Oxygen (O₂) is the third most abundant element found in the sun, also it plays a part within the carbon-nitrogen cycle, the process once thought to provide the sun and stars their energy. Oxygen under excited conditions accounts for the bright red and yellow-green colors of the Aurora Borealis.

A gaseous element, oxygen forms 21% from the atmosphere by volume and is 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, from the earth's crust. About two thirds of the human body and nine tenths of water is oxygen.

Contents 1 History 2 Properties 2.1 Characteristics of Oxygen 2.2 Physical Forms of Oxygen 2.3 Chemical Properties of Oxygen 2.4 Isotopes 2.4.1 Stable isotopes 2.4.2 Radioisotopes 3 Compounds Of Oxygen 3.1 Oxides and other inorganic compounds 3.2 Organic Compounds And Biomolecules 4 Applications Of Oxygen 4.1 Medical Uses 4.1.1 Life Support 4.2 Industrial Uses 4.3 Scientific Uses

[edit] History

For many centuries, workers occasionally realized air was made up of more than one component. The behaviour of oxygen and nitrogen as components of air resulted in 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 would never know how you can collect it, didn't study its properties, and didn't recognize it as being an elementary substance.

Priestley is usually credited using the discovery of oxygen, although Scheele also discovered it independently in 1774.The word originated from the Greek word oxys and genes meaning sharp or acid and born respectively. Combining these Greek words together forms oxy genes which means acid forming.

Its atomic weight was adopted as a standard of comparison for every from the other elements until 1961 once the International Union of Pure and Applied Chemistry adopted carbon 12 since the new basis.

[edit] Properties

The gas is colorless, odorless, and tasteless. The liquid and solid forms are a pale blue color and are strongly paramagnetic.Oxygen is an active, life-sustaining element of the atmosphere; creating 20.94% by volume or 23% by weight from the air we breathe.

Oxygen, which is very reactive, is really a element of hundreds of thousands of organic compounds and combines with many elements.

[edit] Characteristics of Oxygen

• Oxygen can exist in many physical forms. Probably the most commonly found state of oxygen is the diatomic form and the triplet state. The diatomic form is molecular oxygen and the triplet state of commonly called ozone.

• Ozone (O₃), a highly active compound with a name based on the Greek word for 'I smell', is created by the action of the electrical discharge or ultraviolet light on oxygen.

• O₂ exists in most three forms- solid, liquid and gases.

• The liquid and solid forms are a pale blue color and therefore are strongly paramagnetic.

• Liquid oxygen is potentially hazardous about flames and sparks as it will greatly accelerate combustion.

• Nine isotopes of oxygen are known and natural oxygen is really a mixture of three isotopes.

[edit] Physical Forms of Oxygen

Ozone (O₃), a very active compound, is formed through the action of an electrical discharge or ultraviolet light on oxygen.

Ozone's presence within the atmosphere (amounting to the equivalent of a layer 3 mm thick under ordinary pressures and temperatures) aids in preventing harmful ultraviolet sun's rays from reaching the earth's surface. Pollutants within the atmosphere may have a detrimental effect on this ozone layer. Ozone is toxic and exposure shouldn't exceed 0.2 mg/m# (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 features a cubic crystal structure.

• Oxygen gas is colorless, odorless, and tasteless.

• Boiling Point :- 183oC

• Melting Point :- 218.4oC

[edit] Chemical Properties of Oxygen

• It's the essential aspect in the respiratory processes on most from the living cells as well as in combustion processes.

• Oxygen supports combustion, combines with most elements, and is an element of thousands and thousands of organic compounds.

• Oxygen could be separated from air by fractionated liquefaction and distillation.

• Oxygen is a very reactive oxidizing agent.

• Paramagnetism is strong.

• Liquid oxygen is also slightly paramagnetic.

• Ozone is a highly reactive and powerful reactive agent.

• Every element, except fluorine and also the noble gases, combines spontaneously with oxygen at galactic standard temperature and pressure.

• The singlet form of oxygen reacts swiftly with almost all compounds.

• Molecular oxygen is a stable diradical.

• It's two electrons in an unpaired triplet state. Oxygen is the only naturally occurring chemical with this property.

Electronegativity	: 3.4
Heat of vaporization	: 3.4099 kj/mol
Heat of fusion	: 0.22259 kj/mol
First Ionization Energy	: 1314 kJ/mol

[edit] Isotopes

Oxygen has nine isotopes. Natural oxygen is really a mixture of three isotopes.

[edit] Stable isotopes

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

The relative and absolute abundance of ¹⁶O is high because it is a principal product of stellar evolution and because it is a primary isotope, meaning it can be made by stars that were initially made exclusively of hydrogen.Most ¹⁶O is synthesized at the end of the helium fusion process in stars; the triple-alpha reaction creates ¹²C, which captures an additional ⁴He to make ¹⁶O. The neon burning process creates additional ¹⁶O.

Both ¹⁷O and ¹⁸O are secondary isotopes, meaning that their nucleosynthesis requires 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. Most ¹⁸O is produced when ¹⁴N (made abundant from CNO burning) captures a ⁴He nucleus, making ¹⁸O common in the helium-rich zones of stars.Approximately a billion degrees Celsius is required for two oxygen nuclei to undergo nuclear fusion to form the heavier nucleus of sulfur.

[edit] Radioisotopes

Fourteen radioisotopes have been characterized, with the most stable being ¹⁵O with a half-life of 122.24 s and ¹⁴O with a half-life of 70.606 s. 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 (ms). For example, ²⁴O has a half life of 61 ms. The most common decay mode for isotopes lighter than the stable isotopes is β+ decay (to nitrogen) and the most common mode after is β- decay (to fluorine).

An atomic mass 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 mass scales between the two disciplines.

The isotopic composition of oxygen atoms in the Earth's atmosphere is 99.759% ¹⁶O, 0.037% ¹⁷O and 0.204% ¹⁸O. Because water molecules containing the lighter isotope are slightly more likely to evaporate and fall as precipitation, fresh water and polar ice on earth contains slightly less (0.1981%) of the heavy isotope ¹⁸O than air (0.204%) or seawater containing (0.1995%). This disparity allows analysis of temperature patterns via historic ice cores.

[edit] Compounds Of Oxygen

The oxidation state of oxygen is -2 in almost all known compounds of oxygen. The oxidation state -1 is found in several compounds for example peroxides. Compounds containing oxygen in other oxidation states are extremely uncommon: -1/2 (superoxides), -1/3 (ozonides), 0 (elemental, hypofluorous acid), +1/2 (dioxygenyl), +1 (dioxygen difluoride), and +2 (oxygen difluoride).

[edit] Oxides and other inorganic compounds

Water (H₂O) is the oxide of hydrogen and also the most familiar oxygen compound. Hydrogen atoms are covalently bonded to oxygen in a water molecule but additionally come with an additional attraction (about 23.3 kJ·mol-1 per hydrogen atom) for an adjacent oxygen atom in a separate molecule. These hydrogen bonds between water molecules hold them approximately 15% closer than what will be expected in a simple liquid with just van der Waals forces.

Because of its electronegativity, oxygen forms chemical bonds with almost every other elements at elevated temperatures to give corresponding oxides. However, some elements readily form oxides at standard conditions for temperature and pressure; the rusting of iron is definitely an example. The top of metals like aluminium and titanium are oxidized within the presence of air and become coated having a thin film of oxide that passivates the metal and slows further corrosion. A few of the transition metal oxides are simply anyway as non-stoichiometric compounds, with a slightly less metal than the chemical formula would show. For example, the natural occurring FeO (wüstite) is really written as Fe1 -_xO, where x is usually around 0.05.

Oxygen as a compound is present in the atmosphere in trace quantities as carbon dioxide (CO₂). The earth's crustal rock is composed in large part of oxides of silicon (silica SiO₂, present in granite and sand), aluminium (aluminium oxide Al₂O₃, in bauxite and corundum), iron (iron(III) oxide Fe₂O₃, in hematite and rust) along with other metals.

All of those other Earth's crust can also be made from oxygen compounds, particularly calcium carbonate (in limestone) and silicates (in feldspars). Water-soluble silicates as Na₄SiO₄, Na₂SiO₃, and Na₂Si₂O₅ are used as detergents and adhesives.

Oxygen also acts as a ligand for transition metals, forming metal-O₂ bonds using the iridium atom in Vaska's complex, with the platinum in PtF₆, and using the iron center from the heme group of hemoglobin.

[edit] Organic Compounds And Biomolecules

One of the most important classes of organic compounds which contain oxygen are (where "R" is an organic group): alcohols (R-OH); ethers (R-O-R); ketones (R-CO-R); aldehydes (R-CO-H); carboxylic acids (R-COOH); esters (R-COO-R); acid anhydrides (R-CO-O-CO-R); and amides (R-C(O)-NR₂). There are many important organic solvents which contain oxygen, including: acetone, methanol, ethanol, isopropanol, furan, THF, diethyl ether, dioxane, ethyl acetate, DMF, DMSO, acetic acid, and formic acid. Acetone ((CH₃)2CO) and phenol (C₆H₅OH) are used as feeder materials in the synthesis of numerous different substances. Other important organic compounds which contain oxygen are: glycerol, formaldehyde, glutaraldehyde, citric acid, acetic anhydride, and acetamide. Epoxides are ethers where the oxygen atom is a component of the ring of three atoms.

Oxygen reacts spontaneously with many organic compounds at or below room temperature inside a process called autoxidation. Most from the organic compounds which contain oxygen aren't produced by direct action of O₂. Organic compounds essential in industry and commerce which are made by direct oxidation of the precursor include ethylene oxide and peracetic acid.

The element can be found in almost all biomolecules that are vital that you (or generated by) life. Just a few common complex biomolecules, for example squalene and the carotenes, contain no oxygen. Of the organic compounds with biological relevance, carbohydrates retain the largest proportion by mass of oxygen. All fats, fatty acids, amino acids, and proteins contain oxygen (due to the presence of carbonyl groups in these acids and their ester residues). Oxygen also happens in phosphate (PO³_-4) groups within the biologically important energy-carrying molecules ATP and ADP, in the backbone and also the purines (except adenine) and pyrimidines of RNA and DNA, and in bones as calcium phosphate and hydroxylapatite.

[edit] Applications Of Oxygen

[edit] Medical Uses

Uptake of O₂ from the air may be the essential reason for respiration, so oxygen supplementation can be used in medicine. Treatment not just increases oxygen levels within the patient's blood, but has got the secondary effect of decreasing potential to deal with blood circulation in several kinds of diseased lungs, easing work on the heart. Oxygen therapy is accustomed 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 employ gaseous oxygen.

Remedies 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 improve the partial pressure of O₂ around the patient and, if needed, the medical staff. Carbon monoxide poisoning, gas gangrene, and decompression sickness (the 'bends') are occasionally treated using these devices. Increased O₂ concentration in the lungs helps you to displace carbon monoxide in the heme group of hemoglobin. Oxygen gas is poisonous towards the anaerobic bacteria that cause gas gangrene, so increasing its partial pressure helps kill them. Decompression sickness occurs in divers who decompress prematurely after a dive, leading to bubbles of inert gas, mostly nitrogen and helium, forming within their blood. Increasing the pressure of O₂ as quickly as possible is part of the treatment.

Oxygen can also be used medically for patients who require mechanical ventilation, often at concentrations above the 21% present in ambient air.

[edit] Life Support

A notable application of O₂ as a low-pressure breathing gas is within modern space suits, which surround their occupant's body with pressurized air. These devices use nearly pure oxygen at about 1 / 3 normal pressure, producing a normal blood partial pressure of O₂. This trade-off better oxygen concentration for lower pressure is needed to maintain flexible spacesuits.

Scuba divers and submariners also depend on artificially delivered O₂, but many often use normal pressure, and/or mixtures of oxygen and air. Pure or nearly pure O₂ use within diving at higher-than-sea-level pressures is generally restricted to rebreather, decompression, or emergency treatment use at relatively shallow depths (~6 meters depth, or less). Deeper diving requires significant dilution of O₂ along with other gases, for example nitrogen or helium, to assist prevent oxygen toxicity.

People who climb mountains or fly in non-pressurized fixed-wing aircraft sometimes have supplemental O₂ supplies. Passengers traveling in (pressurized) commercial airplanes come with an emergency supply of O₂ automatically supplied to them in the event of cabin depressurization. Sudden cabin pressure loss activates chemical oxygen generators above each seat, causing oxygen masks to drop. Pulling about the masks "to start the flow of oxygen" as cabin safety instructions dictate, forces iron filings to the sodium chlorate within the canister. A steady flow of oxygen gas is then made by the exothermic reaction.

Oxygen, as a supposed mild euphoric, has a good reputation for recreational use within oxygen bars as well as in sports. Oxygen bars are establishments, present in Japan, California, and Las Vegas, Nevada because the late 1990s that offer greater than normal O₂ exposure for a small fee. Professional athletes, especially in American football, also sometimes go off field between plays to put on oxygen masks in order to get a "boost" in performance. The pharmacological effect is doubtful; a 'placebo effect' is really a more likely explanation. Available studies support a performance boost from enriched O₂ mixtures only when they are breathed during actual aerobic exercise.

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

[edit] Industrial Uses

Smelting of iron ore into steel consumes 55% of commercially produced oxygen. On this process, O₂ is injected through a high-pressure lance into molten iron, which removes sulfur impurities and excess carbon since the respective oxides, SO₂ and CO₂. The reactions are exothermic, so the temperature increases to at least one,700 °C.

Another 25% of commercially produced oxygen is used through the chemical industry. Ethylene is reacted with O₂ to produce ethylene oxide, which, consequently, is converted into ethylene glycol; the main feeder material used to manufacture a host of products, including antifreeze and polyester polymers (the precursors of numerous plastics and fabrics).

The majority of the remaining 20% of commercially produced oxygen is used in medical applications, metal cutting and welding, being an oxidizer in rocket fuel, and in water treatment. Oxygen can be used in oxyacetylene welding burning acetylene with O₂ to produce a very hot flame. On this process, metal as much as 60 cm thick is first heated having a small oxy-acetylene flame after which quickly cut by a large stream of O₂. Larger rockets use liquid oxygen as their oxidizer, which is mixed and ignited using the fuel for propulsion.

[edit] Scientific Uses

Paleoclimatologists measure the ratio of oxygen-18 and oxygen-16 in the shells and skeletons of marine organisms to find out what are the climate was like millions of years ago. 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. During periods of lower global temperatures, snow and rain from that evaporated water is commonly higher in oxygen-16, and the seawater left out is commonly higher in oxygen-18. Marine organisms then incorporate more oxygen-18 into their skeletons and shells than they'd in a warmer climate. Paleoclimatologists also directly measure this ratio in the water molecules of ice core samples which are as much as several thousands and thousands of years of age.

Planetary geologists have measured different abundances of oxygen isotopes in samples in the Earth, the Moon, Mars, and meteorites, but were long not able to obtain reference values for that isotope ratios in the Sun, thought to be just like the ones from the primordial solar nebula. However, analysis of the silicon wafer subjected to the solar wind in space and returned through the crashed Genesis spacecraft indicates that the Sun has a higher proportion of oxygen-16 than does the Earth. The measurement signifies that an unknown process depleted oxygen-16 from the Sun's disk of protoplanetary material before the coalescence of dust grains that formed the Earth.

Oxygen presents two spectrophotometric absorption bands peaking in the wavelengths 687 and 760 nm. Some remote sensing scientists have proposed while using measurement of the radiance coming from vegetation canopies in those bands to characterize plant health status from the satellite platform. This approach exploits the truth that in those bands it is possible to discriminate the vegetation's reflectance from the fluorescence, that is much weaker. The measurement is technically difficult owing to the low signal-to-noise ratio and also the physical structure of vegetation; but it continues to be proposed just as one approach to monitoring the carbon cycle from satellites on a global scale. source 1