Oxygen
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↑
O
↓
S
(-218.79 °C, -361.82 °F)
(-182.95 °C, -297.31 °F)
29.378 J·mol−1·K−1
(neutral oxide)
(more) 1st: 1313.9 kJ·mol−1
Oxygen (from the Greek roots ὀξÏÏ‚ (oxys) (acid, literally "sharp," from the taste of acids) and -γενής (-genÄ“s) (producer, literally begetter)is the element with atomic number 8 and represented by the symbol O. It is a member of the chalcogen group on the periodic table, and is a highly reactive nonmetallic period 2 element that readily forms compounds (notably oxides) with almost all other elements. At standard temperature and pressure two atoms of the element bind to form dioxygen, a colorless, odorless, tasteless diatomic gas with the formula O2. Oxygen is the third most abundant element in the universe by mass after hydrogen and helium[1] and the most abundant element by mass in the Earth's crust.[2] Oxygen constitutes 88.8% of the mass of water and 20.9% of the volume of air[clarify].[3]
All major classes of structural molecules in living organisms, such as proteins, carbohydrates, and fats, contain oxygen, as do the major inorganic compounds that comprise animal shells, teeth, and bone. Oxygen in the form of O2 is produced from water by cyanobacteria, algae and plants during photosynthesis and is used in cellular respiration for all complex life. Oxygen is toxic to obligately anaerobic organisms, which were the dominant form of early life on Earth until O2 began to accumulate in the atmosphere 2.5 billion years ago.[4] Another form (allotrope) of oxygen, ozone (O3), helps protect the biosphere from ultraviolet radiation with the high-altitude ozone layer, but is a pollutant near the surface where it is a by-product of smog.
Oxygen was independently discovered by Joseph Priestley in Wiltshire, in 1774, and Carl Wilhelm Scheele, in Uppsala, a year earlier, but Priestley is often given priority because he published his findings first. The name oxygen was coined in 1777 by Antoine Lavoisier,[5] whose experiments with oxygen helped to discredit the then-popular phlogiston theory of combustion and corrosion. Oxygen is produced industrially by fractional distillation of liquefied air, use of zeolites to remove carbon dioxide and nitrogen from air, electrolysis of water and other means. Uses of oxygen include the production of steel, plastics and textiles; rocket propellant; oxygen therapy; and life support in aircraft, submarines, spaceflight and diving.
Contents
Characteristics
Structure
At standard temperature and pressure, oxygen is a colorless, odorless gas with the molecular formula O2, in which the two oxygen atoms are chemically bonded to each other with a spin triplet electron configuration. This bond has a bond order of two, and is often simplified in description as a double bond[6] or as a combination of one two-electron bond and two three-electron bonds.[7]
Triplet oxygen is the ground state of the O2 molecule.[8] The electron configuration of the molecule has two unpaired electrons occupying two degenerate molecular orbitals.[9] These orbitals are classified as antibonding (weakening the bond order from three to two), so the diatomic oxygen bond is weaker than the diatomic nitrogen triple bond in which all bonding molecular orbitals are filled, but some antibonding orbitals are not.[8]
In normal triplet form, O2 molecules are paramagnetic—they form a magnet in the presence of a magnetic field—because of the spin magnetic moments of the unpaired electrons in the molecule, and the negative exchange energy between neighboring O2 molecules.[10] Liquid oxygen is attracted to a magnet to a sufficient extent that, in laboratory demonstrations, a bridge of liquid oxygen may be supported against its own weight between the poles of a powerful magnet.[11][12]
Singlet oxygen, a name given to several higher-energy species of molecular O2 in which all the electron spins are paired, is much more reactive towards common organic molecules. In nature, singlet oxygen is commonly formed from water during photosynthesis, using the energy of sunlight.[13] It is also produced in the troposphere by the photolysis of ozone by light of short wavelength,[14] and by the immune system as a source of active oxygen.[15]Carotenoids in photosynthetic organisms (and possibly also in animals) play a major role in absorbing energy from singlet oxygen and converting it to the unexcited ground state before it can cause harm to tissues.[16]
Allotropes
The common allotrope of elemental oxygen on Earth is called dioxygen, O2. It has a bond length of 121 pm and a bond energy of 498 kJ·mol-1.[17] This is the form that is used by complex forms of life, such as animals, in cellular respiration (see Biological role) and is the form that is a major part of the Earth's atmosphere (see Occurrence). Other aspects of O2 are covered in the remainder of this article.
Trioxygen (O3) is usually known as ozone and is a very reactive allotrope of oxygen that is damaging to lung tissue.[18] Ozone is produced in the upper atmosphere when O2 combines with atomic oxygen made by the splitting of O2 by ultraviolet (UV) radiation.[5] Since ozone absorbs strongly in the UV region of the spectrum, it functions as a protective radiation shield for the planet (see ozone layer).[5] Near the earth's surface, however, it is a pollutant formed as a by-product of automobile exhaust.[19]
The metastable molecule tetraoxygen (O4) was discovered in 2001,[20][21] and was assumed to exist in one of the six phases of solid oxygen. It was proven in 2006 that that phase, created by pressurizing O2 to 20 GPa, is in fact a rhombohedral O8 cluster.[22] This cluster has the potential to be a much more powerful oxidizer than either O2 or O3 and may therefore be used in rocket fuel.[20][21] A metallic phase was discovered in 1990 when solid oxygen is subjected to a pressure of above 96 GPa[23] and it was shown in 1998 that at very low temperatures, this phase becomes superconducting.[24]
Physical properties
- See also: Liquid oxygen and solid oxygen
Oxygen is more soluble in water than nitrogen; water contains approximately 1 molecule of O2 for every 2 molecules of N2, compared to an atmospheric ratio of approximately 1:4. The solubility of oxygen in water is temperature-dependent, and about twice as much (14.6 mg·L−1) dissolves at 0 °C than at 20 °C (7.6 mg·L−1).[25][26] At 25 °C and 1 atm of air, freshwater contains about 6.04 milliliters (mL) of oxygen per liter, whereas seawater contains about 4.95 mL per liter.[27] At 5 °C the solubility increases to 9.0 mL (50% more than at 25 °C) per liter for water and 7.2 mL (45% more) per liter for sea water.
Oxygen condenses at 90.20 K (−182.95 °C, −297.31 °F), and freezes at 54.36 K (−218.79 °C, −361.82 °F).[28] Both liquid and solid O2 are clear substances with a light sky-blue color caused by absorption in the red (in contrast with the blue color of the sky, which is due to Rayleigh scattering of blue light). High-purity liquid O2 is usually obtained by the fractional distillation of liquefied air;[29] Liquid oxygen may also be produced by condensation out of air, using liquid nitrogen as a coolant. It is a highly-reactive substance and must be segregated from combustible materials.[30]
Isotopes and stellar origin
Naturally occurring oxygen is composed of three stable isotopes, 16O, 17O, and 18O, with 16O being the most abundant (99.762% natural abundance).[31] Oxygen isotopes range in mass number from 12 to 28.[31]
Most 16O is synthesized at the end of the helium fusion process in stars but some is made in the neon burning process.[32]17O is primarily made by the burning of hydrogen into helium during the CNO cycle, making it a common isotope in the hydrogen burning zones of stars.[32] Most 18O is produced when 14N (made abundant from CNO burning) captures a 4He nucleus, making 18O common in the helium-rich zones of stars.[32]
Fourteen radioisotopes have been characterized, the most stable being 15O with a half-life of 122.24 seconds (s) and 14O with a half-life of 70.606 s.[31] All of the remaining radioactive isotopes have half-lives that are less than 27 s and the majority of these have half-lives that are less than 83 milliseconds.[31] The most common decay mode of the isotopes lighter than 16O is electron capture to yield nitrogen, and the most common mode for the isotopes heavier than 18O is beta decay to yield fluorine.[31]
Occurrence
- See also: Silicate minerals and Category:Oxide minerals
Oxygen is the most abundant chemical element, by mass, in our biosphere, air, sea and land. Oxygen is the third most abundant chemical element in the universe, after hydrogen and helium.[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).[33][3][34] 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 ultraviolet radiation impacting oxygen-containing molecules such as carbon dioxide.
The unusually high concentration of oxygen gas on Earth is the result of the oxygen cycle. This biogeochemical cycle describes the movement of oxygen within and between its three main reservoirs on Earth: the atmosphere, the biosphere, and the lithosphere. The main driving factor of the oxygen cycle is photosynthesis, which is responsible for modern Earth's atmosphere. Photosynthesis releases oxygen into the atmosphere, while respiration and decay remove it from the atmosphere. In the present equilibrium, production and consumption occur at the same rate of roughly 1/2000th of the entire atmospheric oxygen per year.
Free oxygen also occurs in solution in the world's water bodies. The increased solubility of O2 at lower temperatures (see Physical properties) has important implications for ocean life, as polar oceans support a much higher density of life due to their higher oxygen content.[35]Polluted water may have reduced amounts of O2 in it, depleted by decaying algae and other biomaterials (see eutrophication). Scientists assess this aspect of water quality by measuring the water's biochemical oxygen demand, or the amount of O2 needed to restore it to a normal concentration.[36]
Biological role
Photosynthesis and respiration
In nature, free oxygen is produced by the light-driven splitting of water during oxygenic photosynthesis. Green algae and cyanobacteria in marine environments provide about 70% of the free oxygen produced on earth and the rest is produced by terrestrial plants.[37]
A simplified overall formula for photosynthesis is:[38]
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- 6CO2 + 6H2O + photons → C6H12O6 + 6O2 (or simply carbon dioxide + water + sunlight → glucose + dioxygen)
Photolytic oxygen evolution occurs in the thylakoid membranes of photosynthetic organisms and requires the energy of four photons.[39] Many steps are involved, but the result is the formation of a proton gradient across the thylakoid membrane, which is used to synthesize ATP via photophosphorylation.[40] The O2 remaining after oxidation of the water molecule is released into the atmosphere.[41]
Molecular dioxygen, O2, is essential for cellular respiration in all aerobic organisms. Oxygen is used in mitochondria to help generate adenosine triphosphate (ATP) during oxidative phosphorylation. The reaction for aerobic respiration is essentially the reverse of photosynthesis and is simplified as:
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- C6H12O6 + 6O2 → 6CO2 + 6H2O + 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.[42][18] Other animals use hemocyanin (molluscs and some arthropods) or hemerythrin (spiders and lobsters).[33] A liter of blood can dissolve 200 cc of O2.[33]
Reactive oxygen species, such as superoxide ion (O2−) and hydrogen peroxide (H2O2), are dangerous by-products of oxygen use in organisms.[33] Parts of the immune system of higher organisms, however, create peroxide, superoxide, and singlet oxygen to destroy invading microbes. Reactive oxygen species also play an important role in the hypersensitive response of plants against pathogen attack.[40]
An adult human in rest inhales 1.8 to 2.4 grams of oxygen per minute.[43] This amounts to more than 6 billion tonnes of oxygen inhaled by humanity per year. [44]
Build-up in the atmosphere
Free oxygen gas was almost nonexistent in Earth's atmosphere before photosynthetic archaea and bacteria evolved. Free oxygen first appeared in significant quantities during the Paleoproterozoic era (between 2.5 and 1.6 billion years ago). At first, the oxygen combined with dissolved iron in the oceans to form banded iron formations. Free oxygen started to gas out of the oceans 2.7 billion years ago, reaching 10% of its present level around 1.7 billion years ago.[45]
The presence of large amounts of dissolved and free oxygen in the oceans and atmosphere may have driven most of the anaerobic organisms then living to extinction during the oxygen catastrophe about 2.4 billion years ago. However, cellular respiration using O2 enables aerobic organisms to produce much more ATP than anaerobic organisms, helping the former to dominate Earth's biosphere.[46] Photosynthesis and cellular respiration of O2 allowed for the evolution of eukaryotic cells and ultimately complex multicellular organisms such as plants and animals.
Since the beginning of the Cambrian era 540 million years ago, O2 levels have fluctuated between 15% and 30% per volume.[47] Towards the end of the Carboniferous era (about 300 million years ago) atmospheric O2 levels reached a maximum of 35% by volume,[47] allowing insects and amphibians to grow much larger than today's species. Human activities, including the burning of 7 billion tonnes of fossil fuels each year have had very little effect on the amount of free oxygen in the atmosphere.[10] At the current rate of photosynthesis it would take about 2,000 years to regenerate the entire O2 in the present atmosphere.[48]
History
Early experiments
One of the first known experiments on the relationship between combustion and air was conducted by the second century BCE Greek writer on mechanics, Philo of Byzantium. In his work Pneumatica, Philo observed that inverting a vessel over a burning candle and surrounding the vessel's neck with water resulted in some water rising into the neck.[49] Philo incorrectly surmised that parts of the air in the vessel were converted into the classical element fire and thus were able to escape through pores in the glass. Many centuries later Leonardo da Vinci built on Philo's work by observing that a portion of air is consumed during combustion and respiration.[50]
In the late 17th century, Robert Boyle proved that air is necessary for combustion. English chemist John Mayow refined this work by showing that fire requires only a part of air that he called spiritus nitroaereus or just nitroaereus.[51] 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.[52] From this he surmised that nitroaereus is consumed in both respiration and combustion.
Mayow observed that antimony increased in weight when heated, and inferred that the nitroaereus must have combined with it.[51] He also thought that the lungs separate nitroaereus from air and pass it into the blood and that animal heat and muscle movement result from the reaction of nitroaereus with certain substances in the body.[51] Accounts of these and other experiments and ideas were published in 1668 in his work Tractatus duo in the tract "De respiratione".[52]
Phlogiston theory
Robert Hooke, Ole Borch, Mikhail Lomonosov, and Pierre Bayen all produced oxygen in experiments in the 17th century but none of them recognized it as an element.[25] This may have been in part due to the prevalence of 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 the chemist Georg Ernst Stahl by 1731,[53] 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, or calx.[50]
Highly combustible materials that leave little residue, such as wood or coal, were thought to be made mostly of phlogiston; whereas 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.[50] The fact that a substance like wood actually gains overall weight in burning was hidden by the buoyancy of the gaseous combustion products. Indeed one of the first clues that the phlogiston theory was incorrect was that metals, too, gain weight in rusting (when they were supposedly losing phlogiston).
Discovery
Oxygen was first discovered by Swedish pharmacist Carl Wilhelm Scheele. He had produced oxygen gas by heating mercuric oxide and various nitrates by about 1772.[50][3] Scheele called the gas 'fire air' because it was the only known supporter of combustion. He wrote an account of this discovery in a manuscript he titled Treatise on Air and Fire, which he sent to his publisher in 1775. However, that document was not published until 1777.[54]
In the meantime, an experiment was conducted by the British clergyman Joseph Priestley on August 1, 1774 focused sunlight on mercuric oxide (HgO) inside a glass tube, which liberated a gas he named 'dephlogisticated air'.[3] He noted that candles burned brighter in the gas and that a mouse was more active and lived longer while breathing it. After breathing the gas himself, he wrote: "The feeling of it to my lungs was not sensibly different from that of common air, but I fancied that my breast felt peculiarly light and easy for some time afterwards."[25] Priestley published his findings in 1775 in a paper titled "An Account of Further Discoveries in Air" which was included in the second volume of his book titled Experiments and Observations on Different Kinds of Air.[55][50] Because he published his findings first, Priestley is usually given priority in the discovery.
The noted French chemist Antoine Laurent Lavoisier later claimed to have discovered the new substance independently. However, Priestley visited Lavoisier in October 1774 and told him about his experiment and how he liberated the new gas. Scheele also posted a letter to Lavoisier on September 30, 1774 that described his own discovery of the previously-unknown substance, but Lavoisier never acknowledged receiving it (a copy of the letter was found in Scheele's belongings after his death).[54]
Lavoisier's contribution
What Lavoisier did indisputably do (although this was disputed at the time) was to conduct the first adequate quantitative experiments on oxidation and give the first correct explanation of how combustion works.[3] He used these and similar experiments, all started in 1774, to discredit the phlogiston theory and to prove that the substance discovered by Priestley and Scheele was a chemical element.
In one experiment, Lavoisier observed that there was no overall increase in weight when tin and air were heated in a closed container.[3] He noted that air rushed in when he opened the container, which indicated that part of the trapped air had been consumed. He also noted that the tin had increased in weight and that increase was the same as the weight of the air that rushed back in. This and other experiments on combustion were documented in his book Sur la combustion en général, which was published in 1777.[3] In that work, he proved that air is a mixture of two gases; 'vital air', which is essential to combustion and respiration, and azote (Gk. ἄζωτον "lifeless"), which did not support either. Azote later became nitrogen in English, although it has kept the name in French and several other European languages.[3]
Lavoisier renamed 'vital air' to oxygène in 1777 from the Greek roots ὀξÏÏ‚ (oxys) (acid, literally "sharp," from the taste of acids) and -γενής (-genÄ“s) (producer, literally begetter), because he mistakenly believed that oxygen was a constituent of all acids.[5] Chemists eventually determined that Lavoisier was wrong in this regard, but by that time it was too late, the name had taken. Actually, the gas that could more appropriately have been given the description, "acid producer," is hydrogen.
Oxygen entered the English language despite opposition by English scientists and the fact that the Englishman Priestley had first isolated the gas and written about it. This is partly due to a poem praising the gas titled "Oxygen" in the popular book The Botanic Garden (1791) by Erasmus Darwin, grandfather of Charles Darwin.[54]
Later history
John Dalton's original atomic hypothesis assumed that all elements were monoatomic and that the atoms in compounds would normally have the simplest atomic ratios with respect to one another. For example, Dalton assumed that water's formula was HO, giving the atomic mass of oxygen as 8 times that of hydrogen, instead of the modern value of about 16.[56] In 1805, Joseph Louis Gay-Lussac and Alexander von Humboldt showed that water is formed of two volumes of hydrogen and one volume of oxygen; and by 1811 Amedeo Avogadro had arrived at the correct interpretation of water's composition, based on what is now called Avogadro's law and the assumption of diatomic elemental molecules.[57][58]
By the late 19th century scientists realized that air could be liquefied, and its components isolated, by compressing and cooling it. Using a cascade method, Swiss chemist and physicist Raoul Pierre Pictet evaporated liquid sulfur dioxide in order to liquefy carbon dioxide, which in turn was evaporated to cool oxygen gas enough to liquefy it. He sent a telegram on December 22, 1877 to the French Academy of Sciences in Paris announcing his discovery of liquid oxygen.[59] Just two days later, French physicist Louis Paul Cailletet announced his own method of liquefying molecular oxygen.[59] Only a few drops of the liquid were produced in either case so no meaningful analysis could be conducted. Oxygen was liquified in stable state for the first time on March 29, 1877 by Polish scientists from Jagiellonian University, Zygmunt Wróblewski and Karol Olszewski.[60]
In 1891 Scottish chemist James Dewar was able to produce enough liquid oxygen to study.[10] The first commercially-viable process for producing liquid oxygen was independently developed in 1895 by German engineer Carl von Linde and British engineer William Hampson. Both men lowered the temperature of air until it liquefied and then distilled the component gases by boiling them off one at a time and capturing them.[61] 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.[61]
In 1923 the American scientist Robert H. Goddard became the first person to develop a rocket engine; the engine used gasoline for fuel and liquid oxygen as the oxidizer. Goddard successfully flew a small liquid-fueled rocket 56 m at 97 km/h on March 16, 1926 in Auburn, Massachusetts, USA.[61][62]
Industrial production
- See also: Oxygen evolution and fractional distillation
Two major methods are employed to produce the 100 million tonnes of O2 extracted from air for industrial uses annually.[54] The most common method is to fractionally-distill liquefied air into its various components, with nitrogen N2 distilling as a vapor while oxygen O2 is left as a liquid.[54]
The other major method of producing O2 gas involves passing a stream of clean, dry air through one bed of a pair of identical zeolite molecular sieves, which absorbs the nitrogen and delivers a gas stream that is 90% to 93% O2.[54] Simultaneously, nitrogen gas is released from the other nitrogen-saturated zeolite bed, by reducing the chamber operating pressure and diverting part of the oxygen gas from the producer bed through it, in the reverse direction of flow. After a set cycle time the operation of the two beds is interchanged, thereby allowing for a continuous supply of gaseous oxygen to be pumped through a pipeline. This is known as pressure swing adsorption. Oxygen gas is increasingly obtained by these non-cryogenic technologies (see also the related vacuum swing adsorption).[63]
Oxygen gas can also be produced through electrolysis of acidified water into molecular oxygen and hydrogen. DC electricity must be used: if AC is used, the gases in each limb consist of hydrogen and oxygen in the explosive ratio 2:1. Contrary to popular belief, the 2:1 ratio observed in the DC electrolysis of acidified water does not prove that the empirical formula of water is H2O unless certain assumptions are made about the molecular formulae of hydrogen and oxygen themselves.
A similar method is the electrocatalytic O2 evolution from oxides and oxoacids. Chemical catalysts can be used as well, such as in chemical oxygen generators or oxygen candles that are used as part of the life-support equipment on submarines, and are still part of standard equipment on commercial airliners in case of depressurization emergencies. Another air separation technology involves forcing air to dissolve through ceramic membranes based on zirconium dioxide by either high pressure or an electric current, to produce nearly pure O2 gas.[36]
In large quantities, the price of liquid oxygen in 2001 was approximately $0.21/kg.[64] Since the primary cost of production is the energy cost of liquefying the air, the production cost will change as energy cost varies.
For reasons of economy oxygen is often transported in bulk as a liquid in specially-insulated tankers, since one litre of liquefied oxygen is equivalent to 840 liters of gaseous oxygen at atmospheric pressure and 20 °C.[54] Such tankers are used to refill bulk liquid oxygen storage containers, which stand outside hospitals and other institutions with a need for large volumes of pure oxygen gas. Liquid oxygen is passed through heat exchangers, which convert the cryogenic liquid into gas before it enters the building. Oxygen is also stored and shipped in smaller cylinders containing the compressed gas; a form that is useful in certain portable medical applications and oxy-fuel welding and cutting.[54]
Applications
- See also: Breathing gas, Redox, and Combustion
Medical
Uptake of O2 from the air is the essential purpose of respiration, so oxygen supplementation is used in medicine. Oxygen therapy is used to treat emphysema, pneumonia, some heart disorders, and any disease that impairs the body's ability to take up and use gaseous oxygen.[65] Treatments are flexible enough to be used in hospitals, the patient's home, or increasingly by portable devices.