Oxygen

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This article is about the chemical element and its most stable form, O
2 or dioxygen. For other forms of this element, see Allotropes of oxygen. For other uses, see Oxygen (disambiguation) and O2 (disambiguation).

Oxygen   ₈O

Liquid oxygen, oxygen bubbles
Spectral lines of oxygen
General properties
Name, symbol	oxygen, O
Pronunciation	/ˈɒksɨdʒən/ OK-si-jən
Appearance	colorless gas; pale blue liquid
Oxygen in the periodic table
- ↑ O ↓ S nitrogen ← oxygen → fluorine
Atomic number	8
Standard atomic weight	15.999(4)
Element category	diatomic nonmetal
Group, period, block	group 16 (chalcogens), period 2, p-block
Electron configuration	[He] 2s2 2p4 per shell: 2, 6
Physical properties
Phase	gas
Melting point	54.36 K, ​−218.79 °C, ​−361.82 °F
Boiling point	90.188 K, ​−182.962 °C, ​−297.332 °F
Density	1.429 g/L (at 0 °C, 101.325 kPa)
Liquid density	at b.p.: 1.141 g·cm−3
Triple point	54.361 K, ​0.1463 kPa
Critical point	154.581 K, 5.043 MPa
Heat of fusion	(O2) 0.444 kJ·mol−1
Heat of vaporization	(O2) 6.82 kJ·mol−1
Molar heat capacity	(O2) 29.378 J·mol−1·K−1
Vapor pressure P (Pa) 1 10 100 1 k 10 k 100 k at T (K)       61 73 90
Atomic properties
Oxidation states	2, 1, −1, −2
Electronegativity	3.44 (Pauling scale)
Ionization energies	1st: 1313.9 kJ·mol−1 2nd: 3388.3 kJ·mol−1 3rd: 5300.5 kJ·mol−1 (more)
Covalent radius	66±2 pm
Van der Waals radius	152 pm
Miscellanea
Crystal structure	​cubic
Speed of sound	330 m·s−1 (gas, at 27 °C)
Thermal conductivity	26.58×10−3  W·m−1·K−1
Magnetic ordering	paramagnetic
CAS Number	7782-44-7
History
Discovery	Carl Wilhelm Scheele (1772)
Named by	Antoine Lavoisier (1777)
Most stable isotopes
Main article: Isotopes of oxygen
iso NA half-life DM DE (MeV) DP 16O 99.76% 16O is stable with 8 neutrons 17O 0.039% 17O is stable with 9 neutrons 18O 0.201% 18O is stable with 10 neutrons
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Blue white glow from an oxygen discharge tube.

Oxygen is a chemical element with symbol O and atomic number 8. It is a member of the chalcogen group on the periodic table and is a highly reactive nonmetallic element and oxidizing agent that readily forms compounds (notably oxides) with most elements.^[1] By mass, oxygen is the third-most abundant element in the universe, after hydrogen and helium.^[2] At STP, two atoms of the element bind to form dioxygen, a diatomic gas that is colorless, odorless, and tasteless, with the formula O
2.
Many major classes of organic molecules in living organisms, such as proteins, nucleic acids, carbohydrates, and fats, contain oxygen, as do the major inorganic compounds that are constituents of animal shells, teeth, and bone. Most of the mass of living organisms is oxygen as it is a part of water, the major constituent of lifeforms (for example, about two-thirds of human body mass). Elemental oxygen is produced by cyanobacteria, algae and plants, 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 O
2 began to accumulate in the atmosphere. Free elemental O
2 only began to accumulate in the atmosphere about 2.5 billion years ago during the Great Oxygenation Event, about a billion years after the first appearance of these organisms.^[3][4] Diatomic oxygen gas constitutes 20.8% of the volume of air.^[5] Oxygen is the most abundant element by mass in the Earth's crust as part of oxide compounds such as silicon dioxide, making up almost half of the crust's mass.^[6]
Oxygen is an important part of the atmosphere, and is necessary to sustain most terrestrial life as it is used in respiration. However, it is too chemically reactive to remain a free element in Earth's atmosphere without being continuously replenished by the photosynthetic action of living organisms, which use the energy of sunlight to produce elemental oxygen from water. Another form (allotrope) of oxygen, ozone (O
3), strongly absorbs UVB radiation and consequently the high-altitude ozone layer helps protect the biosphere from ultraviolet radiation, but is a pollutant near the surface where it is a by-product of smog. At even higher low earth orbit altitudes, atomic oxygen is a significant presence and a cause of erosion for spacecraft.^[7] Oxygen is produced industrially by fractional distillation of liquefied air, use of zeolites with pressure-cycling to concentrate oxygen from air, electrolysis of water and other means. Uses of elemental oxygen include the production of steel, plastics and textiles, brazing, welding and cutting of steels and other metals, rocket propellant, oxygen therapy and life support systems in aircraft, submarines, spaceflight and diving.

Oxygen was discovered independently 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 work was published first. The name oxygen was coined in 1777 by Antoine Lavoisier,^[8] whose experiments with oxygen helped to discredit the then-popular phlogiston theory of combustion and corrosion. Its name derives from the Greek roots ὀξύς oxys, "acid", literally "sharp", referring to the sour taste of acids and -γενής -genes, "producer", literally "begetter", because at the time of naming, it was mistakenly thought that all acids required oxygen in their composition.

Characteristics

Structure

Oxygen O₂ molecule.

At standard temperature and pressure, oxygen is a colorless, odorless gas with the molecular formula O
2, 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^[9] or as a combination of one two-electron bond and two three-electron bonds.^[10]

Triplet oxygen (not to be confused with ozone, O
3) is the ground state of the O
2 molecule.^[11] The electron configuration of the molecule has two unpaired electrons occupying two degenerate molecular orbitals.^[a] 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.^[11]

A trickle of liquid oxygen is deflected by a magnetic field, illustrating its paramagnetic property

In normal triplet form, O
2 molecules are paramagnetic. That is, 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 O
2 molecules.^[12] 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.^[13][b]

Singlet oxygen is a name given to several higher-energy species of molecular O
2 in which all the electron spins are paired. It is much more reactive towards common organic molecules than is molecular oxygen per se. In nature, singlet oxygen is commonly formed from water during photosynthesis, using the energy of sunlight.^[14] It is also produced in the troposphere by the photolysis 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 from singlet 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 found mostly in the stratosphere.

The common allotrope of elemental oxygen on Earth is called dioxygen, O
2. It has a bond length of 121 pm and a bond energy of 498 kJ·mol⁻¹.^[18] 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 O
2 are covered in the remainder of this article.
Trioxygen (O
3) is usually known as ozone and is a very reactive allotrope of oxygen that is damaging to lung tissue.^[19] Ozone is produced in the upper atmosphere when O
2 combines with atomic oxygen made by the splitting of O
2 by ultraviolet (UV) radiation.^[8] Since ozone absorbs strongly in the UV region of the spectrum, the ozone layer of the upper atmosphere functions as a protective radiation shield for the planet.^[8] Near the Earth's surface, however, it is a pollutant formed as a by-product of automobile exhaust.^[19] The metastable molecule tetraoxygen (O
4) 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 this phase, created by pressurizing O
2 to 20 GPa, is in fact a rhombohedral O
8 cluster.^[22] This cluster has the potential to be a much more powerful oxidizer than either O
2 or O
3 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

Oxygen is more soluble in water than nitrogen is. Water in equilibrium with air contains approximately 1 molecule of dissolved O
2 for every 2 molecules of N
2, 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 standard atmosphere (101.3 kPa) 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 O
2 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 O
2 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

 

Late in a massive star's life, ¹⁶O concentrates in the 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% natural abundance).^[31]

Most ¹⁶O is synthesized at the end of the helium fusion process in massive stars but some is made in the neon burning process.^{[32] 17}O 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 evolved, massive stars.^[32]

Fourteen radioisotopes have been characterized. The most stable are ¹⁵O with a half-life of 122.24 seconds and ¹⁴O with a half-life of 70.606 seconds.^[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 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]

Occurrence

 

Ten most common elements in the Milky Way Galaxy estimated spectroscopically^[36]

Z	Element	Mass fraction in parts per million
1	Hydrogen	739,000	71 × mass of oxygen (red bar)
2	Helium	240,000	23 × mass of oxygen (red bar)
8	Oxygen	10,400	10400
6	Carbon	4,600	4600
10	Neon	1,340	1340
26	Iron	1,090	1090
7	Nitrogen	960	960
14	Silicon	650	650
12	Magnesium	580	580
16	Sulfur	440	440

Oxygen is the most abundant chemical element by mass in the Earth's biosphere, air, sea and land. Oxygen is the third most abundant chemical element in the universe, after hydrogen and helium.^[2] About 0.9% of the Sun's mass is oxygen.^[5] Oxygen constitutes 49.2% of the Earth's crust by mass^[6] and is the major component of the world's oceans (88.8% by mass).^[5] Oxygen gas is the second most common component of the Earth's atmosphere, taking up 20.8% of its volume and 23.1% of its mass (some 10¹⁵ tonnes).^[5][37][c] 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% O
2 by volume) and Venus have far lower concentrations. However, the O
2 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.

Cold water holds more dissolved O
2.

Free oxygen also occurs in solution in the world's water bodies. The increased solubility of O
2 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.^[38] Water polluted with plant nutrients such as nitrates or phosphates may stimulate growth of algae by a process called eutrophication and the decay of these organisms and other biomaterials may reduce amounts of O
2 in eutrophic water bodies. Scientists assess this aspect of water quality by measuring the water's biochemical oxygen demand, or the amount of O
2 needed to restore it to a normal concentration.^[39]

Analysis

500 million years of climate change vs ¹⁸O

Paleoclimatologists measure the ratio of oxygen-18 and oxygen-16 in the shells and skeletons of marine organisms to determine what the climate was like millions of years ago (see oxygen isotope ratio cycle). Seawater molecules that contain the lighter isotope, oxygen-16, evaporate at a slightly faster rate than water molecules containing the 12% heavier oxygen-18; this disparity increases at lower temperatures.^[40] 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.^[40] Paleoclimatologists also directly measure this ratio in the water molecules of ice core samples that are up to several hundreds of thousands of years old.
Planetary geologists have measured different abundances of oxygen isotopes in samples from the Earth, the Moon, Mars, and meteorites, but were long unable to obtain reference values for the isotope ratios in the Sun, believed to be the same as those of the primordial solar nebula. However, analysis of a silicon wafer exposed to the solar wind in space and returned by the crashed Genesis spacecraft has shown that the Sun has a higher proportion of oxygen-16 than does the Earth. The measurement implies that an unknown process depleted oxygen-16 from the Sun's disk of protoplanetary material prior to the coalescence of dust grains that formed the Earth.^[41]

Oxygen presents two spectrophotometric absorption bands peaking at the wavelengths 687 and 760 nm. Some remote sensing scientists have proposed using the measurement of the radiance coming from vegetation canopies in those bands to characterize plant health status from a satellite platform.^[42] This approach exploits the fact that in those bands it is possible to discriminate the vegetation's reflectance from its fluorescence, which is much weaker. The measurement is technically difficult owing to the low signal-to-noise ratio and the physical structure of vegetation; but it has been proposed as a possible method of monitoring the carbon cycle from satellites on a global scale.

Biological role of O₂

Photosynthesis and respiration

Photosynthesis splits water to liberate O
2 and fixes CO
2 into sugar in what is called a Calvin cycle.

In nature, free oxygen is produced by the light-driven splitting of water during oxygenic photosynthesis. According to some estimates, 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.^[43] 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.^[44]

A simplified overall formula for photosynthesis is:^[45]

6 CO
2 + 6 H
2O + photons → C
6H
12O
6 + 6 O
2

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.^[d] 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.^[46] The O
2 remaining after oxidation of the water molecule is released into the atmosphere.^[e]

Molecular dioxygen, O
2, 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:

C
6H
12O
6 + 6 O
2 → 6 CO
2 + 6 H
2O + 2880 kJ·mol⁻¹

In vertebrates, O
2 diffuses through membranes in the lungs and into red blood cells. Hemoglobin binds O
2, changing its color from bluish red to bright red^[19] (CO
2 is released from another part of hemoglobin through the Bohr effect). Other animals use hemocyanin (molluscs and some arthropods) or hemerythrin (spiders and lobsters).^[37] A liter of blood can dissolve 200 cm³ of O
2.^[37]
Reactive oxygen species, such as superoxide ion (O−
2) and hydrogen peroxide (H
2O
2), are dangerous by-products of oxygen use in organisms.^[37] 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.^[46]

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

Content in body

Partial pressures of oxygen in the human body (PO₂)

Unit	Alveolar pulmonary gas pressures	Arterial blood oxygen	Venous blood gas
kPa	14.2	11[48]-13[48]	4.0[48]-5.3[48]
mmHg	107	75[49]-100[49]	30[50]-40[50]

The free oxygen partial pressure in the body of a living vertebrate organism is highest in the respiratory system, and decreases along any arterial system, peripheral tissues and venous system, respectively. Partial pressure is the pressure which oxygen would have if it alone occupied the volume.^[51]

Build-up in the atmosphere

 

O
2 build-up in Earth's atmosphere: 1) no O
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

Free oxygen gas was almost nonexistent in Earth's atmosphere before photosynthetic archaea and bacteria evolved, probably about 3.5 billion years ago. Free oxygen first appeared in significant quantities during the Paleoproterozoic eon (between 3.0 and 2.3 billion years ago).^[52] For the first billion years, any free oxygen produced by these organisms combined with dissolved iron in the oceans to form banded iron formations. When such oxygen sinks became saturated, free oxygen began to outgas from the oceans 3–2.7 billion years ago, reaching 10% of its present level around 1.7 billion years ago.^[52][53]

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 Great Oxygenation Event (oxygen catastrophe) about 2.4 billion years ago. However, cellular respiration using O
2 enables aerobic organisms to produce much more ATP than anaerobic organisms, helping the former to dominate Earth's biosphere.^[54] Cellular respiration of O
2 occurs in all eukaryotes, including all complex multicellular organisms such as plants and animals.

Since the beginning of the Cambrian period 540 million years ago, O
2 levels have fluctuated between 15% and 30% by volume.^[55] Towards the end of the Carboniferous period (about 300 million years ago) atmospheric O
2 levels reached a maximum of 35% by volume,^[55] which may have contributed to the large size of insects and amphibians at this time.^[56] 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.^[12] At the current rate of photosynthesis it would take about 2,000 years to regenerate the entire O
2 in the present atmosphere.^[57]

History

Early experiments

Philo's experiment inspired later investigators.

One of the first known experiments on the relationship between combustion and air was conducted by the 2nd century BCE Greek writer on mechanics, Philo of Byzantium. In his work Pneumatica, Philo observed that inverting a vessel over a burning candle and surrounding the vessel's neck with water resulted in some water rising into the neck.^[58] 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.^[59]

In the late 17th century, Robert Boyle proved that air is necessary for combustion. English chemist John Mayow (1641–1679) refined this work by showing that fire requires only a part of air that he called spiritus nitroaereus or just nitroaereus.^[60] 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.^[61] 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.^[60] 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.^[60] Accounts of these and other experiments and ideas were published in 1668 in his work Tractatus duo in the tract "De respiratione".^[61]

Phlogiston theory

 

Stahl helped develop and popularize the phlogiston theory.

Robert Hooke, Ole Borch, Mikhail Lomonosov, and Pierre Bayen (fr) all produced oxygen in experiments in the 17th and the 18th century but none of them recognized it as a chemical 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,^[62] 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.^[59]

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.^[59] 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

Carl Wilhelm Scheele beat Priestley to the discovery but published afterwards.

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.^[5][59] Scheele called the gas "fire air" because it was the only known supporter of combustion, and wrote an account of this discovery in a manuscript he titled Treatise on Air and Fire, which he sent to his publisher in 1775. However, that document was not published until 1777.^[63]

Joseph Priestley is usually given priority in the discovery.

In the meantime, on August 1, 1774, an experiment conducted by the British clergyman Joseph Priestley focused sunlight on mercuric oxide (HgO) inside a glass tube, which liberated a gas he named "dephlogisticated air".^[5] 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.^[59][64] 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).^[63]

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.^[5] 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.

Antoine Lavoisier discredited the Phlogiston theory.

In one experiment, Lavoisier observed that there was no overall increase in weight when tin and air were heated in a closed container.^[5] 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.^[5] 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.^[5]

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 he mistakenly believed that oxygen was a constituent of all acids.^[8] Chemists (notably Sir Humphry Davy in 1812) eventually determined that Lavoisier was wrong in this regard (it is in fact hydrogen that forms the basis for acid chemistry), but by that time it was too late; the name had taken.
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.^[63]

Later history

Robert H. Goddard and a liquid oxygen-gasoline rocket

John Dalton's original atomic hypothesis assumed 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, giving the atomic mass of oxygen as 8 times that of hydrogen, instead of the modern value of about 16.^[65] 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.^[66][g]

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.^[67] Just two days later, French physicist Louis Paul Cailletet announced his own method of liquefying molecular oxygen.^[67]
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, 1883 by Polish scientists from Jagiellonian University, Zygmunt Wróblewski and Karol Olszewski.^[68]

In 1891 Scottish chemist James Dewar was able to produce enough liquid oxygen to study.^[12] 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.^[69] Later, in 1901, oxyacetylene welding was demonstrated for the first time by burning a mixture of acetylene and compressed O
2. This method of welding and cutting metal later became common.^[69]

In 1923 the American scientist Robert H. Goddard became the first person to develop a rocket engine that burned liquid fuel; 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, US.^[69][70]