Friendship

Friendship is having a good relationship with someone, who we call a friend. People in a friendship have respect and affection for each other. A friendship is different from a romantic relationship; friends do not necessarily love each other romantically.

A friend is a person that someone likes or knows. People who are friends talk to each other and spend time together. They also help each other when they are in trouble or are hurt. Friends are people that can be looked up to and trusted. Sometimes people can’t tell their secrets to their parents so they tell it to their friends. Usually friends have similar interests. A friend is one who admires a person’s skill and helps or encourages them to make the right choices.

The strength of the bond of friendship between two people can vary. If the bond is very strong, they are called best friends. This can usually be achieved by possessing the elements of friendship, by being kind, generous, loyal, honest and by having fun. With these qualities you can truly enjoy the bliss of friendship. There are expectations, demands and complaints in friendship too. Some people say these should not be there in friendship. But two people interact in this world only when there are expectations. Friends can help you materialistically and morally. But if they don’t match with your expectations, it doesn’t mean you will break the friendship. This is what is different in friendship as compared to love relationship between two couples. Its all about understanding, realizing and helping whenever possible.

Amish Friendship Bread

Amish friendship bread is a type of bread or cake made from a sour dough starter that is often shared in a manner similar to a chain letter. The starter is a replacement for baking yeast. It can be used to make many kinds of yeast-based breads, shared with friends, or frozen for future use. The sweet, cake-like Amish cinnamon bread is a common bread that is made from this starter. It is a simple, stirred quick bread It includes a large amount of sugar and vegetable oil, with a mild cinnamon flavour It has characteristics of both pound cake and coffee cake. The flavour of the finished product can be changed by not using cinnamon.

A common recipe using this starter uses one cup (240 ml) of it to make bread. A second cup is saved to start a new cycle. The remaining two cups are given to friends. The process of sharing the starter makes it somewhat like a chain letter. One cup of starter makes one standard loaf of bread.

It can also be called German friendship cake, or “Herman” the German friendship cake.

The starter mix can be made from a recipe of flour, water and yeast. One cup of the mixture is then passed to other people who get instructions on how to keep it alive, usually just saying to give it more ingredients as time passes.

Synopsis: https://en.wikipedia.org/wiki/Friendship and https://en.wikipedia.org/wiki/Amish_friendship_bread.

Xenon is a chemical element with the symbol Xe and atomic number 54. It is a colourless, heavy, odourless noble gas, that occurs in the Earth’s atmosphere in trace amounts. Although generally unreactive, xenon can undergo a few chemical reactions such as the formation of xenon hexafluoroplatinate, the first noble gas compound to be synthesized.

Naturally occurring xenon consists of eight stable isotopes. There are also over 40 unstable isotopes that undergo radioactive decay. The isotope ratios of xenon are an important tool for studying the early history of the Solar System. Radioactive xenon-135 is produced from iodine-135 as a result of nuclear fission, and it acts as the most significant neutron absorber in nuclear reactors.

Xenon is used in flash lamps and arc lamps, and as a general aesthetic. The first excimer laser design used a xenon dimer molecule (Xe2) as its lasing medium, and the earliest laser designs used xenon flash lamps as pumps. Xenon is also being used to search for hypothetical weakly interacting massive particles and as the propellant for ion thrusters in spacecraft.

History

Xenon was discovered in England by the Scottish chemist William Ramsay and English chemist Morris Travers on July 12, 1898, shortly after their discovery of the elements krypton and neon. They found xenon in the residue left over from evaporating components of liquid air. Ramsay suggested the name xenon for this gas from the Greek word ξένον [xenon], neuter singular form of ξένος [xenos], meaning ‘foreign(er)’, ‘strange(r)’, or ‘guest’. In 1902, Ramsay estimated the proportion of xenon in the Earth’s atmosphere as one part in 20 million. The current symbol for Xenon is Xe, however historically it was also written as X.

During the 1930s, American engineer Harold Edgerton began exploring strobe light technology for high speed photography. This led him to the invention of the xenon flash lamp, in which light is generated by sending a brief electrical current through a tube filled with xenon gas. In 1934, Edgerton was able to generate flashes as brief as one microsecond with this method.

In 1939, American physician Albert R. Behnke Jr. began exploring the causes of “drunkenness” in deep-sea divers. He tested the effects of varying the breathing mixtures on his subjects, and discovered that this caused the divers to perceive a change in depth. From his results, he deduced that xenon gas could serve as an aesthetic. Although Russian toxicologist Nikolay V. Lazarev apparently studied xenon anaesthesia in 1941, the first published report confirming xenon anaesthesia was in 1946 by American medical researcher John H. Lawrence, who experimented on mice. Xenon was first used as a surgical aesthetic. in 1951 by American anaesthesiologist Stuart C. Cullen, who successfully operated on two patients.

Xenon and the other noble gases were for a long time considered to be completely chemically inert and not able to form compounds. However, while teaching at the University of British Columbia, Neil Bartlett discovered that the gas platinum hexafluoride (PtF6) was a powerful oxidizing agent that could oxidise oxygen gas (O2) to form dioxygenyl hexafluoroplatinate (O2+[PtF6]–). Since O2 and xenon have almost the same first ionization potential, Bartlett realized that platinum hexafluoride might also be able to oxidise xenon. On March 23, 1962, he mixed the two gases and produced the first known compound of a noble gas, xenon hexafluoroplatinate. Bartlett thought its composition to be Xe+[PtF6]–, although later work has revealed that it was probably a mixture of various xenon-containing salts. Since then, many other xenon compounds have been discovered, along with some compounds of the noble gases argon, krypton, and radon, including argon fluorohydride (HArF), krypton difluoride (KrF2), and radon fluoride. By 1971, more than 80 xenon compounds were known.

Characteristics

Xenon has atomic number 54; that is, its nucleus contains 54 protons. At standard temperature and pressure, pure xenon gas has a density of 5.761 kg/m3, about 4.5 times the surface density of the Earth’s atmosphere, 1.217 kg/m3. As a liquid, xenon has a density of up to 3.100 g/mL, with the density maximum occurring at the triple point. Under the same conditions, the density of solid xenon, 3.640 g/cm3, is higher than the average density of granite, 2.75 g/cm3. Using gigapascals of pressure, xenon has been forced into a metallic phase.

Solid xenon changes from face-centered cubic (fcc) to hexagonal close packed (hcp) crystal phase under pressure and begins to turn metallic at about 140 GPa, with no noticeable volume change in the hcp phase. It is completely metallic at 155 GPa. When metalized, xenon looks sky blue because it absorbs red light and transmits other visible frequencies. Such behaviour is unusual for a metal and is explained by the relatively small widths of the electron bands in metallic xenon.

Xenon is a member of the zero-valence elements that are called noble or inert gases. It is inert to most common chemical reactions (such as combustion, for example) because the outer valence shell contains eight electrons. This produces a stable, minimum energy configuration in which the outer electrons are tightly bound. However, xenon can be oxidized by powerful oxidizing agents, and many xenon compounds have been synthesized.

In a gas-filled tube, xenon emits a blue or lavenderish glow when the gas is excited by electrical discharge. Xenon emits a band of emission lines that span the visual spectrum, but the most intense lines occur in the region of blue light, which produces the colouration.

Occurrence and production

Xenon is a trace gas in Earth’s atmosphere, occurring at 87±1 parts per billion (nL/L), or approximately 1 part per 11.5 million, and is also found in gases emitted from some mineral springs.

Xenon is obtained commercially as a by-product of the separation of air into oxygen and nitrogen. After this separation, generally performed by fractional distillation in a double-column plant, the liquid oxygen produced will contain small quantities of krypton and xenon. By additional fractional distillation steps, the liquid oxygen may be enriched to contain 0.1–0.2% of a krypton/xenon mixture, which is extracted either via adsorption onto silica gel or by distillation. Finally, the krypton/xenon mixture may be separated into krypton and xenon via distillation. Extraction of a litre of xenon from the atmosphere requires 220 watt-hours of energy.

Worldwide production of xenon in 1998 was estimated at 5,000–7,000 m3. Because of its low abundance, xenon is much more expensive than the lighter noble gases—approximate prices for the purchase of small quantities in Europe in 1999 were 10 €/L for xenon, 1 €/L for krypton, and 0.20 €/L for neon; the much more plentiful argon costs less than a cent per litre.

Within the Solar System, the nucleon fraction of xenon is 1.56 × 10−8, for an abundance of approximately one part in 630 thousand of the total mass. Xenon is relatively rare in the Sun’s atmosphere, on Earth, and in asteroids and comets. The planet Jupiter has an unusually high abundance of xenon in its atmosphere; about 2.6 times as much as the Sun. This high abundance remains unexplained and may have been caused by an early and rapid build-up of planetesimals—small, sub planetary bodies—before the presolar disk began to heat up. (Otherwise, xenon would not have been trapped in the planetesimal ices.) The problem of the low terrestrial xenon may potentially be explained by covalent bonding of xenon to oxygen within quartz, hence reducing the out gassing of xenon into the atmosphere.

Unlike the lower mass noble gases, the normal stellar nucleosynthesis process inside a star does not form xenon. Elements more massive than iron-56 have a net energy cost to produce through fusion, so there is no energy gain for a star when creating xenon. Instead, xenon is formed during supernova explosions, by the slow neutron capture process ( s-process) of red giant stars that have exhausted the hydrogen at their cores and entered the asymptotic giant branch, in classical nova explosions and from the radioactive decay of elements such as iodine, uranium and plutonium.

Isotopes and isotopic studies

Naturally occurring xenon is made of eight stable isotopes, the most of any element with the exception of tin, which has ten. Xenon and tin are the only elements to have more than seven stable isotopes. The isotopes 124Xe and 134Xe are predicted to undergo double beta decay, but this has never been observed so they are considered to be stable. Besides these stable forms, there are over 40 unstable isotopes that have been studied. The longest lived of these isotopes is 136Xe, which has been observed to undergo double beta decay with a half-life of 2.11 x 1021yr. 129Xe is produced by beta decay of 129I, which has a half-life of 16 million years, while 131mXe, 133Xe, 133mXe, and 135Xe are some of the fission products of both 235U and 239Pu, and therefore used as indicators of nuclear explosions.

Nuclei of two of the stable isotopes of xenon, 129Xe and 131Xe, have non-zero intrinsic angular momenta ( nuclear spins, suitable for nuclear magnetic resonance). The nuclear spins can be aligned beyond ordinary polarization levels by means of circularly polarized light and rubidium vapour. The resulting spin polarization of xenon nuclei can surpass 50% of its maximum possible value, greatly exceeding the equilibrium value dictated by the Boltzmann distribution (typically 0.001% of the maximum value at room temperature, even in the strongest magnets). Such non-equilibrium alignment of spins is a temporary condition, and is called hyper polarization. The process of hyper polarizing the xenon is called optical pumping (although the process is different from pumping a laser).

Because a 129Xe nucleus has a spin of 1/2, and therefore a zero electric quadrupole moment, the 129Xe nucleus does not experience any quadrupolar interactions during collisions with other atoms, and thus its hyper polarization. can be maintained for long periods of time even after the laser beam has been turned off and the alkali vapour. removed by condensation on a room-temperature surface. Spin polarization of 129Xe can persist from several seconds for xenon atoms dissolved in blood to several hours in the gas phase and several days in deeply frozen solid xenon. In contrast, 131Xe has a nuclear spin value of 3/2 and a non-zero quadrupole moment, and has T1 relaxation times in the millisecond and second ranges.

Some radioactive isotopes of xenon, for example, 133Xe and 135Xe, are produced by neutron irradiation of fissionable material within nuclear reactors. 135Xe is of considerable significance in the operation of nuclear fission reactors. 135Xe has a huge cross section for thermal neutrons, 2.6×106 barns, so it acts as a neutron absorber or ” poison” that can slow or stop the chain reaction after a period of operation. This was discovered in the earliest nuclear reactors built by the American Manhattan Project for plutonium production. Fortunately the designers had made provisions in the design to increase the reactor’s reactivity (the number of neutrons per fission that go on to fission other atoms of nuclear fuel). 135Xe reactor poisoning played a major role in the Chernobyl disaster. A shut-down or decrease of power of a reactor can result in build-up of 135Xe and getting the reactor into the iodine pit.

Under adverse conditions, relatively high concentrations of radioactive xenon isotopes may be found emanating from nuclear reactors due to the release of fission products from cracked fuel rods, or fissioning of uranium in cooling water.

Because xenon is a tracer for two parent isotopes, xenon isotope ratios in meteorites are a powerful tool for studying the formation of the solar system. The iodine-xenon method of dating gives the time elapsed between nucleosynthesis and the condensation of a solid object from the solar nebula. In 1960, physicist John H. Reynolds discovered that certain meteorites contained an isotopic anomaly in the form of an over-abundance of xenon-129. He inferred that this was a decay product of radioactive iodine-129. This isotope is produced slowly by cosmic ray spallation and nuclear fission, but is produced in quantity only in supernova explosions. As the half-life of 129I is comparatively short on a cosmological time scale, only 16 million years, this demonstrated that only a short time had passed between the supernova and the time the meteorites had solidified and trapped the 129I. These two events (supernova and solidification of gas cloud) were inferred to have happened during the early history of the Solar System, as the 129I isotope was likely generated before the Solar System was formed, but not long before, and seeded the solar gas cloud with isotopes from a second source. This supernova source may also have caused collapse of the solar gas cloud.

In a similar way, xenon isotopic ratios such as 129Xe/130Xe and 136Xe/130Xe are also a powerful tool for understanding planetary differentiation and early out gassing. For example, The atmosphere of Mars shows a xenon abundance similar to that of Earth: 0.08 parts per million, however Mars shows a higher proportion of 129Xe than the Earth or the Sun. As this isotope is generated by radioactive decay, the result may indicate that Mars lost most of its primordial atmosphere, possibly within the first 100 million years after the planet was formed. In another example, excess 129Xe found in carbon dioxide well gases from New Mexico was believed to be from the decay of mantle-derived gases soon after Earth’s formation.

Compounds

After Neil Bartlett’s discovery in 1962 that xenon can form chemical compounds, a large number of xenon compounds have been discovered and described. Almost all known xenon compounds contain the electronegative atoms fluorine or oxygen.

Halides

Three fluorides are known: XeF2, XeF4, and XeF6. XeF is theorized to be unstable. The fluorides are the starting point for the synthesis of almost all xenon compounds.

The solid, crystalline difluoride XeF2 is formed when a mixture of fluorine and xenon gases is exposed to ultraviolet light. Ordinary daylight is sufficient. Long-term heating of XeF2 at high temperatures under an NiF2 catalyst yields XeF6. Pyrolysis of XeF6 in the presence of NaF yields high-purity XeF4.

The xenon fluorides behave as both fluoride acceptors and fluoride donors, forming salts that contain such cations as XeF+ and Xe

2F+

3, and anions such as XeF−5, XeF−7, and XeF2−8. The green, paramagnetic Xe+

2 is formed by the reduction of XeF2 by xenon gas.

XeF2 is also able to form coordination complexes with transition metal ions. Over 30 such complexes have been synthesized and characterized.

Whereas the xenon fluorides are well-characterized, the other halides are not known, the only exception being the dichloride, XeCl2. Xenon dichloride is reported to be an endothermic, colourless, crystalline compound that decomposes into the elements at 80°C, formed by the high-frequency irradiation of a mixture of xenon, fluorine, and silicon or carbon tetrachloride. However, doubt has been raised as to whether XeCl2 is a real compound and not merely a van der Waals molecule consisting of weakly bound Xe atoms and Cl2 molecules. Theoretical calculations indicate that the linear molecule XeCl2 is less stable than the van der Waals complex.

Oxides and oxohalides

Three oxides of xenon are known: xenon trioxide (XeO3) and xenon tetroxide (XeO4), both of which are dangerously explosive and powerful oxidizing agents, and xenon dioxide (XeO2), which was reported in 2011 with a coordination number of four. XeO2 forms when xenon tetrafluoride is poured over ice. Its crystal structure may allow it to replace silicon in silicate minerals. The XeOO+ cation has been identified by infra-red spectroscopy in solid argon.

Xenon does not react with oxygen directly; the trioxide is formed by the hydrolysis of XeF6: XeF6 + 3 H2O → XeO3 + 6 HF

XeO3 is weakly acidic, dissolving in alkali to form unstable xenate salts containing the HXeO−4 anion. These unstable salts easily disproportionate into xenon gas and perxenate salts, containing the XeO4−6 anion.

Barium perxenate, when treated with concentrated sulphuric acid, yields gaseous xenon tetroxide:

Ba2XeO6 + 2 H2SO4 → 2 BaSO4 + 2 H2O + XeO4

To prevent decomposition, the xenon tetroxide thus formed is quickly cooled to form a pale-yellow solid. It explodes above −35.9 °C into xenon and oxygen gas.

A number of xenon oxyfluorides are known, including XeOF2, XeOF4, XeO2F2, and XeO3F2. XeOF2 is formed by the reaction of OF2 with xenon gas at low temperatures. It may also be obtained by the partial hydrolysis of XeF4. It disproportionates at −20 °C into XeF2 and XeO2F2. XeOF4 is formed by the partial hydrolysis of XeF6, or the reaction of XeF6 with sodium perxenate, Na4XeO6. The latter reaction also produces a small amount of XeO3F2. XeOF4 reacts with CsF to form the XeOF−5 anion, while XeOF3 reacts with the alkali metal fluorides KF, RbF and CsF to form the XeOF−4 anion.

Synopsis of https://en.wikipedia.org/wiki/Xenon

Other compounds

Recently, there has been an interest in xenon compounds where xenon is directly bonded to a less electronegative element than fluorine or oxygen, particularly carbon. Electron-withdrawing groups, such as groups with fluorine substitution, are necessary to stabilize these compounds.

Numerous such compounds have been characterized, including:

• C6F5–Xe+–N≡C–CH3, where C6F5 is the pentafluorophenyl group.

• [C6F5]2Xe
• C6F5–Xe–X, where X is CN, F, or Cl.
• R–C≡C–Xe+, where R is C2F−5 or tert-butyl.
• C6F5–XeF+2
• (C6F5Xe)2Cl+

Other compounds containing xenon bonded to a less electronegative element include F–Xe–N(SO2F)2 and F–Xe–BF2. The latter is synthesized from dioxygenyl tetrafluoroborate, O2BF4, at −100 °C.

An unusual ion containing xenon is the tetraxenonogold(II) cation, AuXe2+4, which contains Xe–Au bonds. This ion occurs in the compound AuXe4(Sb2F11)2, and is remarkable in having direct chemical bonds between two notoriously non-reactive atoms, xenon and gold, with xenon acting as a transition metal ligand.

In 1995, M. Räsänen and co-workers, scientists at the University of Helsinki in Finland, announced the preparation of xenon dihydride (HXeH), and later xenon hydride-hydroxide (HXeOH), hydroxenoacetylene (HXeCCH), and other Xe-containing molecules. In 2008, Khriachtchev et al. reported the preparation of HXeOXeH by the photolysis of water within a cryogenic xenon matrix. Deuterated molecules, HXeOD and DXeOH, have also been produced.

Clathrates and excimers

In addition to compounds where xenon forms a chemical bond, xenon can form clathrates—substances where xenon atoms are trapped by the crystalline lattice of another compound. An example is xenon hydrate (Xe•5.75 H2O), where xenon atoms occupy vacancies in a lattice of water molecules. This clathrate has a melting point of 24 °C. The deuterated version of this hydrate has also been produced. Such clathrate hydrates can occur naturally under conditions of high pressure, such as in Lake Vostok underneath the Antarctic ice sheet. Clathrate formation can be used to fractionally distil xenon, argon and krypton.

Xenon can also form endohedral fullerene compounds, where a xenon atom is trapped inside a fullerene molecule. The xenon atom trapped in the fullerene can be monitored via 129Xe nuclear magnetic resonance (NMR) spectroscopy. Using this technique, chemical reactions on the fullerene molecule can be analyzed, due to the sensitivity of the chemical shift of the xenon atom to its environment. However, the xenon atom also has an electronic influence on the reactivity of the fullerene.

While xenon atoms are at their ground energy state, they repel each other and will not form a bond. When xenon atoms becomes energized, however, they can form an excimer (excited dimer) until the electrons return to the ground state. This entity is formed because the xenon atom tends to fill its outermost electronic shell, and can briefly do this by adding an electron from a neighbouring xenon atom. The typical lifetime of a xenon excimer is 1–5 ns, and the decay releases photons with wavelengths of about 150 and 173 nm. Xenon can also form excimers with other elements, such as the halogens bromine, chlorine and fluorine.

Applications

Although xenon is rare and relatively expensive to extract from the Earth’s atmosphere, it has a number of applications.

Illumination and optics

Gas-discharge lamps

Xenon is used in light-emitting devices called xenon flash lamps, which are used in photographic flashes and stroboscopic lamps; to excite the active medium in lasers which then generate coherent light; and, occasionally, in bactericidal lamps. The first solid-state laser, invented in 1960, was pumped by a xenon flash lamp, and lasers used to power inertial confinement fusion are also pumped by xenon flash lamps.

Continuous, short-arc, high pressure xenon arc lamps have a colour temperature closely approximating noon sunlight and are used in solar simulators. That is, the chromaticity of these lamps closely approximates a heated black body radiator that has a temperature close to that observed from the Sun. After they were first introduced during the 1940s, these lamps began replacing the shorter-lived carbon arc lamps in movie projectors. They are employed in typical 35mm, IMAX and the new digital projectors film projection systems, automotive HID headlights, high-end “tactical” flash lights and other specialized uses. These arc lamps are an excellent source of short wavelength ultraviolet radiation and they have intense emissions in the near infra-red, which is used in some night vision systems.

The individual cells in a plasma display use a mixture of xenon and neon that is converted into a plasma using electrodes. The interaction of this plasma with the electrodes generates ultraviolet photons, which then excite the phosphor coating on the front of the display.

Xenon is used as a “starter gas” in high pressure sodium lamps. It has the lowest thermal conductivity and lowest ionization potential of all the non-radioactive noble gases. As a noble gas, it does not interfere with the chemical reactions occurring in the operating lamp. The low thermal conductivity minimizes thermal losses in the lamp while in the operating state, and the low ionization potential causes the breakdown voltage of the gas to be relatively low in the cold state, which allows the lamp to be more easily started.

Lasers

In 1962, a group of researchers at Bell Laboratories discovered laser action in xenon, and later found that the laser gain was improved by adding helium to the lasing medium. The first excimer laser used a xenon dimer (Xe2) energized by a beam of electrons to produce stimulated emission at an ultraviolet wavelength of 176 nm. Xenon chloride and xenon fluoride have also been used in excimer (or, more accurately, exciplex) lasers. The xenon chloride excimer laser has been employed, for example, in certain dermatological uses.

Medical

Anaesthesia

Xenon has been used as a general aesthetic. Although it is expensive, anaesthesia machines that can deliver xenon are about to appear on the European market, because advances in recovery and recycling of xenon have made it economically viable.

Xenon interacts with many different receptors and ion channels and like many theoretically multi-modal inhalation aesthetics these interactions are likely complementary. Xenon is a high-affinity glycine-site NMDA receptor antagonist. However, xenon distinguishes itself from other clinically used NMDA receptor antagonists in its lack of neurotoxicity and ability its to inhibit the neurotoxicity of ketamine and nitrous oxide. Unlike ketamine and nitrous oxide, xenon does not stimulate a dopamine efflux from the nucleus accumbens. Like nitrous oxide and cyclopropane xenon activates the two-pore domain potassium channel TREK-1. A related channel TASK-3 also implicated in aesthetic. actions is insensitive to xenon. Xenon inhibits nicotinic acetylcholine alpha4beta2 receptors which contribute to spinally mediated analgesia. Xenon is an effective inhibitor of plasma membrane Ca2+ ATPase. Xenon inhibits Ca+ ATPase by binding to a hydrophobic pore within the enzyme and preventing the enzyme from assuming active conformations.

Xenon is a competitive inhibitor of serotonin 5HT3. While neither aesthetic. nor antinociceptive this activity reduces anesthesia-emergent nausea and vomiting.

Xenon has a minimum alveolar concentration (MAC) of 72% at age 40, making it 44% more potent than N2O as an aesthetic. Thus it can be used in concentrations with oxygen that have a lower risk of hypoxia. Unlike nitrous oxide (N2O), xenon is not a greenhouse gas and so it is also viewed as environmentally friendly. Xenon vented into the atmosphere is being returned to its original source, so no environmental impact is likely.

Neuroprotectant

Xenon induces robust cardioprotection and neuroprotection through several a variety of mechanisms of action. Through its influence on Ca2+, K+, KATP\HIF and NMDA antagonism xenon is neuroprotective when administered before during & after ischemic insults. Xenon is a high affinity antagonist at the NMDA receptor glycine site. Xenon is cardioprotective in ischemia-reperfusion conditions by inducing pharmacologic non-ischemic preconditioning. Xenon is cardioprotective by activating PKC-epsilon & downstream p38-MAPK. Xenon mimics neuronal ischemic preconditioning by activating ATP sensitive potassium channels.

Xenon allosterically reduces ATP mediated channel activation inhibition independently of the sulfonylurea receptor1 subunit, increasing KATP open-channel time and frequency. Xenon upregulates hypoxia inducible factor 1 alpha (HIF1a).

Xenon gas was added as an ingredient of the ventilation mix for a newborn baby at St. Michael’s Hospital, Bristol, England, whose life chances were otherwise very compromised, and was successful, leading to the authorisation of clinical trials for similar cases. The treatment is done simultaneously with cooling the body temperature to 33.5 °C.

Imaging

Gamma emission from the radioisotope 133Xe of xenon can be used to image the heart, lungs, and brain, for example, by means of single photon emission computed tomography. 133Xe has also been used to measure blood flow.

Xenon, particularly hyperpolarized 129Xe, is a useful contrast agent for magnetic resonance imaging (MRI). In the gas phase, it can be used to image empty space such as cavities in a porous sample or alveoli in lungs. Hyperpolarization renders 129Xe much more detectable via magnetic resonance imaging and has been used for studies of the lungs and other tissues. It can be used, for example, to trace the flow of gases within the lungs. Because xenon is soluble in water and also in hydrophobic solvents, it can be used to image various soft living tissues.

NMR spectroscopy

Because of the atom’s large, flexible outer electron shell, the NMR spectrum changes in response to surrounding conditions, and can therefore be used as a probe to measure the chemical circumstances around the xenon atom. For instance xenon dissolved in water, xenon dissolved in hydrophobic solvent, and xenon associated with certain proteins can be distinguished by NMR.

Hyperpolarized xenon can be used by surface-chemists. Normally, it is difficult to characterize surfaces using NMR, because signals from the surface of a sample will be overwhelmed by signals from the far-more-numerous atomic nuclei in the bulk. However, nuclear spins on solid surfaces can be selectively polarized, by transferrering spin polarization to them from hyperpolarized xenon gas. This makes the surface signals strong enough to measure, and distinguishes them from bulk signals.

Other

In nuclear energy applications, xenon is used in bubble chambers, probes, and in other areas where a high molecular weight and inert nature is desirable. A by-product of nuclear weapon testing is the release of radioactive xenon-133 and xenon-135. The detection of these isotopes is used to monitor compliance with nuclear test ban treaties, as well as to confirm nuclear test explosions by states such as North Korea.

Liquid xenon is being used in calorimeters for measurements of gamma rays as well as a medium for detecting hypothetical weakly interacting massive particles, or WIMPs. When a WIMP collides with a xenon nucleus, it should, theoretically, strip an electron and create a primary scintillation. By using xenon, this burst of energy could then be readily distinguished from similar events caused by particles such as cosmic rays.

However, the XENON experiment at the Gran Sasso National Laboratory in Italy and the ZEPLIN-II and ZEPLIN-III experiments at the Boulby Underground Laboratory in the UK have thus far failed to find any confirmed WIMPs. Even if no WIMPs are detected, the experiments will serve to constrain the properties of dark matter and some physics models. The current detector at the Gran Sasso facility has demonstrated sensitivity comparable to that of the best cryogenic detectors, and the sensitivity was expected to be increased by an order of magnitude in 2009.

Xenon is the preferred propellant for ion propulsion of spacecraft because of its low ionization potential per atomic weight, and its ability to be stored as a liquid at near room temperature (under high pressure) yet be easily converted back into a gas to feed the engine. The inert nature of xenon makes it environmentally friendly and less corrosive to an ion engine than other fuels such as mercury or caesium. Xenon was first used for satellite ion engines during the 1970s. It was later employed as a propellant for JPL’s Deep Space 1 probe, Europe’s SMART-1 spacecraft and for the three ion propulsion engines on NASA’s Dawn Spacecraft.

Chemically, the perxenate compounds are used as oxidizing agents in analytical chemistry. Xenon difluoride is used as an etchant for silicon, particularly in the production of microelectromechanical systems (MEMS). The anticancer drug 5-fluorouracil can be produced by reacting xenon difluoride with uracil. Xenon is also used in protein crystallography.

Applied at pressures from 0.5 to 5 MPa (5 to 50 atm) to a protein crystal, xenon atoms bind in predominantly hydrophobic cavities, often creating a high quality, isomorphous, heavy-atom derivative, which can be used for solving the phase problem.

Precautions

Many oxygen-containing xenon compounds are toxic due to their strong oxidative properties, and explosive due to their tendency to break down into elemental xenon plus diatomic oxygen (O2), which contains much stronger chemical bonds than the xenon compounds.

Xenon gas can be safely kept in normal sealed glass or metal containers at standard temperature and pressure. However, it readily dissolves in most plastics and rubber, and will gradually escape from a container sealed with such materials. Xenon is non- toxic, although it does dissolve in blood and belongs to a select group of substances that penetrate the blood–brain barrier, causing mild to full surgical anaesthesia when inhaled in high concentrations with oxygen.

At 169 m/s, the speed of sound in xenon gas is slower than that in air due to the slower average speed of the heavy xenon atoms compared to nitrogen and oxygen molecules. Hence, xenon lowers the resonant frequencies of the vocal tract when inhaled. This produces a characteristic lowered voice timbre, an effect opposite to the high-timbred voice caused by inhalation of helium. Like helium, xenon does not satisfy the body’s need for oxygen. Xenon is both a simple asphyxiant and an aesthetic. more powerful than nitrous oxide; consequently, many universities no longer allow the voice stunt as a general chemistry demonstration. As xenon is expensive, the gas sulphur hexafluoride, which is similar to xenon in molecular weight (146 versus 131), is generally used in this stunt, and is an asphyxiant without being aesthetic.

It is possible to safely breathe heavy gases such as xenon or sulphur hexafluoride when they are in a mixture with oxygen; the oxygen comprising at least 20% of the mixture. Xenon at 80% concentration along with 20% oxygen rapidly produces the unconsciousness of general anaesthesia (and has been used for this, as discussed above). Breathing mixes gases of different densities very effectively and rapidly so that heavier gases are purged along with the oxygen, and do not accumulate at the bottom of the lungs. There is, however, a danger associated with any heavy gas in large quantities: it may sit invisibly in a container, and if a person enters a container filled with an odourless, colourless gas, they may find themselves breathing it unknowingly. Xenon is rarely used in large enough quantities for this to be a concern, though the potential for danger exists any time a tank or container of xenon is kept in an unventilated space.

Synopsis of https://en.wikipedia.org/wiki/Xenon

Krypton (from Greek: κρυπτός kryptos “the hidden one”) is a chemical element with symbol Kr and atomic number 36. It is a member of group 18 (noble gases) elements. A colourless, odourless, tasteless noble gas, krypton occurs in trace amounts in the atmosphere, is isolated by fractionally distilling liquified air, and is often used with other rare gases in fluorescent lamps. Krypton is inert for most practical purposes.

Krypton, like the other noble gases, can be used in lighting and photography. Krypton light has a large number of spectral lines, and krypton’s high light output in plasmas allows it to play an important role in many high-powered gas lasers (krypton ion and excimer lasers), which pick out one of the many spectral lines to amplify. There is also a specific krypton fluoride laser. The high power and relative ease of operation of krypton discharge tubes caused (from 1960 to 1983) the official length of a meter to be defined in terms of the 605 nm (red-orange) spectral line of krypton-86.

 

History

Krypton was discovered in Britain in 1898 by Sir William Ramsay, a Scottish chemist, and Morris Travers, an English chemist, in residue left from evaporating nearly all components of liquid air. Neon was discovered by a similar procedure by the same workers just a few weeks later.

 

William Ramsay was awarded the 1904 Nobel Prize in Chemistry for discovery of a series of noble gases, including krypton.

 

In 1960, an international agreement defined the meter in terms of wavelength of light emitted by the krypton-86 isotope (wavelength of 605.78 nanometres). This agreement replaced the long-standing standard meter located in Paris, which was a metal bar made of a platinum-iridium alloy (the bar was originally estimated to be one ten-millionth of a quadrant of the Earth’s polar circumference), and was itself replaced by a definition based on the speed of light — a fundamental physical constant. However, in 1927, the International Conference on Weights and Measures had redefined the meter in terms of a red cadmium spectral line (1 m = 1,553,164.13 wavelengths). In October 1983, the same bureau defined the meter as the distance that light travels in a vacuum during 1/299,792,458 s.

 

Characteristics

Krypton is characterized by several sharp emission lines (spectral signatures) the strongest being green and yellow. It is one of the products of uranium fission. Solidified krypton is white and crystalline with a face-centred cubic crystal structure, which is a common property of all noble gases (except helium, with a hexagonal close-packed crystal structure).

 

Isotopes

Naturally occurring krypton is made of six stable isotopes. In addition, about thirty unstable isotopes and isomers are known. 81Kr, the product of atmospheric reactions, is produced with the other naturally occurring isotopes of krypton. Being radioactive, it has a half-life of 230,000 years. Krypton is highly volatile when it is near surface waters but 81Kr has been used for dating old (50,000–800,000 years) groundwater.

 

85Kr is an inert radioactive noble gas with a half-life of 10.76 years. It is produced by the fission of uranium and plutonium, such as in nuclear bomb testing and nuclear reactors. 85Kr is released during the reprocessing of fuel rods from nuclear reactors. Concentrations at the North Pole are 30% higher than at the South Pole due to convective mixing.

 

Chemistry

Like the other noble gases, krypton is chemically non-reactive. However, following the first successful synthesis of xenon compounds in 1962, synthesis of krypton di-fluoride (KrF2) was reported in 1963. In the same year, KrF4 was reported by Grosse, et al., but was subsequently shown to be a mistaken identification. There are also unverified reports of a barium salt of a krypton oxoacid. ArKr+ and KrH+ polyatomic ions have been investigated and there is evidence for KrXe or KrXe+.

 

Compounds with krypton bonded to atoms other than fluorine have also been discovered. The reaction of KrF2 with B(OTeF5)3 produces an unstable compound, Kr(OTeF5)2, that contains a krypton-oxygen bond. A krypton-nitrogen bond is found in the cation [HC≡N–Kr–F]+, produced by the reaction of KrF2 with [HC≡NH]+[AsF−6] below −50 °C. HKrCN and HKrC≡CH (krypton hydride-cyanide and hydro-krypto-acetylene) were reported to be stable up to 40 K.

 

Natural occurrence

The Earth has retained all of the noble gases that were present at its formation except for helium. Krypton’s concentration in the atmosphere is about 1 ppm. It can be extracted from liquid air by fractional distillation. The amount of krypton in space is uncertain, as the amount is derived from the meteoric activity and that from solar winds. The first measurements suggest an over-abundance of krypton in space.

 

Applications

Krypton’s multiple emission lines make ionized krypton gas discharges appear whitish, which in turn makes krypton-based bulbs useful in photography as a brilliant white light source. Krypton is thus used in some types of photographic flashes used in high speed photography. Krypton gas is also combined with other gases to make luminous signs that glow with a bright greenish-yellow light.

 

Krypton mixes with argon as the fill gas of energy saving fluorescent lamps. This reduces their power consumption. Unfortunately this also reduces their light output and raises their cost. Krypton costs about 100 times as much as argon. Krypton (along with xenon) is also used to fill incandescent lamps to reduce filament evaporation and allow higher operating temperatures to be used for the filament. A brighter light results which contains more blue than conventional lamps.

 

Krypton’s white discharge is often used to good effect in coloured gas discharge tubes, which are then simply painted or stained in other ways to allow the desired colour (for example, “neon” type advertising signs where the letters appear in differing colours are often entirely krypton-based). Krypton is also capable of much higher light power density than neon in the red spectral line region, and for this reason, red lasers for high-power laser light-shows are often krypton lasers with mirrors which select out the red spectral line for laser amplification and emission, rather than the more familiar helium-neon variety, which could never practically achieve the multi-watt red laser light outputs needed for this application.

 

Krypton has an important role in production and usage of the krypton fluoride laser. The laser has been important in the nuclear fusion energy research community in confinement experiments. The laser has high beam uniformity, short wavelength, and the ability to modify the spot size to track an imploding pellet.

 

In experimental particle physics, liquid krypton is used to construct quasi-homogeneous electromagnetic calorimeters. A notable example is the calorimeter of the NA48 experiment at CERN containing about 27 tonnes of liquid krypton. This usage is rare, since the cheaper liquid argon is typically used. The advantage of krypton over argon is a small Molière radius of 4.7 cm, which allows for excellent spatial resolution and low degree of overlapping. The other parameters relevant for calorimetry application are: radiation length of X0=4.7 cm, density of 2.4 g/cm3.

 

The sealed spark gap assemblies contained in ignition exciters used in some older jet engines contain a very small amount of Krypton-85 to obtain consistent ionization levels and uniform operation.

 

Krypton-83 has application in magnetic resonance imaging (MRI) for imaging airways. In particular, it may be used to distinguish between hydrophobic and hydrophilic surfaces containing an airway.

 

Although xenon has potential for use in computed tomography (CT) to assess regional ventilation, its anaesthetic properties limit its fraction in the breathing gas to 35%. The use of a breathing mixture containing 30% xenon and 30% krypton is comparable in effectiveness for CT to a 40% xenon fraction, while avoiding the unwanted effects of a high fraction xenon gas.

 

Precautions

Krypton is considered to be a non-toxic asphyxiant. Krypton has a narcotic potency seven times greater than air, so breathing a gas containing 50% krypton and 50% air would cause narcosis similar to breathing air at four times atmospheric pressure. This would be comparable to scuba diving at a depth of 30 m (100 ft) and potentially could affect anyone breathing it. Nevertheless, that mixture would contain only 10% oxygen and hypoxia would be a greater concern.

 

Synopsis: https://en.wikipedia.org/wiki/Krypton

 

ARGON 1 of 2

Argon is a chemical element with symbol Ar and atomic number 18. It is in group 18 (noble gases) of the periodic table. Argon is the third most common gas in the Earth’s atmosphere, at 0.93% (9,300 ppm), making it approximately 23.8 times as abundant as next most common atmospheric gas, carbon dioxide (390 ppm), and more than 500 times as abundant as the next most common noble gas, neon (18 ppm). Nearly all of this argon is radiogenic argon-40 derived from the decay of potassium-40 in the Earth’s crust. In the universe, argon-36 is by far the most common argon isotope, being the preferred argon isotope produced by stellar nucleosynthesis in supernovas.

 

The name “argon” is derived from the Greek word αργον meaning “lazy” or “the inactive one”, a reference to the fact that the element undergoes almost no chemical reactions. The complete octet (eight electrons) in the outer atomic shell makes argon stable and resistant to bonding with other elements. Its triple point temperature of 83.8058 K is a defining fixed point in the International Temperature Scale of 1990.

 

Argon is produced industrially by the fractional distillation of liquid air. Argon is mostly used as an inert shielding gas in welding and other high-temperature industrial processes where ordinarily non-reactive substances become reactive; for example, an argon atmosphere is used in graphite electric furnaces to prevent the graphite from burning. Argon gas also has uses in incandescent and fluorescent lighting, and other types of gas discharge tubes. Argon makes a distinctive blue-green gas laser.

 

Characteristics

Argon has approximately the same solubility in water as oxygen, and is 2.5 times more soluble in water than nitrogen. Argon is colourless, odourless, and non-toxic as a solid, liquid, and gas. Argon is chemically inert under most conditions and forms no confirmed stable compounds at room temperature.

 

Although argon is a noble gas, it has been found to have the capability of forming some compounds. For example, the creation of argon fluorohydride (HArF), a marginally stable compound of argon with fluorine and hydrogen, was 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. Argon-containing ions and excited state complexes, such as ArH+ and ArF, respectively, are known to exist. Theoretical calculations have predicted several argon compounds that should be stable, but for which no synthesis routes are currently known.

 

History

Argon (αργος, Greek meaning “inactive”, in reference to its chemical inactivity) was suspected to be present in air by Henry Cavendish in 1785 but was not isolated until 1894 by Lord Rayleigh and Sir William Ramsay in Scotland in an experiment in which they removed all of the oxygen, carbon dioxide, water and nitrogen from a sample of clean air. They had determined that nitrogen produced from chemical compounds was one-half percent lighter than nitrogen from the atmosphere. The difference seemed insignificant, but it was important enough to attract their attention for many months. They concluded that there was another gas in the air mixed in with the nitrogen. Argon was also encountered in 1882 through independent research of H. F. Newall and W.N. Hartley. Each observed new lines in the colour spectrum of air but were unable to identify the element responsible for the lines. Argon became the first member of the noble gases to be discovered. The symbol for argon is now Ar, but up until 1957 it was A.

 

Occurrence

Argon constitutes 0.934% by volume and 1.28% 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.

 

Isotopes

The main isotopes of argon found on Earth are 40Ar (99.6%), 36Ar (0.34%), and 38Ar (0.06%). Naturally occurring 40K with a half-life of 1.25×109 years, decays to stable 40Ar (11.2%) by electron capture or positron emission, and also to stable 40Ca (88.8%) via beta decay. These properties and ratios are used to determine the age of rocks by the method of K-Ar dating.

 

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 neutron spallation of 40Ca as a result of subsurface nuclear explosions. It has a half-life of 35 days.

 

Argon is notable in that its isotopic composition varies greatly between different locations in the solar system. Where the major source of argon is the decay of 40K in rocks, 40Ar will be the dominant isotope, as it is on Earth. Argon produced directly by stellar nucleosynthesis, in contrast, is dominated by the alpha process nuclide, 36Ar. Correspondingly, solar argon contains 84.6% 36Ar based on solar wind measurements.

The predominance of radiogenic 40Ar is responsible for the fact that the standard atomic weight of terrestrial argon is greater than that of the next element, potassium. This was puzzling at the time when argon was discovered, since Mendeleev had placed the elements in his periodic table in order of atomic weight, although the inertness of argon implies that it must be placed before the reactive alkali metal potassium. Henry Moseley later solved this problem by showing that the periodic table is actually arranged in order of atomic number.

 

The much greater atmospheric abundance of argon relative to the other noble gases is also due to the presence of radiogenic 40Ar. Primordial 36Ar has an abundance of only 31.5 ppmv (= 9340 ppmv x 0.337%), comparable to that of neon (18.18 ppmv).

 

The Martian atmosphere contains 1.6% of 40Ar and 5 ppm of 36Ar. The Mariner space probe 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 40Ar on Titan, the largest moon of Saturn.

 

Compounds

Argon’s complete octet of electrons indicates full s and p sub shells 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 August 2000, the first argon compound was formed by researchers at the University of Helsinki. By shining ultraviolet light onto frozen argon containing a small amount of hydrogen fluoride with caesium iodide, argon fluorohydride (HArF) was formed. It is stable up to 40 kelvin (−233 °C). The metastable ArCF2+ 2 dication, which is valence isoelectronic with carbonyl fluoride, was observed in 2010.

 

Production

Industrial

Argon is produced industrially by the fractional distillation of liquid air in a cryogenic air separation unit; a process that separates liquid nitrogen, which boils at 77.3 K, from argon, which boils at 87.3 K, and liquid oxygen, which boils at 90.2 K. About 700,000 tonnes of argon are produced worldwide every year.

 

In radioactive decays

40Ar, the most abundant isotope of argon, is produced by the decay of 40K with a half-life of 1.25×109 years by electron capture or positron emission. Because of this, it is used in potassium-argon dating to determine the age of rocks.

 

Synopsis from: https://en.wikipedia.org/wiki/Argon

Applications

There are several different reasons argon is used in particular applications:

An inert gas is needed. In particular, argon is the cheapest alternative when nitrogen is not sufficiently inert.

Low thermal conductivity is required.

The electronic properties (ionization and/or the emission spectrum) are necessary.

Other noble gases would probably work as well in most of these applications, but argon is by far the cheapest. Argon is inexpensive since it is a by product of the production of liquid oxygen and liquid nitrogen from a cryogenic air separation unit, both of which are used on a large industrial scale. The other noble gases (except helium) are produced this way as well, but argon is the most plentiful by far, since it has a much higher concentration in the atmosphere. The bulk of argon applications arise simply because it is inert and relatively cheap.

Industrial processes

Argon is used in some high-temperature industrial processes, where ordinarily non-reactive substances become reactive. For example, an argon atmosphere is used in graphite electric furnaces to prevent the graphite from burning.

For some of these processes, the presence of nitrogen or oxygen gases might cause defects within the material. Argon is used in various types of arc welding such as gas metal arc welding and gas tungsten arc welding, as well as in the processing of titanium and other reactive elements. An argon atmosphere is also used for growing crystals of silicon and germanium.

Argon is an asphyxiant in the poultry industry, either for mass culling following disease outbreaks, or as a means of slaughter more humane than the electric bath. Argon’s relatively high density causes it to remain close to the ground during gassing. Its non-reactive nature makes it suitable in a food product, and since it replaces oxygen within the dead bird, argon also enhances shelf life.

Argon is sometimes used for extinguishing fires where damage to equipment is to be avoided.

Scientific research

Argon is used, primarily in liquid form, as the target for direct dark matter searches. The interaction of a hypothetical WIMP particle with the argon nucleus produces scintillation light that is then detected by photomultiplier tubes. Two-phase detectors also use argon gas to detect the ionized electrons produced during the WIMP-nucleus scattering. As with most other liquefied noble gases, argon has a high scintillation light yield (~ 51 photons / keV), is transparent to its own scintillation light, and is relatively easy to purify.

Compared to xenon, argon is cheaper and has a distinct scintillation time profile which allows the separation of electronic recoils from nuclear recoils. On the other hand, its intrinsic gamma-ray background is larger due to 39Ar contamination, unless one uses underground argon sources with a low level of radioactivity. Dark matter detectors currently operating with liquid argon include WArP, ArDM, micro Clean and DEAP-I.

Preservative

Argon is used to displace oxygen- and moisture-containing air in packaging material to extend the shelf-lives of the contents (argon has the European food additive code of E938). Aerial oxidation, hydrolysis, and other chemical reactions which degrade the products are retarded or prevented entirely.

Bottles of high-purity chemicals and certain pharmaceutical products are available in sealed bottles or ampoules packed in argon. In wine making, argon is used to top-off barrels to avoid the aerial oxidation of ethanol to acetic acid during the ageing process.

Argon is also available in aerosol-type cans, which may be used to preserve compounds such as varnish, polyurethane, paint, etc. for storage after opening.

Since 2002, the American National Archives stores important national documents such as the Declaration of Independence and the Constitution within argon-filled cases to retard their degradation. Using argon reduces gas leakage, compared with the helium used in the preceding five decades.

Laboratory equipment

Argon may be used as the inert gas within Schlenk lines and glove boxes The use of argon over comparatively less expensive nitrogen is preferred where nitrogen may react with the experimental reagents or apparatus.

Argon may be used as the carrier gas in gas chromatography and in electro spray ionization mass spectrometry; it is the gas of choice for the plasma used in ICP spectroscopy. Argon is preferred for the sputter coating of specimens for scanning electron microscopy. Argon gas is also commonly used for sputter deposition of thin films as in microelectronics and for wafer cleaning in micro fabrication

Medical use

Cryosurgery procedures such as cryoablation use liquefied argon to destroy cancer cells. In surgery it is used in a procedure called “argon enhanced coagulation” which is a form of argon plasma beam electro surgery The procedure carries a risk of producing gas embolism in the patient and has resulted in the death of one person via this type of accident. Blue argon lasers are used in surgery to weld arteries, destroy tumours, and to correct eye defects. It has also been used experimentally to replace nitrogen in the breathing or decompression mix, to speed the elimination of dissolved nitrogen from the blood.

Lighting

Incandescent lights are filled with argon, to preserve the filaments at high temperature from oxidation. It is used for the specific way it ionizes and emits light, such as in plasma globes and calorimetry in experimental particle physics. Gas-discharge lamps filled with argon provide blue light. Argon is also used for the creation of blue and green laser light.

Safety

Although argon is non-toxic, it is 38% denser than air and is therefore considered a dangerous asphyxiant in closed areas. It is also difficult to detect because it is colourless, odourless, and tasteless. A 1994 incident in which a man was asphyxiated after entering an argon filled section of oil pipe under construction in Alaska highlights the dangers of argon tank leakage in confined spaces, and emphasizes the need for proper use, storage and handling.

Synopsis from: https://en.wikipedia.org/wiki/Argon

Neon is a chemical element with symbol Ne and atomic number 10. It is in group 18 (noble gases) of the periodic table. Neon is a colourless, odourless monatomic gas under standard conditions, with about two-thirds the density of air. It was discovered (along with krypton and xenon) in 1898 as one of the three residual rare inert elements remaining in dry air, after nitrogen, oxygen, argon and carbon dioxide are removed. Neon was the second of these three rare gases to be discovered, and was immediately recognized as a new element from its bright red emission spectrum. Neon’s name is derived from Greek words meaning “new one.” Neon is chemically inert and forms no uncharged chemical compounds.

During cosmic nucleogenesis of the elements, large amounts of neon are built up from the alpha-capture fusion process in stars. Although neon is a very common element in the universe and solar system (it is fifth in cosmic abundance after hydrogen, helium, oxygen and carbon), it is very rare on Earth. It composes about 18.2 ppm of air by volume (this is about the same as the molecular or mole fraction), and a smaller fraction in the crust. The reason for neon’s relative scarcity on Earth and the inner (terrestrial) planets, is that neon forms no compounds to fix it to solids, and is highly volatile, therefore escaping from the planetesimals under the warmth of the newly-ignited Sun in the early Solar System. Even the atmosphere of Jupiter is somewhat depleted of neon, presumably for this reason.

Neon gives a distinct reddish- orange glow when used in either low-voltage neon glow lamps or in high-voltage discharge tubes or neon advertising signs. The red emission line from neon is also responsible for the well known red light of helium-neon lasers. Neon is used in a few plasma tube and refrigerant applications but has few other commercial uses. It is commercially extracted by the fractional distillation of liquid air. It is considerably more expensive than helium, since air is its only source.

History

Neon (Greek ???? (neon) meaning “new one”) was discovered in 1898 by the British chemists Sir William Ramsay (1852–1916) and Morris W. Travers (1872–1961) in London. Neon was discovered when Ramsay chilled a sample of air until it became a liquid, then warmed the liquid and captured the gases as they boiled off. The gases nitrogen, oxygen, and argon had been identified, but the remaining gasses were isolated in roughly their order of abundance, in a six-week period beginning at the end of May 1898. First to be identified was krypton. The next, after krypton had been removed, was a gas which gave a brilliant red light under spectroscopic discharge. This gas, identified in June, was named neon, the Greek analogue of “novum,” (new), the name Ramsay’s son suggested. The characteristic brilliant red-orange colour that is emitted by gaseous neon when excited electrically was noted immediately; Travers later wrote, “the blaze of crimson light from the tube told its own story and was a sight to dwell upon and never forget.” Finally, the same team discovered xenon by the same process, in July.

Neon’s scarcity precluded its prompt application for lighting along the lines of Moore tubes, which used nitrogen and which were commercialized in the early 1900s. After 1902, Georges Claude’s company, Air Liquide, was producing industrial quantities of neon as a by product of his air liquefaction business. In December 1910 Claude demonstrated modern neon lighting based on a sealed tube of neon. Claude tried briefly to get neon tubes to be used for indoor lighting, due to their intensity, but failed, as home owners rejected neon light sources due to their colour. Finally in 1912, Claude’s associate began selling neon discharge tubes as advertising signs, where they were instantly more successful as eye catchers. They were introduced to the U.S. in 1923, when two large neon signs were bought by a Los Angeles Packard car dealership. The glow and arresting red colour made neon advertising completely different from the competition.

Neon played a role in the basic understanding of the nature of atoms in 1913, when J. J. Thomson, as part of his exploration into the composition of canal rays, channelled streams of neon ions through a magnetic and an electric field and measured their deflection by placing a photographic plate in their path.

Thomson observed two separate patches of light on the photographic plate, which suggested two different parabolas of deflection. Thomson eventually concluded that some of the atoms in the neon gas were of higher mass than the rest. Though not understood at the time by Thomson, this was the first discovery of isotopes of stable atoms. It was made by using a crude version of an instrument we now term as a mass spectrometer.

Isotopes

The first evidence for isotopes of a stable element. In the bottom right corner of J. J. Thomson’s photographic plate are the separate impact marks for the two isotopes neon-20 and neon-22.

Neon is the second lightest inert gas. Neon has three stable isotopes: 20Ne (90.48%), 21Ne (0.27%) and 22Ne (9.25%). 21Ne and 22Ne are partly primordial and partly nucleogenic (i.e., made by nuclear reactions of other nuclides with neutrons or other particles in the environment) and their variations in natural abundance are well understood. In contrast, 20Ne (the chief primordial isotope made in stellar nucleosynthesis) is not known to be nucleogenic or radiogenic (save for cluster decay production, which is thought to produce only a small amount). The causes of the variation of 20Ne in the Earth have thus been hotly debated.

The principal nuclear reactions which generate nucleogenic neon isotopes start from 24Mg and 25Mg, which produce 21Ne and 22Ne, respectively, after neutron capture and immediate emission of an alpha particle. The neutrons that produce the reactions are mostly produced by secondary spallation reactions from alpha particles, in turn derived from uranium-series decay chains. The net result yields a trend towards lower 20Ne/22Ne and higher 21Ne/22Ne ratios observed in uranium-rich rocks such as granites. Neon-21 may also be produced in a nucleogenic reaction, when 20Ne absorbs a neutron from various natural terrestrial neutron sources.

In addition, isotopic analysis of exposed terrestrial rocks has demonstrated the cosmogenic (cosmic ray) production of 21Ne. This isotope is generated by spallation reactions on magnesium, sodium, silicon, and aluminium. By analysing all three isotopes, the cosmogenic component can be resolved from magmatic neon and nucleogenic neon. This suggests that neon will be a useful tool in determining cosmic exposure ages of surface rocks and meteorites.

Similar to xenon, neon content observed in samples of volcanic gases is enriched in 20Ne, as well as nucleogenic 21Ne, relative to 22Ne content. The neon isotopic content of these mantle-derived samples represents a non-atmospheric source of neon. The 20Ne-enriched components are attributed to exotic primordial rare gas components in the Earth, possibly representing solar neon. Elevated 20Ne abundances are found in diamonds, further suggesting a solar neon reservoir in the Earth.

Characteristics

Neon is the second-lightest noble gas, after helium. It glows reddish-orange in a vacuum discharge tube. Also, neon has the narrowest liquid range of any element: from 24.55 K to 27.05 K (-248.45 °C to -245.95 °C, or -415.21 °F to -410.71 °F). It has over 40 times the refrigerating capacity of liquid helium and three times that of liquid hydrogen (on a per unit volume basis). In most applications it is a less expensive refrigerant than helium.

Neon plasma has the most intense light discharge at normal voltages and currents of all the noble gases. The average colour of this light to the human eye is red-orange due to many lines in this range; it also contains a strong green line which is hidden, unless the visual components are dispersed by a spectroscope.

Two quite different kinds of neon lighting are in common use. Neon glow lamps are generally tiny, with most operating at about 100–250 volts. They have been widely used as power-on indicators and in circuit-testing equipment, but light-emitting diodes (LEDs) now dominate in such applications. These simple neon devices were the forerunners of plasma displays and plasma television screens. Neon signs typically operate at much higher voltages (2–15 kilovolts), and the luminous tubes are commonly meters long. The glass tubing is often formed into shapes and letters for signage as well as architectural and artistic applications.

Occurrence

Stable isotopes of neon are produced in stars. 20Ne is created in fusing helium and oxygen in the alpha process, which requires temperatures above 100 mega kelvins and masses greater than 3 solar masses.

Neon is abundant on a universal scale; it is the fifth most abundant chemical element in the universe by mass, after hydrogen, helium, oxygen, and carbon. Its relative rarity on Earth, like that of helium, is due to its relative lightness, high vapour pressure at very low temperatures, and chemical inertness, all properties which tend to keep it from being trapped in the condensing gas and dust clouds which resulted in the formation of smaller and warmer solid planets like Earth.

Neon is monatomic, making it lighter than the molecules of diatomic nitrogen and oxygen which form the bulk of Earth’s atmosphere; a balloon filled with neon will rise in air, albeit more slowly than a helium balloon.

Mass abundance in the universe is about 1 part in 750 and in the Sun and presumably in the proto-solar system nebula, about 1 part in 600. The Galileo spacecraft atmospheric entry probe found that even in the upper atmosphere of Jupiter, the abundance of neon is reduced (depleted) by about a factor of 10, to a level of 1 part in 6,000 by mass. This may indicate that even the ice- planetesimals which brought neon into Jupiter from the outer solar system, formed in a region which was too warm for them to have kept their neon (abundances of heavier inert gases on Jupiter are several times that found in the Sun).

Neon is rare on Earth, found in the Earth’s atmosphere at 1 part in 55,000, or 18.2 ppm by volume (this is about the same as the molecule or mole fraction), or 1 part in 79,000 of air by mass. It comprises a smaller fraction in the crust. It is industrially produced by cryogenic fractional distillation of liquefied air.

Applications

Neon is often used in signs and produces an unmistakable bright reddish-orange light. Although still referred to as “neon”, all other colours are generated with the other noble gases or by many colours of fluorescent lighting.

Neon is used in vacuum tubes, high-voltage indicators, lightning arrestors, wave meter tubes, television tubes, and helium-neon lasers. Liquefied neon is commercially used as a cryogenic refrigerant in applications not requiring the lower temperature range attainable with more extreme liquid helium refrigeration.

Both neon gas and liquid neon are relatively expensive – for small quantities, the price of liquid neon can be more than 55 times that of liquid helium. The driver for neon’s expense is the rarity of neon, which unlike helium, can only be obtained from air.

The triple point temperature of neon (24.5561 K) is a defining fixed point in the International Temperature Scale of 1990.

Compounds

Neon is the first p-block noble gas. Neon is generally considered to be inert. No true neutral compounds of neon are known. However, the ions Ne+, (NeAr)+, (NeH)+, and (HeNe+) have been observed from optical and mass spectrometric studies, and there are some unverified reports of an unstable hydrate.

Synopsis from: https://en.wikipedia.org/wiki/Argon

 

The noble gases are the elements in group 18 (also sometimes Group 0 IUPAC Style, or Group 8) of the periodic table. The group is also called the helium family or neon family. Chemically, the noble gases are very stable due to having the maximum number of valence electrons their outer shell can hold. Noble gases rarely react with other elements since they are already stable. Under normal conditions, they occur as odourless, colourless, monatomic gases. Each of them has its melting and boiling point close together, so that only a small temperature range exists for each noble gas in which it is a liquid. Noble gases have numerous important applications in lighting, welding and space technology.

The seven noble gases are: helium, neon, argon, krypton, xenon, radon, and ununoctium.

Etymology

“Noble gas” is the translation of the German Edelgas, which was in use as early as 1898. This refers to the extremely low level of reactivity under normal conditions. The noble gases have also been referred to as inert gases, but these terms are not strictly accurate because several of them do take part in chemical reactions. Another old term is rare gases, although argon forms a fairly considerable part (0.93% by volume, 1.29% by mass) of the Earth’s atmosphere.

History

The existence of noble gases was not known until after the advent of the periodic table. In the late nineteenth century, Lord Rayleigh discovered that some samples of nitrogen from the air were of a different density than nitrogen resulting from chemical reactions. Along with scientist William Ramsay, Lord Rayleigh theorized that the nitrogen extracted from air was associated with another gas, argon. With this discovery, they realized that a whole class of gases was missing from the periodic table. Eventually, all the known noble gases except for helium were discovered in the air, with argon being much more common than the others, and the table was completed. Helium was detected spectrographically in the Sun in 1868. The isolation of helium on Earth had to wait until 1895. Under standard conditions, the noble gases all occur as monatomic gases.

Chemical make-up

Noble gases have full valence electron shells. Valence electrons are the outermost electrons of an atom and are normally the only electrons which can participate in chemical bonding. According to atomic theory derived from quantum mechanics and experimental trends, atoms with full valence electron shells are extraordinarily stable and therefore do not form chemical bonds.

All of them exhibit an extremely low chemical reactivity and very few noble gas compounds have been prepared. No conventional compounds of helium or neon have yet been prepared, while xenon and krypton are known to show some reactivity in the laboratory. Recently argon compounds have also been successfully characterised. The noble gases’ lack of reactivity can be explained in terms of them having a “complete valence shell”. They have little tendency to gain or lose electrons. The noble gases have high ionization energies and negligible electro negativities The noble gases have very weak inter-atomic forces of attraction, and consequently very low melting points and boiling points. This is why they are all monatomic gases under normal conditions, even those with larger atomic masses than many normally solid elements.

Applications

One of the most commonly encountered uses of the noble gases in everyday life is in lighting. Argon is often used as a suitable safe and inert atmosphere for the inside of filament light bulbs, and is also used as an inert atmosphere in the synthesis of air and moisture sensitive compounds (as an alternative for nitrogen). Some of the noble gases glow distinctive colours when used inside lighting tubes ( neon lights). Helium, due to its non reactivity (compared with flammable hydrogen) and lightness, is often used in blimps and balloons. Helium and argon are commonly used to shield a welding arc, and the surrounding base metal from the atmosphere during welding. Krypton is also used in lasers, which are used by doctors for eye surgery. Xenon is used in xenon arc lamps, and it has anaesthetic properties.

Noble gas notation

The noble gases can be used in conjunction with the electron configuration notation to make what is called the Noble Gas Notation. For example: while the electron notation of the element carbon is 1s²2s² 2p², the Noble Gas notation would be [He] 2s²2p².

This notation makes the identification of elements faster, and is shorter and easier than writing out the full notation of orbitals.

Synopsis from: https://en.wikipedia.org/wiki/Noble_gas

Helium (He) is a colourless, odourless, tasteless, non-toxic, inert monatomic chemical element that heads the noble gas series in the periodic table and whose atomic number is 2. Its boiling and melting points are the lowest among the elements and it exists only as a gas except in extreme conditions. Extreme conditions are also needed to create the small handful of helium compounds, which are all unstable at standard temperature and pressure. In its most common form, helium-4, it has two neutrons in its nucleus, while a second, rarer, stable isotope called helium-3 contains just one neutron. The behaviour of liquid helium-4’s two fluid phases, helium I and helium II, is important to researchers studying quantum mechanics (in particular the phenomenon of super fluidity) and to those looking at the effects that temperatures near absolute zero have on matter (such as superconductivity).

In 1868 the French astronomer Pierre Janssen first detected helium as an unknown yellow spectral line signature in light from a solar eclipse. Since then large reserves of helium have been found in the natural gas fields of the United States, which is by far the largest supplier of the gas. It is used in cryogenics, in deep-sea breathing systems, to cool superconducting magnets, in helium dating, for inflating balloons, for providing lift in airships and as a protective gas for many industrial uses (such as arc welding and growing silicon wafers). A much less serious use is to temporarily change the timbre and quality of one’s voice by inhaling a small volume of the gas (see precautions section below).

Helium is the second most abundant and second lightest element in the known universe, and is one of the elements believed to have been created in the Big Bang. In the modern universe almost all new helium is created as a result of the nuclear fusion of hydrogen in stars. On Earth helium is rare, and almost all of that which exists was created by the radioactive decay of much heavier elements ( alpha particles are helium nuclei). After its creation, part of it was trapped with natural gas in concentrations up to 7% by volume, from which it is extracted commercially by fractional distillation. Large reserves of helium have been found in the natural gas fields of the United States (the largest supplier) but helium is known in gas reserves of a few other countries.

 

Notable characteristics

 

Gas and plasma phases

 

Helium is the least reactive member of the noble gas elements, and thus also the least reactive of all elements; it is inert and monatomic in virtually all conditions. Due to helium’s relatively low molar (molecular) mass, in the gas phase it has a thermal conductivity, specific heat, and sound conduction velocity that are all greater than any gas, except hydrogen. For similar reasons, and also due to the small size of its molecules, helium’s diffusion rate through solids is three times that of air and around 65% that of hydrogen.

 

Helium is less water soluble than any other gas known, and helium’s index of refraction is closer to unity than that of any other gas. Helium has a negative Joule-Thomson coefficient at normal ambient temperatures, meaning it heats up when allowed to freely expand. Only below its Joule-Thomson inversion temperature (of about 40 K at 1 atmosphere) does it cool upon free expansion. Once pre cooled below this temperature, helium can be liquefied through expansion cooling.

 

Throughout the universe, helium is found mostly in a plasma state whose properties are quite different from atomic helium. In a plasma, helium’s electrons and protons are not bound together, resulting in very high electrical conductivity, even when the gas is only partially ionized. The charged particles are highly influenced by magnetic and electric fields. For example, in the solar wind together with ionized hydrogen, they interact with the Earth’s magnetosphere giving rise to Birkeland currents and the aurora.

 

Solid and liquid phases

 

Helium solidifies only under great pressure. The resulting colourless, almost invisible solid is highly compressible; applying pressure in a laboratory can decrease its volume by more than 30%. With a bulk modulus on the order of 5×107 Pa it is 50 times more compressible than water. Unlike any other element, helium will fail to solidify and remain a liquid down to absolute zero at normal pressures. This is a direct effect of quantum mechanics: specifically, the zero point energy of the system is too high to allow freezing. Solid helium requires a temperature of 1–1.5 K (about −272 °C or −457 °F) and about 25 bar (2.5 MPa) of pressure. It is often hard to distinguish solid from liquid helium since the refractive index of the two phases are nearly the same. The solid has a sharp melting point and has a crystalline structure.

 

Solid helium has a density of 0.214 ±0.006 g/ml (1.15 K, 66 atm) with a mean isothermal compressibility of the solid at 1.15 K between the solidus and 66 atm of 0.0031 ±0.0008/atm. Also, no difference in density was noted between 1.8 K and 1.5 K. This data projects that T=0 solid helium under 25 bar of pressure (the minimum required to freeze helium) has a density of 0.187 ±0.009 g/ml.

 

Helium I state

 

Below its boiling point of 4.22 kelvin and above the lambda point of 2.1768 kelvin, the isotope helium-4 exists in a normal colourless liquid state, called helium I. Like other cryogenic liquids, helium I boils when it is heated. It also contracts when its temperature is lowered until it reaches the lambda point, when it stops boiling and suddenly expands. The rate of expansion decreases below the lambda point until about 1 K is reached; at which point expansion completely stops and helium I starts to contract again.

 

Helium I has a gas-like index of refraction of 1.026 which makes its surface so hard to see that floats of Styrofoam are often used to show where the surface is. This colourless liquid has a very low viscosity and a density one-eighth that of water, which is only one-fourth the value expected from classical physics. Quantum mechanics is needed to explain this property and thus both types of liquid helium are called quantum fluids, meaning they display atomic properties on a macroscopic scale. This is probably due to its boiling point being so close to absolute zero, which prevents random molecular motion (heat) from masking the atomic properties.

 

Helium II state

 

Liquid helium below its lambda point begins to exhibit very unusual characteristics, in a state called helium II. Boiling of helium II is not possible due to its high thermal conductivity; heat input instead causes evaporation of the liquid directly to gas. The isotope helium-3 also has a super fluid phase, but only at much lower temperatures; as a result, less is known about such properties in the isotope helium-3.

 

Helium II is a super fluid, a quantum-mechanical state of matter with strange properties. For example, when it flows through even capillaries of 10−7 to 10−8 m width it has no measurable viscosity. However, when measurements were done between two moving discs, a viscosity comparable to that of gaseous helium was observed. Current theory explains this using the two-fluid model for helium II. In this model, liquid helium below the lambda point is viewed as containing a proportion of helium atoms in a ground state, which are super fluid and flow with exactly zero viscosity, and a proportion of helium atoms in an excited state, which behave more like an ordinary fluid.

 

Helium II also exhibits a creeping effect. When a surface extends past the level of helium II, the helium II moves along the surface, seemingly against the force of gravity. Helium II will escape from a vessel that is not sealed by creeping along the sides until it reaches a warmer region where it evaporates. It moves in a 30 nm-thick film regardless of surface material. This film is called a Rollin film and is named after the man who first characterized this trait, Bernard V. Rollin. As a result of this creeping behaviour and helium II’s ability to leak rapidly through tiny openings, it is very difficult to confine liquid helium. Unless the container is carefully constructed, the helium II will creep along the surfaces and through valves until it reaches somewhere warmer, where it will evaporate. Waves propagating across a Rollin film are governed by the same equation as gravity waves in shallow water, but rather than gravity, the restoring force is the Van der Waals force. These waves are known as third sound.

 

In the fountain effect, a chamber is constructed which is connected to a reservoir of helium II by a sintered disc through which super fluid helium leaks easily but through which non superfluid helium cannot pass. If the interior of the container is heated, the super fluid helium changes to non superfluid helium. In order to maintain the equilibrium fraction of super fluid helium, super fluid helium leaks through and increases the pressure, causing liquid to fountain out of the container.

 

The thermal conductivity of helium II is greater than that of any other known substance, a million times that of helium I and several hundred times that of copper. This is because heat conduction occurs by an exceptional quantum-mechanical mechanism. Most materials that conduct heat well have a valence band of free electrons which serve to transfer the heat. Helium II has no such valence band but nevertheless conducts heat well. The flow of heat is governed by equations that are similar to the wave equation used to characterize sound propagation in air. So when heat is introduced, it will move at 20 meters per second at 1.8 K through helium II as waves in a phenomenon called second sound.

 

Synopsis of: https://en.wikipedia.org/wiki/Helium

 

Applications

Helium is used for many purposes that require some of its unique properties, such as its low boiling point, low density, low solubility, high thermal conductivity, or inertness. Helium is commercially available in either liquid or gaseous form. As a liquid, it can be supplied in small containers called dewars which hold up to 1,000 litres of helium, or in large ISO containers which have nominal capacities as large as 11,000 gallons (41,637 litres). In gaseous form, small quantities of helium are supplied in high pressure cylinders holding up to 300 standard cubic feet, while large quantities of high pressure gas are supplied in tube trailers which have capacities of up to 180,000 standard cubic feet.

 

Because it is lighter than air, airships and balloons are inflated with helium for lift. In airships, helium is preferred over hydrogen because it is not flammable and has 92.64% of the buoyancy (or lifting power) of the alternative hydrogen.

For its low solubility in water, the major part of human blood, mixtures of helium with oxygen and nitrogen ( trimix), with oxygen only ( heliox), with common air ( heliair), and with hydrogen and oxygen ( hydreliox), are used in deep-sea breathing systems to reduce the high-pressure risk of nitrogen narcosis.

 

At extremely low temperatures, liquid helium is used to cool certain metals to produce superconductivity, such as in superconducting magnets used in magnetic resonance imaging. Helium at low temperatures is also used in cryogenics.

For its inertness and high thermal conductivity, neutron transparency, and because it does not form radioactive isotopes under reactor conditions, helium is used as a coolant in some nuclear reactors, such as pebble-bed reactors.

Helium is used as a shielding gas in arc welding processes on materials that are contaminated easily by air. It is especially useful in overhead welding, because it is lighter than air and thus floats, whereas other shielding gases sink.

Because it is inert, helium is used as a protective gas in growing silicon and germanium crystals, in titanium and zirconium production, in gas chromatography, and as an atmosphere for protecting historical documents. This property also makes it useful in supersonic wind tunnels.

 

In rocketry, helium is used as an ullage medium to displace fuel and oxidisers in storage tanks and to condense hydrogen and oxygen to make rocket fuel. It is also used to purge fuel and oxidizer from ground support equipment prior to launch and to pre-cool liquid hydrogen in space vehicles. For example, the Saturn V booster used in the Apollo program needed about 13 million cubic feet (370,000 m³) of helium to launch.

 

The gain medium of the helium-neon laser is a mixture of helium and neon.

Because it diffuses through solids at a rate three times that of air, helium is used as a tracer gas to detect leaks in high-vacuum equipment and high-pressure containers, as well as in other applications with less stringent requirements such as heat exchangers, valves, gas panels, etc.

 

Because of its extremely low index of refraction, the use of helium reduces the distorting effects of temperature variations in the space between lenses in some telescopes.

 

The age of rocks and minerals that contain uranium and thorium, radioactive elements that emit helium nuclei called alpha particles, can be discovered by measuring the level of helium with a process known as helium dating.

 

The high thermal conductivity and sound velocity of helium is also desirable in thermo acoustic refrigeration. The inertness of helium adds to the environmental advantage of this technology over conventional refrigeration systems which may contribute to ozone depleting and global warming effects.

Because helium alone is less dense than atmospheric air, it will change the timbre (not pitch) of a person’s voice when inhaled. However, inhaling it from a typical commercial source, such as that used to fill balloons, can be dangerous due to the risk of asphyxiation from lack of oxygen, and the number of contaminants that may be present. These could include trace amounts of other gases, in addition to aerosolized lubricating oil.

 

History

 

Scientific discoveries

Evidence of helium was first detected on August 18, 1868 as a bright yellow line with a wavelength of 587.49 nanometres in the spectrum of the chromosphere of the Sun, by French astronomer Pierre Janssen during a total solar eclipse in Guntur, India. This line was initially assumed to be sodium. On October 20 of the same year, English astronomer Norman Lockyer observed a yellow line in the solar spectrum, which he named the D3 line, for it was near the known D1 and D2 lines of sodium, and concluded that it was caused by an element in the Sun unknown on Earth. He and English chemist Edward Frankland named the element with the Greek word for the Sun, ????? (helios).

 

On 26 March 1895 British chemist William Ramsay isolated helium on Earth by treating the mineral cleveite with mineral acids. Ramsay was looking for argon but, after separating nitrogen and oxygen from the gas liberated by sulphuric acid, noticed a bright-yellow line that matched the D3 line observed in the spectrum of the Sun. These samples were identified as helium by Lockyer and British physicist William Crookes. It was independently isolated from cleveite the same year by chemists Per Teodor Cleve and Abraham Langlet in Uppsala, Sweden, who collected enough of the gas to accurately determine its atomic weight. Helium was also isolated by the American geochemist William Francis Hillebrand prior to Ramsay’s discovery when he noticed unusual spectral lines while testing a sample of the mineral uraninite. Hillebrand, however, attributed the lines to nitrogen. His letter of congratulations to Ramsay offers an interesting case of discovery and near-discovery in science.

 

In 1907, Ernest Rutherford and Thomas Royds demonstrated that alpha particles are helium nuclei, by allowing them to penetrate the thin glass wall of a evacuated tube, then creating a discharge in the tube to study the spectra of the new gas inside. In 1908, helium was first liquefied by Dutch physicist Heike Kamerlingh Onnes by cooling the gas to less than one kelvin. He tried to solidify it by further reducing the temperature but failed, because helium does not have a triple point temperature where the solid, liquid, and gas phases are at equilibrium. It was first solidified in 1926 by his student Willem Hendrik Keesom by subjecting helium to 25 atmospheres of pressure.

 

In 1938, Russian physicist Pyotr Leonidovich Kapitsa discovered that helium-4 (a boson) has almost no viscosity at temperatures near absolute zero, a phenomenon now called super fluidity This phenomenon is related to Bose-Einstein condensation. In 1972, the same phenomenon was observed in helium-3, but at temperatures much closer to absolute zero, by American physicists Douglas D. Osheroff, David M. Lee, and Robert C. Richardson. The phenomenon in helium-3 is thought to be related to pairing of helium-3 fermions to make bosons, in analogy to Cooper pairs of electrons producing superconductivity.

 

Extraction and uses

After an oil drilling operation in 1903 in Dexter, Kansas, U.S. produced a gas geyser that would not burn, Kansas state geologist Erasmus Haworth collected samples of the escaping gas and took them back to the University of Kansas at Lawrence where, with the help of chemists Hamilton Cady and David McFarland, he discovered that the gas contained, by volume, 72% nitrogen, 15% methane—insufficient to make the gas combustible, 1% hydrogen, and 12% of an unidentifiable gas. With further analysis, Cady and McFarland discovered that 1.84% of the gas sample was helium. Far from being a rare element, helium was present in vast quantities under the American Great Plains, available for extraction from natural gas.

 

This put the United States in an excellent position to become the world’s leading supplier of helium. Following a suggestion by Sir Richard Threlfall, the United States Navy sponsored three small experimental helium production plants during World War I. The goal was to supply barrage balloons with the non-flammable lifting gas. A total of 200,000 cubic feet (5700 m³) of 92% helium was produced in the program even though only a few cubic feet (less than 100 litres) of the gas had previously been obtained. Some of this gas was used in the world’s first helium-filled airship, the U.S. Navy’s C-7, which flew its maiden voyage from Hampton Roads, Virginia to Bolling Field in Washington, D.C. on 1 December 1921.

 

Although the extraction process, using low-temperature gas liquefaction, was not developed in time to be significant during World War I, production continued. Helium was primarily used as a lifting gas in lighter-than-air craft. This use increased demand during World War II, as well as demands for shielded arc welding. Helium was also vital in the atomic bomb Manhattan Project.

 

The government of the United States set up the National Helium Reserve in 1925 at Amarillo, Texas with the goal of supplying military airships in time of war and commercial airships in peacetime. Due to a US military embargo against Germany that restricted helium supplies, the Hindenburg was forced to use hydrogen as the lift gas. Helium use following World War II was depressed but the reserve was expanded in the 1950s to ensure a supply of liquid helium as a coolant to create oxygen/hydrogen rocket fuel (among other uses) during the Space Race and Cold War. Helium use in the United States in 1965 was more than eight times the peak wartime consumption.

 

After the “Helium Acts Amendments of 1960” (Public Law 86–777), the U.S. Bureau of Mines arranged for five private plants to recover helium from natural gas. For this helium conservation program, the Bureau built a 425-mile (684 km) pipeline from Bushton, Kansas to connect those plants with the government’s partially depleted Cliffside gas field, near Amarillo, Texas. This helium-nitrogen mixture was injected and stored in the Cliffside gas field until needed, when it then was further purified.

 

By 1995, a billion cubic metres of the gas had been collected and the reserve was US$1.4 billion in debt, prompting the Congress of the United States in 1996 to phase out the reserve. The resulting “Helium Privatization Act of 1996” (Public Law 104–273) directed the United States Department of the Interior to start liquidating the reserve by 2005.

 

Helium produced before 1945 was about 98% pure (2% nitrogen), which was adequate for airships. In 1945 a small amount of 99.9% helium was produced for welding use. By 1949 commercial quantities of Grade A 99.995% helium were available.

 

For many years the United States produced over 90% of commercially usable helium in the world. Extraction plants created in Canada, Poland, Russia, and other nations produced the remaining helium. In the mid 1990s, A new plant in Arzew, Algeria producing 600mmcf came on stream, with enough production to cover all of Europe’s demand. Subsequently, in 2004–2006 two additional plants, one in Ras Laffen, Qatar and the other in Skikda, Algeria were built, but as of early 2007, Ras Laffen is functioning at 50%, and Skikda has yet to start up.

 

Algeria quickly became the second leading producer of helium. Through this time, both helium consumption and the costs of producing helium increased and during 2007 the major suppliers, Air Liquide, Airgas and Praxair all raised prices from 10 to 30%.

 

Synopsis of: https://en.wikipedia.org/wiki/Helium