Oxygen - WikiHQ

Oxygen is a chemical element; it has the symbolO and its atomic number is8. It is a member of the chalcogen group in the periodic table. It is highly reactive, a nonmetal, and a potent oxidizing agent that readily forms oxides with most elements as well as with other compounds. Oxygen is the most abundant element in Earth's crust, making up almost half of the Earth's crust in the form of various oxides such as water, carbon dioxide, iron oxides, and silicates. It is also the third-most abundant element in the universe after hydrogen and helium.

At standard temperature and pressure, two oxygen atoms will bind covalently to form dioxygen, a colorless and odorless diatomic gas with the chemical formula . Dioxygen gas currently constitutes approximately 20.95% molar fraction of the Earth's atmosphere, though this has changed considerably over long periods of time in Earth's history. The much rarer allotrope of oxygen, ozone (), strongly absorbs the UVB and UVC wavelengths and forms a protective ozone layer at the lower stratosphere, which shields the biosphere from ionizing ultraviolet radiation. However, ozone present at the surface is a corrosive byproduct of smog and thus an air pollutant.

All eukaryotic organisms, including plants, animals, fungi, algae, and most protists, need oxygen for cellular respiration, a process that extracts chemical energy by the reaction of oxygen with organic molecules derived from food and releases carbon dioxide as a waste product.

Many major classes of organic molecules in living organisms contain oxygen atoms, such as proteins, nucleic acids, carbohydrates, and fats, as do the major constituent inorganic compounds of animal shells, teeth, and bone. Most of the mass of living organisms is oxygen as a component of water, the major constituent of lifeforms. Oxygen in Earth's atmosphere is produced by biotic photosynthesis, in which photon energy in sunlight is captured by chlorophyll to split water molecules and then react with carbon dioxide to produce carbohydrates, with oxygen released as a byproduct. Oxygen is too chemically reactive to remain a free element in air without being continuously replenished by the photosynthetic activities of autotrophs such as cyanobacteria, chloroplast-bearing algae, and land plants.

Oxygen was isolated by Michael Sendivogius before 1604, but it is commonly believed that the element was discovered independently by Carl Wilhelm Scheele, in Uppsala, in 1773 or earlier, and Joseph Priestley in Wiltshire, 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.

Common industrial uses of oxygen include 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.

History of study

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.

Early experiments

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 ', Philo observed that inverting a vessel over a burning candle and surrounding the vessel's neck with water resulted in some water rising into the neck. Philo incorrectly surmised that parts of the air in the vessel were converted into the classical element fire and thus were able to escape through pores in the glass. Many centuries later Ibn 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.

Leonardo da Vinci observed that a portion of air is consumed during combustion and respiration.

Polish alchemist, philosopher, and physician Michael Sendivogius (Michał Sędziwój), writing in 1604, described a substance contained in air, referring to it as ('food of life'); During his experiments, performed between 1598 and 1604, Sendivogius properly recognized that the substance is equivalent to the gaseous byproduct released by the thermal decomposition of potassium nitrate. However, this important connection was not understood by contemporary scientists like Robert Boyle. In one experiment, he found that placing either a mouse or a lit candle in a closed container over water caused the water to rise and replace one-fourteenth of the air's volume before extinguishing the subjects. From this, he surmised that is consumed in both respiration and combustion.

Phlogiston theory

Robert Hooke, Ole Borch, Mikhail Lomonosov, and Pierre Bayen all produced oxygen in experiments in the 17th and the 18th century but none of them recognized it as a chemical element. 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, 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. Scheele had produced oxygen gas by heating mercuric oxide (HgO) and various nitrates in 1771–1772.

In the meantime, on August 1, 1774, an experiment conducted by the British clergyman Joseph Priestley focused sunlight on mercuric oxide contained in a glass tube, which liberated a gas he named "dephlogisticated air". 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, Priestley 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."

The French chemist Antoine 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). Chemists (such as Sir Humphry 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. 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 diatomic elemental molecules in those gases.

In 1879 the French brothers Quentin and Arthur

Brin discovered a commercially viable reaction to create oxygen. They realized that the known reversible reaction

(s) + (g) ↔ (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, at which point the more economical fractional distillation began to be used.

thumb|left|upright|[[Robert H. Goddard and a liquid-oxygen–gasoline rocket|alt=A metal frame structure stands on the snow near a tree. A middle-aged man wearing a coat, boots, leather gloves and a cap stands by the structure and holds it with his right hand.]]

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. On December 22, 1877, he sent a telegram to the French Academy of Sciences in Paris announcing his discovery of liquid oxygen. Just two days later, French physicist Louis Paul Cailletet announced his own method of liquefying molecular oxygen.

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

Characteristics

Properties and molecular structure

[[File:Oxygen molecule orbitals diagram-en.svg|thumb|left|upright=1.2|Orbital diagram, after Barrett (2002),

As dioxygen, two oxygen atoms are chemically bound to each other. The bond can be variously described based on level of theory, but is reasonably and simply described as a covalent double bond that results from the filling of molecular orbitals formed from the atomic orbitals of the individual oxygen atoms, the filling of which results in a bond order of two. More specifically, the double bond is the result of sequential, low-to-high energy, or Aufbau, 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.

This combination of cancellations and σ and π overlaps results in dioxygen's double-bond character and reactivity, as well as the presence of a triplet electronic ground state. An electron 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 a spin triplet state. Hence, the ground state of the molecule is referred to as triplet oxygen. The highest-energy, partially filled orbitals are antibonding, 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.

thumb|left|upright|Liquid oxygen, temporarily suspended in a magnet owing to its paramagnetism

In the triplet form, molecules are paramagnetic. That is, they impart magnetic character to oxygen when it is in the presence of a magnetic field, because of the spin of the unpaired electrons in the molecule and the negative exchange energy between neighboring molecules. Oxygen's paramagnetism can be used in paramagnetic oxygen gas analysers that determine gaseous oxygen concentration, especially in industrial process control and medicine.

Singlet oxygen is a name given to several higher-energy species of molecular in which all the electron spins are paired. It is much more reactive with common organic molecules than normal (triplet) molecular oxygen. In nature, singlet oxygen is commonly formed from water during photosynthesis, using the energy of sunlight. It is also produced in the troposphere by the photolysis of ozone by light of short wavelength and by the immune system as a source of active oxygen. Carotenoids in photosynthetic organisms (and possibly animals) play a major role in absorbing energy from singlet oxygen and converting it to the unexcited ground state before it can cause harm to tissues.

Allotropes

thumb|right|upright=0.9|[[Space-filling model representation of dioxygen (O2) molecule]]

The common allotrope of elemental oxygen on Earth is called dioxygen, , the allotrope that is major part of the Earth's atmospheric oxygen (see occurrence). O2 has a bond length of 121 pm and a bond energy of 498 kJ/mol.

Trioxygen () is usually known as ozone and is a very reactive allotrope of oxygen that is damaging to lung tissue. Ozone is produced in the upper atmosphere when combines with atomic oxygen made by the splitting of by ultraviolet (UV) radiation.

The metastable molecule tetraoxygen () was discovered in 2001 and was assumed to exist in one of the six phases of solid oxygen. In 2006, this phase, created by pressurizing to 20 GPa, was shown to form a rhombohedral cluster. This cluster has the potential to be a much more powerful oxidizer than either or and may therefore be used in rocket fuel.

Physical properties

thumb|Liquid oxygen boiling (O2)|alt=A transparent beaker containing a light blue fluid with gas bubbles.

Oxygen dissolves more readily in water than nitrogen does. Water in equilibrium with air contains approximately 1 molecule of dissolved for every 2 molecules of (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 () dissolves at than at ().

At and in air, freshwater can dissolve about 6.04 milliliters (mL) of oxygen per liter, while seawater contains about 4.95 mL per liter.

At the solubility increases to 9.0 mL (50% more than at ) per liter for freshwater and 7.2 mL (45% more) per liter for sea water.

{| class="wikitable" style="float:left; margin-right:2em"

|+Oxygen gas dissolved in water at sea-level (milliliters per liter)

|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). Both liquid and solid are clear substances with a light sky-blue color caused by absorption in the red (in contrast with the blue color of the sky, which is due to Rayleigh scattering of blue light). High-purity liquid is usually obtained by the fractional distillation of liquefied air. Liquid oxygen may also be condensed from air using liquid nitrogen as a coolant. Liquid oxygen is a highly reactive substance and must be segregated from combustible materials. The absorption in the Herzberg continuum and Schumann–Runge bands in the ultraviolet produces atomic oxygen that is important in the chemistry of the middle atmosphere. Excited-state singlet molecular oxygen is responsible for red chemiluminescence in solution.

Table of thermal and physical properties of oxygen (O2) at atmospheric pressure:

{|class="wikitable mw-collapsible mw-collapsed"

|Temperature (K)

|Density (kg/m3)

|Specific heat (kJ/(kg·K))

|Dynamic viscosity (kg/(m·s))

|Kinematic viscosity (m2/s)

|Thermal conductivity (W/(m·K))

|Thermal diffusivity (m2/s)

|Prandtl Number

|100

|3.945

|0.962

|7.64E-06

|1.94E-06

|0.00925

|2.44E-06

|0.796

|150

|2.585

|0.921

|1.15E-05

|4.44E-06

|0.0138

|5.80E-06

|0.766

|200

|1.93

|0.915

|1.48E-05

|7.64E-06

|0.0183

|1.04E-05

|0.737

|250

|1.542

|0.915

|1.79E-05

|1.16E-05

|0.0226

|1.60E-05

|0.723

|300

|1.284

|0.92

|2.07E-05

|1.61E-05

|0.0268

|2.27E-05

|0.711

|350

|1.1

|0.929

|2.34E-05

|2.12E-05

|0.0296

|2.90E-05

|0.733

|400

|0.962

|1.0408

|2.58E-05

|2.68E-05

|0.033

|3.64E-05

|0.737

|450

|0.8554

|0.956

|2.81E-05

|3.29E-05

|0.0363

|4.44E-05

|0.741

|500

|0.7698

|0.972

|3.03E-05

|3.94E-05

|0.0412

|5.51E-05

|0.716

|550

|0.6998

|0.988

|3.24E-05

|4.63E-05

|0.0441

|6.38E-05

|0.726

|600

|0.6414

|1.003

|3.44E-05

|5.36E-05

|0.0473

|7.35E-05

|0.729

|700

|0.5498

|1.031

|3.81E-05

|6.93E-05

|0.0528

|9.31E-05

|0.744

|800

|0.481

|1.054

|4.15E-05

|8.63E-05

|0.0589

|1.16E-04

|0.743

|900

|0.4275

|1.074

|4.47E-05

|1.05E-04

|0.0649

|1.41E-04

|0.74

|1000

|0.3848

|1.09

|4.77E-05

|1.24E-04

|0.071

|1.69E-04

|0.733

|1100

|0.3498

|1.103

|5.06E-05

|1.45E-04

|0.0758

|1.96E-04

|0.736

|1200

|0.3206

|1.0408

|5.33E-05

|1.661E-04

|0.0819

|2.29E-04

|0.725

|1300

|0.296

|1.125

|5.88E-05

|1.99E-04

|0.0871

|2.62E-04

|0.721

Isotopes and stellar origin

thumb|Late in a massive star's life, 16O concentrates in the O-shell, 17O in the H-shell and 18O in the He-shell.|alt=A concentric-sphere diagram, showing, from the core to the outer shell, iron, silicon, oxygen, neon, carbon, helium and hydrogen layers.

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

16O is one of the dominant fusion products in massive stars. It is synthesized at the end of the triple-alpha process with some synthesis in the neon burning process. Both 17O and 18O require seed nuclei. 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. The most stable are 15O with a half-life of 122.24 seconds and 14O with a half-life of 70.606 seconds.

Occurrence

{| class="wikitable sortable" style="float:left; margin-right: 20px"

|+Ten most common elements in the Milky Way Galaxy estimated spectroscopically

!Z !! Element !! colspan="2"|Mass fraction in parts per million

| 1 || Hydrogen || align="right"|

| 2 || Helium || align="right"|

| 8 || Oxygen || align="right"|

| 6 || Carbon || align="right"|

| 10 || Neon || align="right"|

| 26 || Iron || align="right"|

| 7 || Nitrogen || align="right"|

| 14 || Silicon || align="right"|

| 12 || Magnesium || align="right"|

| 16 || Sulfur || align="right"|

Oxygen is the third most abundant chemical element in the universe, after hydrogen and helium. About 0.9% of the Sun's mass is oxygen. It is also the major component of the world's oceans (88.8% by mass).

Earth is unusual among the planets of the Solar System in having such a high concentration of oxygen gas in its atmosphere. Mars (with 0.1% by volume) and Venus have much less. The surrounding those planets is produced solely by the action of ultraviolet radiation on 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, decay, and combustion remove it from the atmosphere. In the present equilibrium, production and consumption occur at the same rate.

thumb|right|Cold water holds more dissolved .|alt=World map showing that the sea-surface oxygen is depleted around the equator and increases towards the poles.

Free oxygen also occurs in solution in the world's water bodies. The increased solubility of at lower temperatures (see Physical properties) has important implications for ocean life, as polar oceans support a much higher density of life due to their higher oxygen content. Scientists assess this aspect of water quality by measuring the water's biochemical oxygen demand, or the amount of needed to restore it to a normal concentration.

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 support macrofauna.

Analysis

Paleoclimatologists measure the ratio of oxygen-18 and oxygen-16 in the shells and skeletons of marine organisms to determine the climate 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, and this disparity increases at lower temperatures. 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.

Planetary geologists have measured the relative quantities 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. Analysis of a silicon wafer exposed to the solar wind in space and returned by the crashed Genesis spacecraft has shown that the Sun has a higher proportion of oxygen-16 than does the Earth. The measurement implies that an unknown process depleted oxygen-16 from the Sun's disk of protoplanetary material prior to the coalescence of dust grains that formed the Earth.

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

Biological production and role of O2

Photosynthesis and respiration

thumb|Photosynthesis splits water to liberate and fixes into sugar in what is called a [[Calvin cycle.|alt=A diagram of photosynthesis processes, including income of water and carbon dioxide, illumination and release of oxygen. Reactions produce ATP and NADPH in a Calvin cycle with a sugar as a by product.]]

In nature, free oxygen is produced as a byproduct of light-driven splitting of water during chlorophyllic photosynthesis. According to some estimates, marine photoautotrophs such as red/green algae and cyanobacteria provide about 70% of the free oxygen produced on Earth, and the rest is produced in terrestrial environments by plants. 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.

A simplified overall formula for photosynthesis is

: 6 + 6 + photons → + 6

or simply

: carbon dioxide + water + sunlight → glucose + dioxygen

Photolytic oxygen evolution occurs in the thylakoid membranes of photosynthetic organisms and requires the energy of four photons. Many steps are involved, but the result is the formation of a proton gradient across the thylakoid membrane, which is used to synthesize adenosine triphosphate (ATP) via photophosphorylation. The remaining (after production of the water molecule) is released into the atmosphere.

Oxygen is used in mitochondria of eukaryotes to generate ATP during oxidative phosphorylation. The reaction for aerobic respiration is essentially the reverse of photosynthesis and is simplified as

: + 6 → 6 + 6 + 2880 kJ/mol

In aquatic animals, dissolved oxygen in water is absorbed by gills, through the skin, or via the gut; in terrestrial animals such as tetrapods, oxygen in air is actively taken into the body via lungs, where gas exchange takes place to diffuse oxygen into the blood and carbon dioxide out, and the body's circulatory system then transports the oxygen to other tissues where cellular respiration takes place. However, in insects, the most successful and biodiverse terrestrial clade, oxygen is directly conducted to the internal tissues via a deep network of airways. Hemoglobin in red blood cells binds , changing color from bluish red to bright red

Reactive oxygen species, such as superoxide ion () and hydrogen peroxide (), are reactive by-products of oxygen use in organisms.

An adult human at rest inhales 1.8 to 2.4 grams of oxygen per minute. This amounts to more than 6 billion tonnes of oxygen inhaled by humanity per year.

Living organisms

{|class="wikitable" style="float:right; margin-left:25px"

|+Partial pressures of oxygen in the human body (PO2)

! Unit !! Alveolar pulmonary gas pressures !! Arterial blood oxygen !! Venous blood gas

| kPa || 14.2 || 11-13 || 4.0-5.3

| mmHg || 107 || 75-100

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 that oxygen would have if it alone occupied the volume.

Build-up in the atmosphere

thumb|left|upright=1.35| build-up in Earth's atmosphere: 1) no produced; 2) produced, but absorbed in oceans & seabed rock; 3) starts to gas out of the oceans, but is absorbed by land surfaces and formation of ozone layer; 4–5) sinks filled and the gas accumulates|alt=A graph showing time evolution of oxygen pressure on Earth; the pressure increases from zero to 0.2 atmospheres.

Before photosynthesis evolved, Earth's atmosphere had little free diatomic elemental oxygen (O2). The concentrations of O2 attained were less than 10% of today's and probably fluctuated greatly. Around 500Mya a second event known as the Neoproterozoic Oxygenation Event lead to oxygen levels similar or even higher than the present. 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. The increase in oxygen concentrations had wide-ranging and significant impacts on Earth's geochemistry and biosphere. However, detailed connections between oxygen and evolution remain elusive.

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.

Extraterrestrial free oxygen

Small amounts of oxygen have been detected in Europa's and Ganymede's thin oxygen atmospheres, particularly around the polars. This oxygen is believe to result from photodissociation of water.

In the field of astrobiology and in the search for extraterrestrial life, oxygen is considered the strongest biosignature, or sign of biological activity. Oxygen meets the three criteria for a biosignature. Its presence is a reliable indicator of life based on the observation that almost all of Earth's oxygen has biological origin. Second, oxygen persists in the atmosphere over long periods of time, on the order of a billion years. Third, oxygen mixes well in the atmosphere and has strong distinctive absorption lines which can be detected by remote telescopes. One potential issue is that oxygen may be produced abiotically on celestial bodies with processes and conditions (such as a peculiar hydrosphere) which allow build up of free oxygen.

Industrial production

thumb|upright|[[Hofmann voltameter|Hofmann electrolysis apparatus used in electrolysis of water|alt=A drawing of three vertical pipes connected at the bottom and filled with oxygen (left pipe), water (middle) and hydrogen (right). Anode and cathode electrodes are inserted into the left and right pipes and externally connected to a battery.]]

Every year, one hundred million tonnes of are extracted from air for industrial uses.

alt=An experiment setup with test tubes to prepare oxygen|left|thumb|280x280px|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.

Oxygen gas can also be produced through electrolysis of water into molecular oxygen and hydrogen. DC electricity must be used: if AC is used, the gases in each limb consist of hydrogen and oxygen in the explosive ratio 2:1. A similar method is the electrocatalytic evolution from oxides and oxoacids. Chemical catalysts can be used as well, such as in chemical oxygen generators or oxygen candles that are used as part of the life-support equipment on submarines, and are still part of standard equipment on commercial airliners in case of depressurization emergencies. Another air separation method is forcing air to dissolve through ceramic membranes based on zirconium dioxide by either high pressure or an electric current to produce nearly pure gas.

Treatments are flexible enough to be used in hospitals, the patient's home, or increasingly by portable devices. Oxygen tents were once commonly used in oxygen supplementation, but have since been replaced mostly by the use of oxygen masks or nasal cannulas.

Hyperbaric (high-pressure) medicine uses special oxygen chambers to increase the partial pressure of around the patient and, when needed, the medical staff. Carbon monoxide poisoning, gas gangrene, and decompression sickness (the 'bends') are sometimes addressed with this therapy. Increased concentration in the lungs helps to displace carbon monoxide from the heme group of hemoglobin. Oxygen gas is poisonous to the anaerobic bacteria that cause gas gangrene, so increasing its partial pressure helps kill the bacteria and alleviates gas gangrene. 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 of 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. 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.

Life support and recreational use

thumb|Low-pressure pure is used in [[space suits.]]

In modern space suits, which surround their occupant's body, oxygen gas is used as a low-pressure breathing gas. These devices use nearly pure oxygen at about one-third normal pressure, resulting in a normal blood partial pressure of . This trade-off of higher oxygen concentration for lower pressure is needed to maintain suit flexibility.

Scuba and surface-supplied underwater divers and submarines also rely on artificially delivered . Submarines, submersibles, and atmospheric 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 pure use in diving at pressures higher than atmospheric is usually limited to rebreathers, or decompression at relatively shallow depths (~6 meters depth, or less), or medical 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 of with other gases, such as nitrogen or helium, to prevent oxygen toxicity. Pressurized commercial airplanes have an emergency supply of automatically supplied to the passengers in case of cabin depressurization. Sudden cabin pressure loss activates chemical oxygen generators above each seat, causing oxygen masks to drop. Pulling on the masks "to start the flow of oxygen" as cabin safety instructions dictate, forces iron filings into the sodium chlorate inside the canister. Professional athletes, especially in American football, sometimes go off-field between plays to don oxygen masks to boost performance. The pharmacological effect is doubted; a placebo effect is a more likely explanation.

Other recreational uses that do not involve breathing include pyrotechnic applications, such as George Goble's five-second ignition of barbecue grills.

Industrial

thumb|Most commercially produced is used to [[smelting|smelt and/or decarburize iron.|alt=An elderly worker in a helmet is facing his side to the viewer in an industrial hall. The hall is dark but is illuminated yellow glowing splashes of a melted substance.]]

Smelting of iron ore into steel consumes 55% of commercially produced oxygen.

Compounds

thumb|upright|[[Water () is the most familiar oxygen compound.|alt=Water flowing from a bottle into a glass.]]

The oxidation state of oxygen is −2 in almost all known compounds of oxygen. The oxidation state −1 is found in a few compounds such as peroxides.

Oxides and other inorganic compounds

Water () is an oxide of hydrogen and the most familiar oxygen compound. Hydrogen atoms are covalently bonded to oxygen in a water molecule but also have an additional attraction (about 23.3 kJ/mol per hydrogen atom) to an adjacent oxygen atom in a separate molecule. These hydrogen bonds between water molecules hold them approximately 15% closer than what would be expected in a simple liquid with just van der Waals forces.

thumb|left|Oxides, such as [[iron oxide or rust, form when oxygen combines with other elements.|alt=A rusty piece of a bolt.]]

Due to its electronegativity, oxygen forms chemical bonds with almost all other elements to give corresponding oxides. The surface of most metals, such as aluminium and titanium, are oxidized in the presence of air and become coated with a thin film of oxide that passivates the metal and slows further corrosion. Many oxides of the transition metals are non-stoichiometric compounds, with slightly less metal than the chemical formula would show. For example, the mineral FeO (wüstite) is written as <math chem>\ce{Fe}_{1-x}\ce{O}</math>, where x is usually around 0.05.

Oxygen is present in the atmosphere in trace quantities in the form of carbon dioxide (). The Earth's crustal rock is composed in large part of oxides of silicon (silica , as found in granite and quartz), aluminium (aluminium oxide , in bauxite and corundum), iron (iron(III) oxide , in hematite and rust), and calcium carbonate (in limestone). The rest of the Earth's crust is also made of oxygen compounds, in particular various complex silicates (in silicate minerals). The Earth's mantle, of much larger mass than the crust, is largely composed of silicates of magnesium and iron.

Water-soluble silicates in the form of , , and are used as detergents and adhesives.

Oxygen also acts as a ligand for transition metals, forming transition metal dioxygen complexes, which feature metal–. This class of compounds includes the heme proteins hemoglobin and myoglobin. An exotic and unusual reaction occurs with platinum hexafluoride|, which oxidizes oxygen to give , dioxygenyl hexafluoroplatinate.

Organic compounds

thumb|[[Acetone is an important feeder material in the chemical industry.

|alt=A ball structure of a molecule. Its backbone is a zig-zag chain of three carbon atoms connected in the center to an oxygen atom and on the end to 6 hydrogens.]]

Among the most important classes of organic compounds that contain oxygen are (where "R" is an organic group): alcohols (); ethers (); ketones (); aldehydes (); carboxylic acids (); esters (); acid anhydrides (); and amides (). There are many important organic solvents that contain oxygen, including: acetone, methanol, ethanol, isopropanol, furan, THF, diethyl ether, dioxane, ethyl acetate, DMF, DMSO, acetic acid, and formic acid. Acetone () and phenol () are used as feeder materials in the synthesis of many different substances. Other important organic compounds that contain oxygen are: glycerol, formaldehyde, glutaraldehyde, citric acid, acetic anhydride, and acetamide. Epoxides are ethers in which the oxygen atom is part of a ring of three atoms. The element is similarly found in almost all biomolecules that are important to (or generated by) life.

Oxygen reacts spontaneously with many organic compounds at or below room temperature in a process called autoxidation. Most of the organic compounds that contain oxygen are not made by direct action of . Organic compounds important in industry and commerce that are made by direct oxidation of a precursor include ethylene oxide and peracetic acid.

Toxicity

thumb|left|upright=1.35|Main symptoms of oxygen toxicity|alt=A diagram showing a male torso and listing symptoms of oxygen toxicity: Eyes – visual field loss, nearsightedness, cataract formation, bleeding, fibrosis; Head – seizures; Muscles – twitching; Respiratory system – jerky breathing, irritation, coughing, pain, shortness of breath, tracheobronchitis, acute respiratory distress syndrome.

Oxygen gas () can be toxic at elevated partial pressures, leading to convulsions and other health problems. 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-level partial pressure of about 21 kPa. This is not a problem except for patients on mechanical ventilators, since gas supplied through oxygen masks in medical applications is typically composed of only 30–50% by volume (about 30 kPa at standard pressure). In the case of spacesuits, the partial pressure in the breathing gas is, in general, about 30 kPa (1.4 times normal), and the resulting partial pressure in the astronaut's arterial blood is only marginally more than normal sea-level partial pressure.

Oxygen toxicity to the lungs and central nervous system can also occur in deep scuba diving and surface-supplied diving. Exposure to an 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% at or more of depth; the same thing can occur by breathing 100% at only .

Combustion and other hazards

thumb|right|The interior of the [[Apollo 1 Command Module. Pure at higher than normal pressure and a spark led to a fire and the loss of the Apollo 1 crew.|alt=The inside of a small spaceship, charred and apparently destroyed.]]

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. Concentrated will allow combustion to proceed rapidly and energetically.

Liquid oxygen spills, if allowed to soak into organic matter such as wood, petrochemicals, and asphalt, can cause these materials to detonate unpredictably on subsequent mechanical impact.