARSS.2003.1293855.ISBN0-7803-7929-2.
^Hall, John (2011).Guyton and Hall textbook of medical physiology (12th ed.). Philadelphia, Pa.: Saunders/Elsevier. p. 5.ISBN978-1-4160-4574-8.
^Pocock, Gillian; Richards, Christopher D. (2006).Human physiology : the basis of medicine (3rd ed.). Oxford: Oxford University Press. p. 311.ISBN978-0-19-856878-0.
^Ward, Peter D.; Brownlee, Donald (2000).Rare Earth: Why Complex Life is Uncommon in the Universe. Copernicus Books (Springer Verlag). p. 217.ISBN978-0-387-98701-9.
^abNormal Reference Range TableArchived December 25, 2011, at theWayback Machine from The University of Texas Southwestern Medical Center at Dallas. Used in Interactive Case Study Companion to Pathologic basis of disease.
^Undersea and Hyperbaric Medical Society."Carbon Monoxide". Archived fromthe original on July 25, 2008. RetrievedSeptember 22, 2008.
^Piantadosi CA (2004)."Carbon monoxide poisoning".Undersea Hyperb Med.31 (1):167–77.PMID15233173. Archived from the original on February 3, 2011. RetrievedSeptember 22, 2008.
^Webb JT; Olson RM; Krutz RW; Dixon G; Barnicott PT (1989). "Human tolerance to 100% oxygen at 9.5 psia during five daily simulated 8-hour EVA exposures".Aviat Space Environ Med.60 (5):415–21.doi:10.4271/881071.PMID2730484.
^Dharmeshkumar N Patel; Ashish Goel; SB Agarwal; Praveenkumar Garg; et al. (2003)."Oxygen Toxicity"(PDF).Indian Academy of Clinical Medicine.4 (3): 234. Archived fromthe original(PDF) on September 22, 2015. RetrievedApril 26, 2009.
Cook, Gerhard A.; Lauer, Carol M. (1968)."Oxygen". In Clifford A. Hampel (ed.).The Encyclopedia of the Chemical Elements. New York: Reinhold Book Corporation. pp. 499–512.LCCN68-29938.
Emsley, John (2001)."Oxygen".Nature's Building Blocks: An A–Z Guide to the Elements. Oxford, England: Oxford University Press. pp. 297–304.ISBN978-0-19-850340-8.
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