8
  O  
15.999400
Oxygen

Name: Oxygen
Symbol: O
Atomic Number: 8
Atomic Weight: 15.999400
Family: Non Metals
CAS RN: 7782-44-7
Description: A colorless gas, liquid is pale blue.
State (25C): Gas
Oxidation states: -2
, +4, +6
Molar Volume: 14 cm3/mole
Valence Electrons: 2p4

Boiling Point:  90.33K, -182.82C, -297.08F
Melting Point:
50.5K, -222.65C, -368.77F
Electrons Energy Level: 2, 6
Isotopes: 10 + 3 Stable
Heat of Vaporization: 3.4099 kJ/mol
Heat of Fusion: 0.22259 kJ/mol
Density: 1.429 g/L @ 273K & 1atm
Specific Heat: 0.92 J/gK
Atomic Radius: 0.65
Ionic Radius: 1.4
Electronegativity: 3.44 (Pauling); 3.5 (Allrod Rochow)
Named from the Greek oxys + genes, "acid-former" oxygen was first described by Michal Sedziwoj, a Polish alchemist and philosopher in the late 16th century. Sedziwj thought of the gas given off by warm nitre (saltpeter) as "the elixir of life".

Oxygen was more quantitatively discovered by the Swedish pharmacist Carl Wilhelm Scheele some time before 1773, but the discovery was not published until after the independent discovery by Joseph Priestley on August 1, 1774, who called the gas dephlogisticated air.  Priestley published discoveries in 1775 and Scheele in 1777; consequently Priestley is usually given the credit. Both Scheele and Priestley produced oxygen by heating mercuric oxide (HgO).

Scheele called the gas 'fire air' because it was the only known supporter of combustion. It was later called 'vital air' because it was and is vital for the existence of animal life.

The gas was named by Antoine Laurent Lavoisier, after Priestley's publication in 1775, from Greek roots meaning "acid-former".  As noted, the name reflects the then-common incorrect belief that all acids contain oxygen. 

1s2 2s2p4

8
O
15.99
16
S
32.06
34
Se
78.96
52
Te
127.6
84
Po
210.0
116
Uuh
289.0

Characteristics

Oxygen is the third most abundant element in the universe and the most common element in the earth's crust (two thirds of the mass of the human body and nine tenths of the mass of water).and makes up about 20% of the air we breathe.  Large amounts of oxygen can be extracted from liquefied air through a process known as fractional distillation. Oxygen can also be produced through the electrolysis of water or by heating potassium chlorate (KClO3).  Historically the discovery of oxygen as an element essential for combustion stands at the heart of the phlogiston controversy.

Oxygen is a gas at room temperature and is colorless, odorless and tasteless. Liquid oxygen has a slight blue color.  Virtually every element known forms a compound with oxygen except for some of the noble gases.  Two allotropes of oxygen are known.  The ordinary oxygen in the air (O2) and the ordinary ozone in air and above it (O3).  Ozone occurs naturally in the upper atmosphere where it acts as a shield protecting the earth's biosphere from ultraviolet radiation.   Closer to the surface ozone can be a nuisance as it is very reactive.  But in a controlled environment it is also handy: it makes a good bleach and disinfectant.

1s2
2s2 2p4

Oxygen is one of the two major components of air.  It is produced by plants during photosynthesis, and is necessary for aerobic respiration in animals.  The word oxygen derives from two roots in Greek, (oxys) (acid, sharp) and (-genes) (born of).  In the early 18th century, Antoine Lavoisier coined the name oxygen from the Greek roots mentioned above because he erroneously thought that it was a constituent of all acids.  (The definition of acid has since been revised).  On Earth it is usually bonded to other elements covalently or ionically.

Unbound oxygen (also called molecular oxygen, or dioxygen, O2, a diatomic molecule) first appeared in significant quantities in Earth's atmosphere during the Paleoproterozoic Era (between 2.5 billion years ago and 1.6 billion years ago) as a product of the metabolic action of early anaerobes ( archaea and bacteria).  The presence of large amounts of free oxygen in the atmosphere may have driven most of the organisms then living to extinction.  The atmospheric abundance of free oxygen in later geological epochs and its gradual increase up to the present has been largely due to synthesis by photosynthetic organisms; organisms; about three quarters of the free element being produced by algae and green microorganisms in the oceans, and one quarter from terrestrial plants.

At standard temperature and pressure, oxygen exists as a diatomic molecule with the formula O2, in which the two oxygen atoms are bonded to each other with the electron configuration of triplet oxygen.  This bond has a bond order of two, and is thus often very grossly simplified in description as a double bond.  Triplet oxygen is the ground state of the oxygen molecule. The electron configuration of the molecule has two unpaired electrons occupying two degenerate molecular orbitals.  These orbitals are classified as antibonding, so the diatomic oxygen bond is weaker than the diatomic nitrogen bond, where all bonding molecular orbitals are filled.  Though unpaired electrons are commonly associated with high reactivity in chemical compounds, triplet oxygen is relatively (and fortunately) unreactive by comparison with most radicals.

Singlet oxygen, a name given to several higher energy species of molecular oxygen in which all the electron spins are paired, is much more reactive towards common organic molecules.  In nature, singlet oxygen is commonly formed from water during photosynthesis, using the energy of sunlight.  It is also produced by the immune system as a source of active oxygen.  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.

Liquid O2 and solid O2 are clear substances with a light sky-blue color.  In normal triplet form they are paramagnetic due to the spin magnetic moments of the unpaired electrons in the molecule, and the negative exchange energy between neighboring O2 molecules.  Liquid oxygen is attracted to a magnet to a sufficient extent that a bridge of liquid oxygen may be supported against its own weight between the poles of a powerful magnet, in laboratory demonstrations.  Liquid O2 is usually obtained by the fractional distillation of liquid air.

Oxygen is slightly soluble in water, but naturally occurring dissolved amounts are enough to support animal life.  O2 has a bond length of 121 pm and a bond energy of 498 kJ/mol.

Allotropes

100px-Dioxygen-montage1.jpg (5631 bytes)

Dioxygen, O2, is a gas at standard conditions, consisting of 2-atom molecules.  Elemental oxygen is most commonly encountered in this form, as 21% of Earth's atmosphere.  Note that the double bond depicted here is an oversimplification.

100px-Ozone-montage2.jpg (5930 bytes)

Ozone, O3, is a gas at standard conditions, consisting of 3-atom molecules.

This oxygen allotrope is rare on Earth and is found mostly in the stratosphere.   The common allotrope of elemental oxygen on Earth, O2, is known as dioxygen.

Ozone, the less common triatomic allotrope of oxygen, is a poisonous gas with a sharp odor.  It is thermodynamically unstable toward the more common dioxygen form.   It is formed continuously in the upper atmosphere of the Earth by short-wave UV radiation, and also functions as a shield against UV radiation reaching the ground.   Ozone has recently been found to be produced by the immune system as an antimicrobial (see below).  Liquid and solid O3 (ozone) have a deeper blue color than ordinary oxygen, and they are unstable and explosive.

A recently discovered allotrope of oxygen, tetraoxygen (O4), is a deep red solid that is created by pressurizing O2 to the order of 20 GPa.  Its properties are being studied for use in rocket fuels  and similar applications, as it is a much more powerful oxidizer than either O2 or O3.

Occurrence

Oxygen is the third most abundant chemical element in the universe by mass, after hydrogen and helium.  Some of this Oxygen was produced during stellar nucleosynthesis as a step in the CNO-II branch of the CNO cycle.  However Oxygen is primarily produced in massive stars. In stars with at least four times the Sun's mass,  16O nuclei are produced during the Carbon burning process.   16O can also be produced in stars with at least 8 times the Sun's mass as a result of photodisintegration during the Neon burning process.

Oxygen is the most common component of the Earth's crust (49% by mass), the second most common component of the Earth as a whole (28% by mass), the most common component of the world's oceans (86% by mass), and the second most common component of the Earth's atmosphere (20.947% by volume), second to nitrogen.

Elemental oxygen occurs not only in the atmosphere, but also as solution in the world's water bodies. At 25 C under 1 atm of air, a liter of water will dissolve about 6.04 cc (8.63 mg, 0.270 mmol) of oxygen, whereas sea water will dissolve about 4.9 cc (7.0 mg, 0.22 mmol). At 0 C the solubility's increase to 10.29 cc (14.7 mg, 0.460 mmol) for water and 8.0 cc (11.4 mg, 0.36 mmol) for sea water. This difference has important implications for ocean life, as polar oceans support a much higher density of life due to their oxygen content.

Applications

Uptake of oxygen from the air is the essential purpose of respiration, so oxygen supplementation has found use in medicine (as oxygen therapy).  People who climb mountains or fly in non-pressurized airplanes sometimes have supplemental oxygen supplies; the reason is that increasing the proportion of oxygen in the breathing gas at low pressure acts to augment the inspired oxygen partial pressure nearer to that found at sea-level.  A notable application of oxygen as a very low-pressure breathing gas, is in modern spacesuits, where use of nearly pure oxygen at a total ambient pressure of about one third normal, results in normal blood partial pressures of oxygen.  This trade-off of breathing gas content and needed pressure is important for space applications, because the issue of flexible spacesuits working at Earth sea-level pressures remains a technological challenge of aerospace technology.

Oxygen is used in welding (such as the oxyacetylene torch), and in the industrial production of steel and methanol.  Also, liquid oxygen finds use as a classic oxidizer in rocket propulsion.

Oxygen presents two spectrophotometric absorption bands peaking at the wavelengths 687 and 760 nanometers. Some scientists have proposed to use the measurement of the radiance coming from vegetation canopies in those oxygen bands to characterize plant health status from a satellite platform.  This is because in those bands, it is possible to discriminate the vegetation's reflectance from the vegetation's fluorescence, which is much weaker. The measurement presents several technical difficulties due to the low signal to noise ratio and due to the vegetation's architecture, but it has been proposed as a possibility to monitor the carbon cycle from satellites on a global scale.

Oxygen, as a supposed mild euphoric, has a history of recreational use.  However, the reality of a pharmacological effect is doubtful being a metabolic boost the most plausible explanation. Controlled tests of high oxygen mixtures in diving and other activities, even at higher than normal pressures, demonstrated no particular effects on humans other than promotion of an increased tolerance to aerobic exercise.

In the 19th century, oxygen was often mixed with nitrous oxide to temper its analgesic effect.  A stable 50% gaseous mixture (Entonox) is commonly used in medicine today as an analgesic. However, the common basic anaesthetic mixture is 30% oxygen with 70% nitrous oxide; the pain-suppressing effects, obviously, are due to the nitrous oxide and not to oxygen.

Oxygen is a highly reactive element and is capable of combining with most other elements. It is required by most living organisms and for most forms of combustion. Impurities in molten pig iron are burned away with streams of high pressure oxygen to produce steel. Oxygen can also be combined with acetylene (C2H2) to produce an extremely hot flame used for welding. Liquid oxygen, when combined with liquid hydrogen, makes an excellent rocket fuel. Ozone (O3) forms a thin, protective layer around the earth that shields the surface from the sun's ultraviolet radiation. Oxygen is also a component of hundreds of thousands of organic compounds.

Compounds

The most familiar of oxygen compounds is water.

Due to its electronegativity, oxygen forms chemical bonds with almost all other elements hence the original definition of oxidation.  The only elements known to escape the possibility of oxidation are a few of the noble gases, and fluorine.   Other than water (H2O), well known examples include compounds of carbon and oxygen, such as carbon dioxide (CO2), alcohols  (R-OH), carbonyls, (R-CO-H or R-CO-R)), and carboxylic acids (R-COOH).  Oxygenated radicals such as chlorates ClO3-), perchlorates (ClO4-), chromates (CrO42-), dichromates (Cr2O72-), permanganates (MnO4-), and nitrates (NO3-) are strong oxidizing agents in and of themselves.  Many metals such as iron bond with oxygen atoms, iron (III) oxide (Fe2O3).  Ozone (O3) is formed by electrostatic discharge in the presence of molecular oxygen.  A double oxygen molecule (O2)2 is known and is found as a minor component of liquid oxygen.  Epoxides are ethers in which the oxygen atom is part of a ring of three atoms.

One unexpected oxygen compound is dioxygen hexafluoroplatinate O2+PtF6-.   It was discovered when Neil Bartlett was studying the properties of PtF6.   He noticed a change in color when this compound was exposed to atmospheric air. Bartlett reasoned that xenon should be oxidized by PtF6.  This led him to the discovery of xenon hexafluoroplatinate Xe+PtF6-.

Isotopes

Oxygen has ten known isotopes with atomic masses ranging from 12.03 u to 28.06 u.  Three are stable, 16O, 17O, and 18O, of which 16O is the most abundant (over 99.7%). The radioisotopes all have half-lives of less than three minutes.

An atomic weight of 16 was assigned to oxygen prior to the definition of the unified atomic mass unit based upon 12C.  Since physicists referred to 16O only, while chemists meant the naturally abundant mixture of isotopes, this led to slightly different atomic weight scales.

atom.gif (700 bytes)

Isotope Atomic Mass Half-Life
O12 12.0344 0.4 MeV
O13 13.0248 8.58 ms
O14 14.0086 70.606 seconds
O15 15.0031 122.24 seconds
O16 15.9949 Stable
O17 16.9991 Stable
O18 17.9992 Stable
O19 19.0036 26.91 seconds
O20 20.0041 13.51 seconds
O21 21.0087 3.42 seconds
O22 22.01 2.25 seconds
O23 23.016 82 ms
O24 24.02 61 ms
O25 25.029  
O26 26.038  

Precautions

40px-Skull_and_crossbones.svg.jpg (1420 bytes) Oxygen can be toxic at elevated partial pressures.  Since oxygen partial pressure is the fraction of oxygen times the total pressure, elevated partial pressures can occur either from high oxygen fraction in breathing gas, or from high breathing gas pressure, or a combination of both.

Oxygen toxicity usually begins to occur at partial pressures more than 0.5 atmospheres, or 2.5 times the normal sea-level oxygen partial pressure of about 0.2 atmospheres or bars.  This means that at sea-level pressures, mixtures containing less than 50% oxygen are essentially non-toxic.  However in medical applications (such as in ventilation gas mixtures in hospital applications) mixtures containing more than 50% oxygen can be expected to show lung toxicity, causing slow damage to the lungs over periods of days, with the rate of damage rising rapidly from mixtures between 50% and 100% oxygen.  On the other hand, breathing 100% oxygen in space applications (such as in some modern spacesuits, or in early spacecraft such as the Apollo spacecraft), causes no damage due to the low total pressures (30% to 33% sea-level) used.  In the case of spacesuits, oxygen partial pressure in the breathing gas is typically about 0.30 bar (1.4 times normal), and oxygen partial pressure in the astronaut's blood (due to downward adjustments due to water vapor and CO2 in the alveoli) is close to sea-level normal of 0.14 bar.

In deep scuba diving and surface supplied diving and when using equipment which can provide high partial pressures of oxygen, such as rebreathers, oxygen toxicity to the lungs can occur, just as in medical applications. Due to the higher total pressures in these applications, the fraction of oxygen which produces lung damage may be considerably less than 50%. More importantly, under pressures higher than normal sea-level, a far more serious form of oxygen toxicity in the central nervous system may lead to generalized seizures. This form of oxygen toxicity usually occurs after several hours exposure to oxygen partial pressures over about 1.4 atmospheres (bars) (i.e. 7 times normal), with the time decreasing for higher pressures above this, and with great variation from person to person. At over three bars of oxygen partial pressure (15 times normal), seizures typically occur within minutes.

Toxicity of Other Oxygen Forms

Certain derivatives of oxygen, such as ozone (O3), singlet oxygen, hydrogen peroxide, hydroxyl radicals and superoxide, are also highly toxic.  Cells have developed various mechanisms to protect against all of these toxic compounds. For instance, the naturally-occurring glutathione can act as an antioxidant, as can bilirubin which is normally a breakdown product of hemoglobin.  To protect against the destructive nature of peroxides, nearly every organism on earth has developed some form of the enzyme catalase, which very quickly disproportionates hydrogen peroxide into water and dioxygen. Another nearly universally present enzyme in living organisms (except for a few species of bacteria which use Mn2+ ions directly for the job) is superoxide dismutase.  This family of enzymes disproportionates superoxide to oxygen and peroxide, which is then in turn dealt with, by catalase.

Immune systems of higher organisms have long made use of reactive forms of oxygen which they produce. Not only do antibodies catalyze production of peroxide from oxygen, it is now known that immune cells produce peroxide, superoxide, and singlet oxygen in the course of an immune response. Recently, singlet oxygen has been found to be a source of biologically-produced ozone: this reaction proceeds through an unusual compound dihydrogen trioxide, also known as trioxadane, (HOOOH) which is an antibody-catalyzed product of singlet oxygen and water. This compound in turn disproportionates to ozone and peroxide, providing two powerful antibacterials. The body's range of defense against all of these active oxidizing agents is hardly surprising, then, given their "deliberate" employment as antimicrobial agents in the immune response.

Oxygen derivatives are prone to form free radicals, especially in metabolic processes. Because they can cause severe damage to cells and their DNA before they are dealt with, they form part of many theories of carcinogenesis and aging.

Combustion Hazard

80px-Flammable.jpg (2186 bytes) Highly concentrated sources of oxygen promote rapid combustion and therefore are fire and explosion hazards in the presence of fuels.  Oxygen itself is not the fuel, but as a reactant, concentrated oxygen may allow combustion to proceed dangerously rapidly.  

The fire that killed the Apollo 1 crew on a test launch pad spread so rapidly because the capsule was pressurized with pure oxygen as would be usual in an actual flight, but to maintain positive pressure in the capsule, this was at slightly more than atmospheric pressure instead of the normal pressure that would be used in flight.

Hazards also apply to compounds of oxygen with a high oxidative potential, such as high concentration peroxides, chlorates, perchlorates, and dichromates; they also can often cause chemical burns.

atom.gif (700 bytes)

Oxygen Data

Atomic Radius (): 0.65
Atomic Volume cm3/mol : 14cm3/mol
Covalent Radius: 0.73
Crystal Structure: Cubic
Ionic Radius: 1.4

Chemical Properties

Electrochemical Equivalents: 0.29847g/amp-hr
Electron Work Function: unknown
Electronegativity: 3.44 (Pauling); 3.5 (Allrod Rochow)
Heat of Fusion: 0.22259 kJ/mol
Incompatibilities: oxidizable materials
First Ionization Potential: 13.618
Second Ionization Potential: 35.117
Third Ionization Potential: 54.934
Valence Electron Potential: -20.6
Ionization Energy (eV): 13.618 eV

Physical Properties

Atomic Mass Average: 15.9994
Boiling Point: 90.33K, -182.82C, -297.08F
Melting Point: 50.5K, -222.65C, -368.77F
Heat of Vaporization: 3.4099 kJ/mol
Coefficient of Lineal Thermal Expansion/K-1: N/A
Electrical Conductivity: unknown
Thermal Conductivity: 0.0002674 W/cmK
Density: 1.429 g/L @ 273K & 1atm
Enthalpy of Atomization: 249.4 kJ/mole @ 25C
Enthalpy of Fusion: 0.22 kJ/mole
Enthalpy of Vaporization: 3.41 kJ/mole
Flammability Class: Non-flammable gas (Oxidizer)
Molar Volume: 14 cm3/mole
Optical Refractive Index: 1.000271 (gas) 1.221 (liquid)
Relative Gas Density (Air=1): unknown
Specific Heat: 0.92 J/gK
Vapor Pressure: unknown
Estimated Crustal Abundance: 4.61105 milligrams per kilogram
Estimated Oceanic Abundance: 8.57105 milligrams per liter


(Gr. oxys, sharp, acid, and genes, forming; acid former) For many centuries, workers occasionally realized air was composed of more than one component.  The behavior of oxygen and nitrogen as components of air led to the advancement of the phlogiston theory of combustion, which captured the minds of chemists for a century.  Oxygen was prepared by several workers, including Bayen and Borch, but they did not know how to collect it, did not study its properties, and did not recognize it as an elementary substance.  Priestley is generally credited with its discovery, although Scheele also discovered it independently.  Oxygen is the third most abundant element found in the sun, and it plays a part in the carbon-nitrogen cycle, once process thought to give the sun and stars their energy.  Oxygen under excited conditions is responsible for the bright red and yellow-green colors of the Aurora.  Oxygen, as a gaseous element, forms 21% of the atmosphere by volume from which it can be obtained by liquefaction and fractional distillation.  The atmosphere of Mars contains about 0.15% oxygen.   The element and its compounds make up 49.2%, by weight, of the earth's crust.   About two thirds of the human body and nine tenths of water is oxygen.  In the laboratory it can be prepared by the electrolysis of water or by heating potassium chlorate with manganese dioxide as a catalyst.  The gas is colorless, odorless, and tasteless.  The liquid and solid forms are a pale blue color and are strongly paramagnetic.  Ozone (O3), a highly active compound, is formed by the action of an electrical discharge or ultraviolet light on oxygen.  Ozone's presence in the atmosphere (amounting to the equivalent of a layer 3 mm thick at ordinary pressures and temperatures) is of vital importance in preventing harmful ultraviolet rays of the sun from reaching the earth's surface.  There has been recent concern that pollutants in the atmosphere may have a detrimental effect on this ozone layer.  Ozone is toxic and exposure should not exceed 0.2 mg/m3 (8-hour time-weighted average - 40-hour work week). Undiluted ozone has a bluish color.  Liquid ozone is bluish black and solid ozone is violet-black.  Oxygen is very reactive and capable of combining with most elements. It is a component of hundreds of thousands of organic compounds.  It is essential for respiration of all plants and animals and for practically all combustion.   In hospitals it is frequently used to aid respiration of patients.  Its atomic weight was used as a standard of comparison for each of the other elements until 1961 when the International Union of Pure and Applied Chemistry adopted carbon 12 as the new basis.   Oxygen has ten isotopes.  Natural oxygen is a mixture of three isotopes.   Oxygen 18 occurs naturally, is stable, and is available commercially.  Water (H2O with 15% 18O) is also available.  Commercial oxygen consumption in the U.S. is estimated to be 20 million short tons per year and the demand is expected to increase substantially in the next few years.  Oxygen enrichment of steel blast furnaces accounts for the greatest use of the gas.  Large quantities are also used in making synthesis gas for ammonia and methanol, ethylene oxide, and for oxy-acetylene welding.  Air separation plants produce about 99% of the gas, electrolysis plants about 1%.  The gas costs 5 cents / ft3 in small quantities, and about $15/ton in large quantities.

Source: CRC Handbook of Chemistry and Physics, 1913-1995. David R. Lide, Editor in Chief. Author: C.R. Hammond