Oxygen Is Odorless And Colorless

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oxygen (O), nonmetallic chemical element of Group 16 (VIa, or the oxygen grouping) of the periodic table. Oxygen is a colourless, odourless, tasteless gas essential to living organisms, existence taken up by animals, which catechumen it to carbon dioxide; plants, in plow, employ carbon dioxide as a source of carbon and render the oxygen to the temper. Oxygen forms compounds by reaction with practically any other element, every bit well as by reactions that displace elements from their combinations with each other; in many cases, these processes are accompanied by the evolution of rut and light and in such cases are called combustions. Its most important chemical compound is water.

Element Properties
atomic number	8
atomic weight	15.9994
melting point	−218.iv °C (−361.1 °F)
boiling point	−183.0 °C (−297.four °F)
density (1 atm, 0 °C)	1.429 m/litre
oxidation states	−1, −2, +2 (in compounds with fluorine)
electron config.	onesouth ²twodue south ^two2p ⁴

History

Oxygen was discovered about 1772 by a Swedish pharmacist, Carl Wilhelm Scheele, who obtained it by heating potassium nitrate, mercuric oxide, and many other substances. An English pharmacist, Joseph Priestley, independently discovered oxygen in 1774 by the thermal decomposition of mercuric oxide and published his findings the same yr, 3 years before Scheele published. In 1775–eighty, French pharmacist Antoine-Laurent Lavoisier, with remarkable insight, interpreted the role of oxygen in respiration too as combustion, discarding the phlogiston theory, which had been accepted upwardly to that time; he noted its trend to form acids by combining with many different substances and accordingly named the element oxygen (oxygène) from the Greek words for "acrid former."

Illustration of molecules. (molecular, chemistry, science)

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Occurrence and properties

At 46 per centum of the mass, oxygen is the most plentiful element in World's crust. The proportion of oxygen past volume in the atmosphere is 21 per centum and past weight in seawater is 89 percent. In rocks, it is combined with metals and nonmetals in the form of oxides that are acidic (such equally those of sulfur, carbon, aluminum, and phosphorus) or basic (such as those of calcium, magnesium, and iron) and as saltlike compounds that may be regarded equally formed from the acidic and basic oxides, as sulfates, carbonates, silicates, aluminates, and phosphates. Plentiful as they are, these solid compounds are not useful as sources of oxygen, because separation of the element from its tight combinations with the metallic atoms is too expensive.

Beneath −183 °C (−297 °F), oxygen is a stake blue liquid; information technology becomes solid at nigh −218 °C (−361 °F). Pure oxygen is 1.ane times heavier than air.

During respiration, animals and some bacteria take oxygen from the atmosphere and return to it carbon dioxide, whereas past photosynthesis, green plants assimilate carbon dioxide in the presence of sunlight and evolve gratuitous oxygen. Almost all the free oxygen in the temper is due to photosynthesis. About 3 parts of oxygen past book deliquesce in 100 parts of fresh water at 20 °C (68 °F), slightly less in seawater. Dissolved oxygen is essential for the respiration of fish and other marine life.

Natural oxygen is a mixture of three stable isotopes: oxygen-sixteen (99.759 percent), oxygen-17 (0.037 percent), and oxygen-xviii (0.204 percent). Several artificially prepared radioactive isotopes are known. The longest-lived, oxygen-15 (124-2d half-life), has been used to written report respiration in mammals.

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Allotropy

Oxygen has two allotropic forms, diatomic (O₂) and triatomic (O₃, ozone). The properties of the diatomic form propose that six electrons bond the atoms and two electrons remain unpaired, accounting for the paramagnetism of oxygen. The three atoms in the ozone molecule do not lie forth a straight line.

Ozone may be produced from oxygen according to the equation: Chemical equation.

The procedure, as written, is endothermic (energy must be provided to brand information technology continue); conversion of ozone dorsum into diatomic oxygen is promoted by the presence of transition metals or their oxides. Pure oxygen is partly transformed into ozone by a silent electric discharge; the reaction is also brought nigh by absorption of ultraviolet low-cal of wavelengths around 250 nanometres (nm, the nanometre, equal to 10^−ix metre); occurrence of this process in the upper atmosphere removes radiation that would be harmful to life on the surface of the Earth. The pungent odor of ozone is noticeable in confined areas in which in that location is sparking of electric equipment, as in generator rooms. Ozone is lite blue; its density is 1.658 times that of air, and it has a boiling betoken of −112 °C (−170 °F) at atmospheric pressure.

Ozone is a powerful oxidizing agent, capable of converting sulfur dioxide to sulfur trioxide, sulfides to sulfates, iodides to iodine (providing an belittling method for its estimation), and many organic compounds to oxygenated derivatives such as aldehydes and acids. The conversion by ozone of hydrocarbons from automotive exhaust gases to these acids and aldehydes contributes to the irritating nature of smog. Commercially, ozone has been used equally a chemical reagent, as a disinfectant, in sewage treatment, h2o purification, and bleaching textiles.

Preparative methods

Production methods chosen for oxygen depend upon the quantity of the element desired. Laboratory procedures include the following:

1. Thermal decomposition of sure salts, such as potassium chlorate or potassium nitrate: Chemical equations.

The decomposition of potassium chlorate is catalyzed past oxides of transition metals; manganese dioxide (pyrolusite, MnO_two) is frequently used. The temperature necessary to effect the evolution of oxygen is reduced from 400 °C to 250 °C by the goad.

ii. Thermal decomposition of oxides of heavy metals: Chemical equations.

Scheele and Priestley used mercury(II) oxide in their preparations of oxygen.

3. Thermal decomposition of metallic peroxides or of hydrogen peroxide: Chemical equations.

An early commercial procedure for isolating oxygen from the atmosphere or for manufacture of hydrogen peroxide depended on the formation of barium peroxide from the oxide as shown in the equations.

4. Electrolysis of water containing small proportions of salts or acids to let conduction of the electric electric current: Chemical equation.

Commercial production and use

When required in tonnage quantities, oxygen is prepared by the fractional distillation of liquid air. Of the master components of air, oxygen has the highest boiling bespeak and therefore is less volatile than nitrogen and argon. The process takes advantage of the fact that when a compressed gas is allowed to aggrandize, it cools. Major steps in the functioning include the following: (1) Air is filtered to remove particulates; (2) moisture and carbon dioxide are removed by absorption in alkali; (3) the air is compressed and the heat of compression removed by ordinary cooling procedures; (four) the compressed and cooled air is passed into coils contained in a chamber; (5) a portion of the compressed air (at virtually 200 atmospheres pressure) is immune to expand in the bedchamber, cooling the coils; (half-dozen) the expanded gas is returned to the compressor with multiple subsequent expansion and pinch steps resulting finally in liquefaction of the compressed air at a temperature of −196 °C; (vii) the liquid air is immune to warm to dribble first the light rare gases, and then the nitrogen, leaving liquid oxygen. Multiple fractionations volition produce a product pure enough (99.5 pct) for near industrial purposes.

The steel manufacture is the largest consumer of pure oxygen in "blowing" high carbon steel—that is, volatilizing carbon dioxide and other nonmetal impurities in a more rapid and more easily controlled process than if air were used. The treatment of sewage by oxygen holds promise for more than efficient handling of liquid effluents than other chemical processes. Incineration of wastes in closed systems using pure oxygen has become important. The so-chosen LOX of rocket oxidizer fuels is liquid oxygen; the consumption of LOX depends upon the activity of infinite programs. Pure oxygen is used in submarines and diving bells.

Commercial oxygen or oxygen-enriched air has replaced ordinary air in the chemic industry for the industry of such oxidation-controlled chemicals as acetylene, ethylene oxide, and methanol. Medical applications of oxygen include use in oxygen tents, inhalators, and pediatric incubators. Oxygen-enriched gaseous anesthetics ensure life support during general anesthesia. Oxygen is significant in a number of industries that use kilns.

Chemical properties and reactions

The large values of the electronegativity and the electron analogousness of oxygen are typical of elements that evidence but nonmetallic behaviour. In all of its compounds, oxygen assumes a negative oxidation land equally is expected from the two half-filled outer orbitals. When these orbitals are filled by electron transfer, the oxide ion Oⁱⁱ⁻ is created. In peroxides (species containing the ion O₂ ⁱⁱ⁻) information technology is causeless that each oxygen has a charge of −one. This holding of accepting electrons by complete or partial transfer defines an oxidizing amanuensis. When such an agent reacts with an electron-altruistic substance, its own oxidation land is lowered. The change (lowering), from the zero to the −2 state in the case of oxygen, is chosen a reduction. Oxygen may be thought of every bit the "original" oxidizing agent, the nomenclature used to describe oxidation and reduction existence based upon this behaviour typical of oxygen.

As described in the section on allotropy, oxygen forms the diatomic species, O_ii, under normal conditions and, as well, the triatomic species ozone, O₃. In that location is some evidence for a very unstable tetratomic species, O₄. In the molecular diatomic grade there are ii unpaired electrons that lie in antibonding orbitals. The paramagnetic behaviour of oxygen confirms the presence of such electrons.

The intense reactivity of ozone is sometimes explained by suggesting that 1 of the three oxygen atoms is in an "atomic" state; on reacting, this cantlet is dissociated from the O₃ molecule, leaving molecular oxygen.

The molecular species, O₂, is not especially reactive at normal (ambient) temperatures and pressures. The atomic species, O, is far more reactive. The energy of dissociation (O₂ → 2O) is large at 117.2 kilocalories per mole.

Oxygen has an oxidation land of −2 in most of its compounds. It forms a big range of covalently bonded compounds, amongst which are oxides of nonmetals, such every bit h2o (H₂O), sulfur dioxide (And so₂), and carbon dioxide (CO₂); organic compounds such as alcohols, aldehydes, and carboxylic acids; common acids such as sulfuric (H₂SO₄), carbonic (H₂CO_iii), and nitric (HNO_three); and corresponding salts, such every bit sodium sulfate (Na₂And then_four), sodium carbonate (Na₂CO₃), and sodium nitrate (NaNO₃). Oxygen is present as the oxide ion, O² ^-, in the crystalline structure of solid metallic oxides such as calcium oxide, CaO. Metal superoxides, such equally potassium superoxide, KO₂, incorporate the O₂ ^- ion, whereas metallic peroxides, such every bit barium peroxide, BaO₂, comprise the O₂ ^2- ion.

Robert C. Brasted