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Oxygen is a chemical element; it has the symbol O and its atomic number is 8.
Oxygen is a member of the chalcogen group in the periodic table.
Oxygen is highly reactive, a nonmetal, and a potent oxidizing agent that readily forms oxides with most elements as well as with other compounds.
CAS: 7782-44-7
MF: O2
MW: 32
EINECS: 231-956-9
Synonyms
hyperoxia;Liquid-oxygen-;AUTOMOTIVE STANDARD R99-5;AUTOMOTIVE STANDARD R99-6;HYDROGEN AND OXYGEN;OXYGEN, PRESSURE TIN WITH 1 L;OXYGEN, 99.6+%;OXYGEN-16O2 (GAS) (18O-DEPLETED) 99.98%
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.
Oxygen 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 O2.
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 (O3), 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.
Oxygen was discovered by Joseph Priestley in 1774 by the thermal decomposition of mercuric oxide and published his findings the same year, three years before Scheele published.
In 1775–80, French chemist Antoine-Laurent Lavoisier, with remarkable insight, interpreted the role of oxygen in respiration as well as combustion, discarding the phlogiston theory, which had been accepted up to that time; he noted its tendency to form acids by combining with many different substances and accordingly named the element oxygen (oxygène) from the Greek words for“acid former.”
Oxygen is a very prevalent and important element and is necessary for sustaining life on this planet.
This element is the third most abundant in mass behind helium and hydrogen in the universe.
The diatomic form (O2) is the most common pure form.
With a boiling point at -183 °C, O2 exists as a colorless and odorless gas at standard temperature and pressure.
In the process of cellular respiration, the highly reactive O2 is used as the oxidant in breaking down food molecules to produce energy.
In turn, photosynthetic organisms generate O2 by using energy from the sun and water.
Other allotropes of pure oxygen exist, including the trioxygen (O3) form known as ozone, as well as other, less common allotropes of oxygen such as O4 and O8.
These oxygen allotropes are formed under high pressure and low temperatures and are solid.
Symbol O.
A colourlessodourless gaseous element belongingto group 16 (formerly VIB) of the periodictable; a.n. 8; r.a.m. 15.9994; d.1.429 g dm–3; m.p. –218.4°C; b.p.–183°C.
Oxygen is the most abundant elementin the earth’s crust (49.2% byweight) and is present in the atmosphere(28% by volume).
Atmosphericoxygen is of vital importance for allorganisms that carry out aerobic respiration.
For industrial purposes Oxygen isobtained by fractional distillation of liquid air.
Oxygen is used in metallurgicalprocesses, in high-temperatureflames (e.g. for welding), and inbreathing apparatus.
The commonform is diatomic (dioxygen, O2);there is also a reactive allotropeozone (O3).
Chemically, oxygen reactswith most other elements formingoxides.
The element wasdiscovered by Joseph Priestley in1774.
Oxygen Chemical Properties
Melting point: −218 °C(lit.)
Boiling point: −183 °C(lit.)
density: 1.429（0℃）
vapor density: 1.11 (vs air)
vapor pressure: >760 mmHg at 20 °C
storage temp.: -20°C
solubility: At 20 °C and at a pressure of 101 kPa, 1 volume dissolves in about 32 volumes of water.
form: colorless gas
color: Colorless gas, liquid, or hexagonal crystals
Odor: Odorless gas
biological source: rabbit
Water Solubility: one vol gas dissolves in 32 volumes H2O (20°C), in 7 volumes alcohol (20°C); soluble other organic liq, usually higher solubility than in H2O [MER06]
Thermal Conductivity: 0.02658 W/(m·K)
Merck: 13,7033
Henry's Law Constant: 1.3×10-5 mol/(m3Pa) at 25℃, Burkholder et al. (2019)
Dielectric constant: 1.5（-193℃）
Stability: Stable. Vigorously supports combustion. Incompatible with phosphorus, organic materials, many powdered metals.
Cosmetics Ingredients Functions: SKIN CONDITIONING
InChI: 1S/O2/c1-2/i1+0,2+0
InChIKey: MYMOFIZGZYHOMD-ZCWHFVSRSA-N
Surface tension: 15.85 mN/m at 80.0K
CAS DataBase Reference: 7782-44-7(CAS DataBase Reference)
NIST Chemistry Reference: Oxygen(7782-44-7)
EPA Substance Registry System: Oxygen (7782-44-7)
Oxygen, O2, is a colorless, tasteless, gaseous element essential to almost all forms of life.
Oxygen promotes respiration and combustion.
Oxygen comprises 20% of the earth's atmosphere and is the most abundant element in seawater and in the earth's crust.
Oxygen is slightly soluble in water and alcohol, but combines readily with most other elements to form oxides.
The electrolysis of water produces both oxygen and hydrogen.
Characteristics
At standard temperature and pressure, oxygen is a colorless, odorless, and tasteless gas with the molecular formula O2, referred to as dioxygen.
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.
Oxygen is, without a doubt, the most essential element on Earth.
Oxygen is required to supportall plant and animal life, and it forms more compounds with other elements than any otherelement.
Oxygen is soluble in both water and alcohol.
Contrary to what many people believe, oxygenis NOT combustible (it will not burn), but rather it actively supports the combustion ofmany other substances.
After all, if oxygen burned, every time a fire was lit, all the O2 in theatmosphere would be consumed!
Burning is a form of oxidation wherein oxygen chemically combines with a substance rapidlyenough to produce adequate heat to cause fire and light, or to maintain a fire once started.
The oxidation of iron is called rusting.
Rusting in an example of “slow oxidation,” which isthe reaction of O2 with Fe to form Fe2O3 or Fe3O4.
This chemical reaction is so slow that theheat it produces is dissipated; thus, there is no fire.
Recently a new allotrope of oxygen was discovered.
When O2 is subjected to great pressure, it is converted intoO4, which is a deep red solid that is a much more powerful oxidizer thanthe other forms of oxygen.
Physical properties
Oxygen dissolves more readily in water than nitrogen does.
Water in equilibrium with air contains approximately 1 molecule of dissolved O
2 for every 2 molecules of N
2 (1:2), compared with an atmospheric ratio of approximately 1:4.
The solubility of oxygen in water is temperature-dependent, and about twice as much (14.6 mg/L) dissolves at 0 °C (32 °F) than at 20 °C (68 °F) (7.6 mg/L).
At 25 °C (77 °F) and 1 standard atmosphere (101.325 kPa) in air, freshwater can dissolve about 6.04 milliliters (mL) of oxygen per liter, while seawater contains about 4.95 mL per liter.
At 5 °C (41 °F) the solubility increases to 9.0 mL (50% more than at 25 °C (77 °F)) per liter for freshwater and 7.2 mL (45% more) per liter for sea water.
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 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.
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 spectroscopy of molecular oxygen is associated with the atmospheric processes of aurora and airglow.
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.
There are three allotropes (different forms) of oxygen: (1) atomic oxygen (O), sometimesreferred to as nascent or “newborn” oxygen; (2) diatomic oxygen (O2), or molecular oxygen(gas); and (3) ozone (O3), also a gas.
The atmospheric oxygen that we breathe is a very reactive nonmetal and is colorless, odorless,and tasteless, but it is essential to all living organisms.
Oxygen readily forms compounds withmost other elements.
With six electrons in its outer valence shell, it easily gains two moreelectrons to form a negative (–2) ion; or as covalent, it can share electrons with other elementsto complete its outer shell.
Almost all the oxygen in the atmosphere (21%) is the allotropic form of molecular oxygen(O2).
This essential gas we breathe is the result of photosynthesis, which is how green plants(with chlorophyll) use the energy of the sun to convert carbon dioxide (CO2) and water tostarches and sugars with molecular oxygen as the by-product.
Liquid oxygen has a slightly bluish cast to it.
As it boils, pure oxygen gas is released.
The melting point for oxygen is –218.79°C, its boiling point is –182.95°C, and its density is0.001429 g/cm3.
Uses
In oxyhydrogen or oxyacetylene flame for welding metals and for lighting (calcium light, etc); submarine work by divers, propellant for rockets.
In the production of synthesis gas which can be used in the Fischer-Tropsch process for liquid fuels.
Oxygen has many uses due to its high electronegativity with the ability to oxidize manyother substances.
Only fluorine has higher electronegativity and is thus a stronger oxidizer.
Besides the essential use to support life, oxygen has many other uses.
Oxygen is used in the smelting process to free metals from their ores.
Oxygen is particularly importantin the oxygen-converter process in the production of steel from iron ore.
Oxygen is used in making several important synthetic gases and in the production ofammonia, methyl alcohol, and so on.
Oxygen is the oxidizer for liquid rocket fuels, and as a gas, oxygen is used in a mixture withhelium to support the breathing of astronauts and divers and to aid patients who have difficultybreathing.
Oxygen is use to treat (oxidize) sewage and industrial organic wastes.
Oxygen has many uses because of its ability to accept electrons from other elements to formionic bonds or to share electrons with other elements to form covalent bonds.
The most obvious use of oxygen is to support life.
The process of aerobic respiration uses Oxygen in an oxidation reaction that produces H2O and energy required for metabolic reactions.
Uses of commercially produced O2 gas are mainly industrial and medical. O2 is a highly reactive element that reacts readily with most other chemicals.
The process of combustion uses oxygen as the oxidizing element; for instance, the combustion of hydrocarbons in cars requires O2 in a chemical reaction to produce H2O, CO2, and energy as byproducts.
Similarly, rocket fuel containing H2 gas uses liquid O2 in a combustion reaction to produce water vapor and heat.
Oxygen is used in a wide variety of commercial industries.
Most commercially produced Oxygen is used in the production of steel from iron ore.
Oxygen is used in the production of other metals as well, such as zinc, copper, and lead.
Oxygen is used in the gasification of coal to promote more complete combustion in incinerators; in the production of various chemicals such as ethylene oxide, propylene oxide, and nitric acid; as a fuel component in blow torches for the welding and cutting of metals. i.e., oxyacetylene torch; and in the papermaking industry to bleach pulp.
Oxygen is used in wastewater treatment facilities and to reduce hydrogen sulfide in sewers.
Oxygen has many therapeutic uses in medicine.
Highpurity Oxygen gas is used in life support in surgeries, in intensive care units to assist breathing, for premature babies, and with inhalation therapies with chronic conditions. Hyperbaric oxygen chambers provide oxygen at high partial pressures for numerous conditions, including instances of CO poisoning or in treatment for decompression sickness, or ‘the bends,’ in divers.
Medical
Uptake of Oxygen from the air is the essential purpose of respiration, so oxygen supplementation is used in medicine.
Treatment not only increases oxygen levels in the patient's blood but has the secondary effect of decreasing resistance to blood flow in many types of diseased lungs, easing work load on the heart.
Oxygen therapy is used to treat emphysema, pneumonia, some heart disorders (congestive heart failure), some disorders that cause increased pulmonary artery pressure, and any disease that impairs the body's ability to take up and use gaseous oxygen.
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 Oxygen 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 O2 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 Oxygen 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
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 Oxygen.
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 Oxygen.
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 Oxygen 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 Oxygen with other gases, such as nitrogen or helium, to prevent oxygen toxicity.
People who climb mountains or fly in non-pressurized fixed-wing aircraft sometimes have supplemental Oxygen supplies.
Pressurized commercial airplanes have an emergency supply of Oxygen 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.
A steady stream of oxygen gas is then produced by the exothermic reaction.
Oxygen, as a mild euphoric, has a history of recreational use in oxygen bars and in sports.
Oxygen bars are establishments found in the United States since the late 1990s that offer higher than normal Oxygen exposure for a minimal fee.
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.
Available studies support a performance boost from oxygen enriched mixtures only if it is inhaled during aerobic exercise.
Other recreational uses that do not involve breathing include pyrotechnic applications, such as George Goble's five-second ignition of barbecue grills.
Industrial
Smelting of iron ore into steel consumes 55% of commercially produced oxygen.
In this process, Oxygen is injected through a high-pressure lance into molten iron, which removes sulfur impurities and excess carbon as the respective oxides, SO2 and CO2.
The reactions are exothermic, so the temperature increases to 1,700 °C.
Another 25% of commercially produced oxygen is used by the chemical industry.
Ethylene is reacted with O2 to create ethylene oxide, which, in turn, is converted into ethylene glycol; the primary feeder material used to manufacture a host of products, including antifreeze and polyester polymers (the precursors of many plastics and fabrics).
Most of the remaining 20% of commercially produced oxygen is used in medical applications, metal cutting and welding, as an oxidizer in rocket fuel, and in water treatment.
Oxygen is used in oxyacetylene welding, burning acetylene with Oxygen to produce a very hot flame.
In this process, metal up to 60 cm (24 in) thick is first heated with a small oxy-acetylene flame and then quickly cut by a large stream of O2.
Preparation
Most commercial oxygen at present is obtained from air by cryogenic separation processes.
Although design of oxygen manufacturing plants and process conditions may vary depending on production capacity, purity desired, and cost, basic steps are similar.
Air first is filtered to remove dust particles.
Water and carbon dioxide and most trace impurities are removed by silica gel (or other effective adsorbent) at a temperature slightly above 0°C.
Acetylene and other hydrocarbons also can be removed by such adsorption processes. Alternatively, clean air is compressed and cooled to freeze out water and carbon dioxide, which can be trapped and removed in reversing exchangers.
Compression and cooling of air is a critical step in its liquefaction.
When cooled compressed air is allowed to expand it cools further (Joule-Thomson effect), converting the gaseous air to liquid air at about -196°C.
Liquefied air is subjected to fractional distillation.
More volatile argon and nitrogen distill out on warming, leaving behind oxygen with trace quantities of hydrogen, helium, and other inert gases.
Oxygen may be produced by electrolysis of water.
In such electrolytic procedure, small amounts of H2SO4 or NaOH may be added to water.
Electrolysis methods, however, are not used as much commercially as are air liquefaction processes which cost less.
However, in making hydrogen from water by electrolysis, oxygen is obtained as a by-product.
In the laboratory oxygen may be prepared by several chemical methods that involve thermal decomposition of solid oxides or oxo salts.
The most convenient method of preparing oxygen is to heat potassium chlorate in the presence of manganese dioxide catalyst:
Early preparation of oxygen involved thermal dissociation of metal oxides, notably mercury(II) oxide, which was used independently by both Priestley and Scheele.
Also, oxides of lead, silver, and barium or potassium nitrate and permanganate were used by these and later investigators to prepare oxygen.
Some reactions that yield oxygen by thermal decomposition of metal oxides and metal oxo salts, are highlighted below:
2HgO (s) → 2Hg (l) + O2 (g)
2BaO2 (s) → 2BaO (s) + O2 (g)
2PbO2 (s) → 2PbO (s) + O2 (g)
2KNO3 (s) → 2KNO2 (s) + O2 (g)
2Ag2O (s) → 4Ag (s) + O2 (g)
2KMnO4 (s) → K2MnO4 (s) + MnO2 (s) + O2 (g)
2K2S2O8 (s) → 2K2SO4 (s) + 2SO2 (g) + O2 (g)
Barium peroxide was used in commercial production of oxygen in the past.
Heating barium oxide in air at 500°C forms barium peroxide, which decomposes at 800°C to yield oxygen:
Oxygen can be prepared chemically at ordinary temperatures.
Several reactions in solution are known that may produce small quantities of oxygen at room temperatures.
One such convenient method of producing oxygen is to slowly add water to sodium peroxide.
The reaction is exothermic; therefore, the addition of water must be done cautiously.
2Na2O2 (s) + 2H2O (l) → 4NaOH (aq) + O2 (g)
Oxygen also is liberated when an acidified solution of potassium permangnate, acidified with sulfuric acid, is treated with a solution of hydrogen peroxide:
2MnO4ˉ (aq) + 5H2O2 (aq) + 6H+ (aq) → 2Mn2+ (aq) + 8H2O (l) + 5O2 (g).
