|
|
Atomic Number: |
8 |
Melting Point: | -218.79 ºC | |
| Atomic Symbol: | O | Boiling Point: | -182.95 ºC | ||
| Atomic Weight: | 15.9994 amu | Density: | 1.429
kg/m 3 |
||
| Atomic Radius: |
66 pm |
Oxidation States: | -2 | ||
| Covalent Radius: | 73 pm | Electron Configuration: | [He]2s22p4 | ||
| van der Waals Radius: |
152 |
State of Matter: | gas (paramagnetic) |
(Gr. oxys: acid, and genes: forming) 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.
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.
The gas is colorless, odorless, and tasteless. The liquid and solid forms are a pale blue color and are strongly paramagnetic.
Molecular oxygen (O2) on Earth is thermodynamically unstable. Its initial appearance was due to the action of photosynthetic anaerobes and its abundance in later epochs has been largely facilitated by terrestrial plants, which release oxygen during photosynthesis.
Oxygen, which is very reactive, is a component of hundreds of thousands of organic compounds and combines with most elements. The only elements to escape the possibility of oxidation are a few of the noble gases.
The most famous of these oxides is of course hydrogen oxide, or water (H2O). Other well known examples include compounds of carbon and oxygen, such as carbon dioxide (CO2), alcohols (R-OH), aldehydes, (R-CHO), 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, found as a minor component of liquid oxygen. Epoxides are ethers in which the oxygen atom is part of a ring of three atoms.
Ozone's presence in the atmosphere (amounting to the equivalent of a layer 3 mm thick under ordinary pressures and temperatures) helps prevent harmful ultraviolet rays of the sun from reaching the earth's surface. Pollutants in the atmosphere may have a detrimental effect on this ozone layer. Undiluted ozone has a bluish color. Liquid ozone is bluish black and solid ozone is violet-black.
Oxygen is the third most abundant element found in the sun, and it plays a part in the carbon-nitrogen cycle, the process once 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 Borealis.
A gaseous element, oxygen forms 21% of 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, 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. Air separation plants produce about 99% of the gas, while electrolysis plants produce about 1%. Liquid oxygen is usually obtained by the fractional distillation of liquid air.
Plants and animals rely on oxygen for respiration. Hospitals frequently prescribe oxygen for patients with respiratory ailments. Oxygen finds considerable use as an oxidizer, with liquid oxygen being used as an oxidizer in rocket propulsion. It is used in welding, and in the making of steel and methanol.
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.
Oxygen has three stable isotopes and ten known radioactive isotopes. The radioisotopes all have half lives of less than three minutes. Natural oxygen is a mixture of three isotopes.
Natural occurring oxygen-18 is stable and available commercially, as is water (H2O with 15% 18O). Commercial oxygen consumption in the U.S. is estimated at 20 million short tons per year and the demand is expected to increase substantially.
Certain derivatives of oxygen, such as ozone (O3), hydrogen peroxide, hydroxyl radicals and superoxide, are also highly toxic. Ozone is toxic and exposure should not exceed 0.2 mg/m 3 (8-hour time-weighted average - 40-hour work week). The body has developed mechanisms to protect against these toxic species. For instance, the naturally-occurring glutathione can act as an antioxidant, as can bilirubin which is normally a breakdown product of hemoglobin. Highly concentrated sources of oxygen promote rapid combustion and therefore are fire and explosion hazards in the presence of fuels. This is true as well of compounds of oxygen such as chlorates, perchlorates, dichromates, etc. Compounds with a high oxidative potential can often cause chemical burns.
Oxygen can be toxic at elevated partial pressures. The fire that killed the Apollo 1crew on a test launch pad spread so rapidly because the pure oxygen atmosphere was at normal atmospheric pressure instead of the one third pressure that would be used during an actual launch.
