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This article pertains to the chemical element. For other uses, see argon (disambiguation).

Argon (IPA: /ˈɑːgɒn/) is a chemical element designated by the symbol Ar. Argon has atomic number 18 and is the third element in group 18 of the periodic table (noble gases). Argon is present in the Earth's atmosphere at slightly less than 1%, making it the most common noble gas on Earth.

Characteristics

Argon and oxygen have approximately the same solubility in water and are 2.5 times more soluble in water than nitrogen.^[1] This highly stable chemical element is colorless and odorless in both its liquid and gaseous forms. Argon is inert under most conditions and forms no confirmed stable compounds at room temperature. The creation of argon hydrofluoride (HArF), a metastable compound of argon with fluorine and hydrogen, was first reported by researchers at the University of Helsinki in 2000.

Although the neutral ground-state chemical compounds of argon are presently limited to HArF, argon can form clathrates with water when atoms of it are trapped in a lattice of the water molecules. Also argon-containing ions e.g. ArH+ and excited state complexes e.g. ArF are well known. Theoretical calculations on computers have shown several argon compounds that should be stable but for which no synthesis routes are currently known.

Applications

An argon filled discharge tube glows brightly when an electric current is passed through it.

Argon is used in incandescent lighting and other applications in which diatomic nitrogen is not sufficiently inert. Argon will not react with the filament of light bulbs even at high temperatures. Other uses:

Argon is used as an inert gas shield in many forms of welding, including metal inert gas welding and tungsten inert gas welding.
as the gas of choice for the plasma used in ICP spectroscopy.
as a non-reactive blanket in the manufacture of titanium and other reactive elements.
as a protective atmosphere for growing silicon and germanium crystals.
as a gas for use in plasma globes.
as a gas for thermal insulation in energy efficient windows.
Argon-39 has been used for a number of applications, primarily ice coring. It has also been used for ground water dating.
Cryosurgery procedures such as cryoablation use liquified argon to destroy cancer cells.
Liquid argon is used in calorimetry in experimental particle physics.
Argon is used in technical scuba diving to inflate the dry suit, because it is inert and has low thermal conductivity.
Blue argon lasers are used in surgery to weld arteries, destroy tumors, and to correct eye defects.
Due to its inert qualities, it is commonly used by museum conservators to protect old materials or documents, which are prone to gradual oxidation in the presence of air. ^{[citation needed]}
Argon is used in the tires of luxury cars to protect the rubber from oxygen and to ensure quiet tires when the car moves at fast speeds. ^{[citation needed]}

Occurrence

Argon constitutes 0.934% by volume and 1.29% by mass of the Earth's atmosphere, and air is the primary raw material used by industry to produce purified argon products. Argon is isolated from air by fractionation, most commonly by cryogenic fractional distillation, a process that also produces purified nitrogen, oxygen, neon, krypton and xenon.

The Martian atmosphere in contrast contains 1.6% of argon-40 and 5 ppm of argon-36. The Mariner spaceprobe fly-by of the planet Mercury in 1973 found that Mercury has a very thin atmosphere with 70% argon, believed to result from releases of the gas as a decay product from radioactive materials on the planet. In 2005, the Huygens probe also discovered the presence of argon-40 on Titan, the largest moon of Saturn.^[2]

Compounds

Argon’s complete octet of electrons indicates full s and p subshells. This full outer energy level makes argon very stable and extremely resistant to bonding with other elements. Before 1962, argon and the other noble gases were considered to be chemically inert and unable to form compounds; however, compounds of the heavier noble gases have since been synthesized. In 2000, the first argon compounds were formed by researchers at the University of Helsinki. By shining ultraviolet light onto frozen argon containing a small amount of hydrogen fluoride, argon hydrofluoride (HArF) was formed.^[3] It is stable up to 40 kelvins.

Isotopes

The main isotopes of argon found on Earth are ⁴⁰Ar, ³⁶Ar, and ³⁸Ar. Naturally occurring ⁴⁰K with a half-life of 1.250×10⁹ years, decays to stable ⁴⁰Ar (11.2%) by electron capture and by positron emission, and also transforms to stable ⁴⁰Ca (88.8%) via beta decay. These properties and ratios are used to determine the age of rocks.

In the Earth's atmosphere, ³⁹Ar is made by cosmic ray activity, primarily with ⁴⁰Ar. In the subsurface environment, it is also produced through neutron capture by ³⁹K or alpha emission by calcium. ³⁷Ar is created from the decay of ⁴⁰Ca as a result of subsurface nuclear explosions. It has a half-life of 35 days.

External links

^ http://www.lenntech.com/Periodic-chart-elements/Ar-en.htm
^ http://www.esa.int/esaCP/SEMHB881Y3E_index_0.html
^ see http://pubs.acs.org/cen/80th/noblegases.html in its paragraph starting "Many recent findings"

[1] ttp://www.lenntech.com/Periodic-chart-elements/Ar-en.htm

[2] ttp://www.esa.int/esaCP/SEMHB881Y3E_index_0.html

[3] see http://pubs.acs.org/cen/80th/noblegases.html in its paragraph starting "Many recent findings"

[1]

[2]

[3]

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 In the Earth's atmosphere, <sup>39</sup>Ar is made by [[cosmic ray]] activity, primarily with <sup>40</sup>Ar. In the subsurface environment, it is also produced through [[neutron capture]] by <sup>39</sup>K or [[alpha decay|alpha emission]] by [[calcium]]. <sup>37</sup>Ar is created from the decay of <sup>40</sup>Ca as a result of subsurface [[nuclear testing|nuclear explosions]]. It has a half-life of 35 days.
-== References ==
-<references/>
-*[http://periodic.lanl.gov/elements/18.html Los Alamos National Laboratory &ndash; Argon]
-*[http://wwwrcamnl.wr.usgs.gov/isoig/period/ar_iig.html USGS Periodic Table - Argon]
-*Emsley, J. Nature’s Building Blocks; Oxford University Press: Oxford, NY, 2001; pp 35-39.
-*Brown, T.L.; Bursten, B.E.; LeMay, H.E. In Chemistry: The Central Science, 10th ed.; Challice, J.; Draper, P.; Folchetti, N. et al.; Eds.; Pearson Education, Inc.: Upper Saddle River, NJ, 2006; pp 276 and 289.
 == External links ==