Noble gases and their properties. Noble gases and their properties Noble gases what causes the inertness of these gases

Opening:

In 1893, attention was drawn to the discrepancy between the densities of nitrogen from the air and nitrogen obtained from the decomposition of nitrogen compounds: a liter of nitrogen from the air weighed 1.257 g, and that obtained chemically weighed 1.251 g. A very accurate study of the composition of the air carried out to clarify this mysterious circumstance showed that after all the oxygen and nitrogen were removed, there was a small residue (about 1%) that did not react chemically with anything.

The discovery of a new element, called argon (Greek for inactive), thus represented the “triumph of the third decimal place.” The molecular weight of argon turned out to be 39.9 g/mol.

The next inert gas to be discovered, helium (“solar”), was discovered on the Sun earlier than on Earth. This turned out to be possible thanks to the spectral analysis method developed in the 50s of the last century.

A few years after the discovery of argon and helium (in 1898), three more noble gases were isolated from the air: neon (“new”), krypton (“hidden”) and xenon (“alien”). How difficult it was to detect them can be seen from the fact that 1 m 3 of air, along with 9.3 liters of argon, contains only 18 ml of neon, 5 ml of helium, 1 ml of krypton and 0.09 ml of xenon.

The last inert gas, radon, was discovered in 1900 while studying certain minerals. Its content in the atmosphere is only 6-10 -18% by volume (which corresponds to 1-2 atoms per cubic centimeter). It has been estimated that the entire earth's atmosphere contains only 374 liters of radon.

Physical properties:

All noble gases are colorless and consist of monatomic molecules. The separation of inert gases is based on the difference in their physical properties.

Inert gases are colorless and odorless. They are present in small quantities in the air. Inert gases are not poisonous. However, an atmosphere with an increased concentration of inert gases and a corresponding decrease in oxygen concentration can have a suffocating effect on a person, including loss of consciousness and death. There are known cases of death due to argon leaks.

Melting point, °C
Boiling point, °C

The amount of heat required to transfer a substance from a solid to a liquid state is called the heat of fusion, and to transfer from a liquid to a vapor state is called the heat of evaporation. Both quantities are usually referred to as transitions occurring under normal pressure. For inert gases they have the following values ​​(kcal/g-atom):

Heat of Melting
Heat of vaporization

Below are compared critical temperatures inert gases and those pressures that are necessary and sufficient for their transfer at these temperatures from a gaseous state to a liquid state, - critical pressures:

Critical temperature, °C
Critical pressure, atm

This is interesting :

The question of the atomicity of the argon molecule was resolved using kinetic theory. According to it, the amount of heat that needs to be expended to heat a gram-molecule of a gas by one degree depends on the number of atoms in its molecule. At constant volume, a gram-molecule of a monatomic gas requires 3 feces, diatomic - 5 cal. For argon the experiment gave 3 feces, which indicated the monoatomic nature of its molecule. The same applies to other inert gases.

Helium was the last gas to be converted into a liquid and solid state. In relation to it, there were special difficulties due to the fact that as a result of expansion at ordinary temperatures, helium does not cool, but heats up. Only below -250 °C does it begin to behave “normally”. It follows that the usual liquefaction process could be applied to helium only after it had been very strongly cooled beforehand. On the other hand, the critical temperature of helium is extremely low. Due to these circumstances, favorable results when working with helium were obtained only after mastering the technique of operating with liquid hydrogen, using the evaporation of which only it was possible to cool helium to the required temperatures. It was possible to obtain liquid helium for the first time in 1908, solid helium-V1926

Chemical properties:

Inert gases are characterized by a complete (He, Ne, Ar) or almost complete (Kr, Xe, Rn) lack of chemical activity. In the periodic table they form a special group (VIII). Soon after the discovery of inert gases, the new group they formed in the periodic table was called zero, in order to emphasize the zero valency of these elements, i.e., their lack of chemical activity. This name is often used at the present time, however, in essence of the periodic law, it is more correct to consider the group of inert gases as the eighth group, since the corresponding periods do not begin with these elements, but end.

The absence of complete chemical inertness in heavy inert gases was discovered only in 1962. It turned out that they are capable of combining with the most active metalloid - fluorine (and only with it). Xenon (and radon) react quite easily, krypton much more difficult. XeF 2 , XeF 4 , XeF 6 and low-stable KrF 2 were obtained. All of them are colorless volatile crystalline substances.

Xenon difluoride(XeF 2) - is slowly formed under the influence of daylight on a mixture of Xe and F 2 at zero conditions. It has a characteristic nauseating odor. The formation of a molecule requires excitation of the xenon atom from 5s 2 5p 6 to the nearest divalent state 5s 2 5p 5 s 1 - 803 kJ/mol, to 5s 2 5p 5 6p 1 -924 kJ/mol, 25s 2 5p 1 6d 1 - 953 kJ/ mole.

Xe+F 2 →XeF 2

0.15 mol/l dissolves in water. The solution is a very strong oxidizing agent. The solution decomposes according to the following scheme:

XeF 2 +H 2 O →HF+Xe+O 2 (the process occurs faster in an alkaline environment, slower in an acidic environment).

Xenontetrafluoride- formed from simple substances, the reaction is highly exothermic, and is the most stable of all fluorides.

XeF 4 +2Hg=2HgF 2 +Xe

XeF 4 +Pt=PtF 4 +Xe

Qualitative reaction to xenon tetrafluoride :

XeF 4 +4KI=4KF+2I 2 ↓+Xe

Xenon tetrafluoride decomposes according to the following schemes:

3Xe 4+ →Xe 6+ +2Xe 0 (in acidic medium).

Xe 4+ →Xe 0 +Xe 8+ (in an alkaline medium).

Xenon hexafluoride is colorless, known in 3 crystalline modifications. At 49 ℃, turning into a yellow liquid, when hardening it becomes discolored again. The vapors are pale yellow in color. Explosively decomposes. Under the influence of moist air hydrolyze:

XeF 6 +H 2 O→2HF+OXeF 4

OXeF 4 is a colorless liquid, less reactive than XeF 6. Forms crystalline hydrates with alkali metal fluorides, for example: KF∙OXeF 4

Further hydrolysis can produce xenon trioxide:

XeF 6 +3H 2 O→XeO 3 +6HF

XeO 3 is a colorless explosive substance that diffuses in air. It disintegrates explosively, but when gently heated at 40 degrees Celsius, the reaction occurs:

2XeO 3 →2Xe+3O 2

There is an acid that formally corresponds to this oxide - H 2 XeO 4. There are salts corresponding to this acid: MHXeO 4 or MH 5 XeO 6, an acid (M - from sodium to cesium) corresponding to the last salt was obtained:

3XeF 4 +6Ca(OH) 2 →6CaF 2 ↓+Xe+2H 2 XeO 6

In a strongly alkaline environment, Xe 6+ dismutates:

4Xe 6+ →Xe 0 +3Xe 8+

Krypton difluoride- volatile, colorless crystals , a chemically active substance. At elevated temperatures it decomposes into fluorine krypton . It was first obtained by the action of an electric discharge on a mixture of substances, at -188℃:

F 2 +Kr→KrF 2

Decomposes with water according to the following scheme:

2KrF 2 +2H 2 O→O 2 +4HF+2Kr

Application of inert gases:

Inert gases find quite a variety of practical applications. In particular, the role of helium in obtaining low temperatures is extremely important, since liquid helium is the coldest of all liquids. Artificial air, in which nitrogen is replaced by helium, was first used to ensure the breathing of divers. The solubility of gases increases greatly with increasing pressure, therefore, when a diver descends into water and is supplied with ordinary air, the blood dissolves more nitrogen than under normal conditions. During ascent, when the pressure drops, dissolved nitrogen begins to be released and its bubbles partially clog small blood vessels, thereby disrupting normal blood circulation and causing attacks of “caisson sickness.” Thanks to the replacement of nitrogen with helium, painful effects are sharply weakened due to the much lower solubility of helium in the blood, which is especially noticeable at high pressures. Working in an atmosphere of “helium” air allows divers to descend to great depths (over 100 m) and significantly extend their stay under water.

Since the density of such air is approximately three times less than that of normal air, it is much easier to breathe. This explains the great medical importance of helium air in the treatment of asthma, suffocation, etc., when even short-term relief of a patient’s breathing can save his life. Similar to helium, “xenon” air (80% xenon, 20% oxygen) has a strong narcotic effect when inhaled, which can be used medically.

Neon and argon are widely used in the electrical industry. When an electric current passes through glass tubes filled with these gases, the gas begins to glow, which is used to design illuminated inscriptions.

High-power neon tubes of this type are especially suitable for lighthouses and other signaling devices, since their red light is little blocked by fog. The color of the helium glow changes from pink through yellow to green as its pressure in the tube decreases. Ar, Kr and Xe are characterized by different shades of blue.

Argon (usually mixed with 14% nitrogen) is also used to fill electric lamps. Due to their significantly lower thermal conductivity, krypton and xenon are even better suited for this purpose: electric lamps filled with them provide more light with the same energy consumption, withstand overload better and are more durable than conventional ones.

Editor: Galina Nikolaevna Kharlamova

- (a. inert gasses; n. Inertgase, Tragergase; f. gaz inertes; i. gases inertes) noble, rare gases, monatomic gases without color and odor: helium (He), neon (Ne) ... Geological encyclopedia

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Books

• Set of tables. Chemistry. Nonmetals (18 tables), . Educational album of 18 sheets. Art. 5-8688-018 Halogens. Chemistry of halogens. Sulfur. Allotropy. Chemistry of sulfur. Sulfuric acid. Chemistry of nitrogen. Nitrogen oxides. Nitric acid is an oxidizing agent. Phosphorus.…
• Inert gases, Fastovsky V.G.. The book discusses the basic physical and physico-chemical properties of the inert gases helium, neon, argon, krypton and xenon, as well as their areas of application in chemical, metallurgical,…

Page 1
Noble (inert) gases.

	2 He	10 Ne	18Ar	36 Kr	54 Xe	86 Rn
Atomic mass	4,0026	20,984	39,948	83,80	131,30
Valence electrons	1s 2	(2)2s 2 2p 6	(8)3s 2 3p 6	(18)4s 2 4p 6	(18)5s 2 5p 6	(18)6s 2 6p
Atomic radius	0,122	0,160	0,192	0,198	0,218	0,22
Ionization energy E - → E +	24,59	21,57	15,76	14,00	12,13	10,75
Content in the earth's atmosphere, %	5*10 -4	1,8*10 -3	9,3*10 -1	1,1*10 -4	8,6*10 -6	6*10 -20

Noble (inert) gases are the elements of the main subgroup of group VIII: helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe) and radon (Rn) (a radioactive element). Each noble gas completes the corresponding period in the Periodic Table and has a stable, fully completed external electronic level - ns 2 n.p. 6 . – this explains the unique properties of the elements of the subgroup. The noble gases are considered to be completely inert. This is where their second name comes from – inert.

All noble gases are part of the atmosphere, their content in the atmosphere by volume (%) is: helium - 4.6 * 10 -4; argon – 0.93; krypton – 1.1* 10 -4; xenon – 0.8 * 10 -6 and radon – 6 * 10 -8. Under normal conditions, all of them are odorless and colorless gases, poorly soluble in water. Their boiling and melting points increase with increasing atomic sizes. The molecules are monatomic.

Properties	He	Ne	Ar	Kr	Xe	Rn
Atomic radius, nm	0,122	0,160	0,191	0,201	0,220	0,231
Ionization energy of atoms, eV	24,58	21,56	15,76	14,00	12,13	10,75
Boiling point, o C	-268,9	-245,9	-185,9	-153,2	-181,2	Near
Melting point, o C	-272.6(under pressure)	-248,6	-189,3	-157,1	-111,8	Near
Solubility in 1 liter of water at 0 o C, ml	10	-	60	-	50	-

§1. Helium

Helium was discovered in 1868. Using the method of spectral analysis of solar radiation (Lockyer and Frankland, England; Jansen, France). Helium was discovered on Earth in 1894. In the mineral kleveite (Ramsay, England).

From Greek ἥλιος - "Sun" (see Helios). It is curious that the name of the element used the ending “-i”, characteristic of metals (in Latin “-um” - “Helium”), since Lockyer assumed that the element he discovered was a metal. By analogy with other noble gases, it would be logical to give it the name “Helion”. In modern science, the name “helion” is assigned to the nucleus of a light isotope of helium - helium-3.

The special stability of the electronic structure of the atom distinguishes helium from all other chemical elements of the periodic table.

Helium is closest in physical properties to molecular hydrogen. Due to the negligible polarizability of helium atoms, it has the lowest boiling and melting points.

Helium is less soluble than other gases in water and other solvents. Under normal conditions, helium is chemically inert, but with strong excitation of atoms it can form molecular ions. Under normal conditions these ions are unstable; I capture the missing electron, they split into two neutral atoms. The formation of ionized molecules is also possible. Helium is the most difficult to compress of all gases.

Helium can be converted into a liquid state only at a temperature approaching absolute zero, i.e. -273.15. Liquid helium at a temperature of about 2K has a unique property - superfluidity, which in 1938. P.L. was opened. Kapitsa and theoretically substantiated by L.D. Landau, who created the quantum theory of convolution. Liquid helium exists in two modifications: helium I, which behaves like an ordinary liquid, and helium II, a superthermal conductive and supervolatile liquid. Helium II conducts heat 10 7 times better than helium I (and 1000 times better than silver). It has virtually no viscosity, instantly passes through narrow capillaries, and spontaneously overflows through the walls of blood vessels in the form of a thin film. He atoms in the superfluid state behave in much the same way as electrons in superconductors.

In the earth's crust, helium accumulates due to the decay of particles of radioactive elements, and is found dissolved in minerals and native metals.

Helium nuclei are extremely stable and are widely used to carry out various nuclear reactions.

In industry, helium is mainly isolated from natural gases by deep cooling. At the same time, it, as the lowest boiling substance, remains in the form of a gas, while all other gases condense.

Helium gas is used to create an inert atmosphere when welding metals, preserving food products, etc. Liquid helium is used in the laboratory as a coolant in low-temperature physics.

§2. Neon

Neon was discovered in June 1898 by Scottish chemist William Ramsay and English chemist Maurice Travers. They isolated this inert gas by “exclusion,” after oxygen, nitrogen, and all the heavier components of the air were liquefied. The element was given the simple name “neon,” which translated from Greek means “new.” In December 1910, French inventor Georges Claude made a gas-discharge lamp filled with neon.

The name comes from the Greek. νέος - new.

There is a legend according to which the name of the element was given by Ramsay's thirteen-year-old son, Willie, who asked his father what he was going to call the new gas, noting that he would like to give it a name novum(Latin - new). His father liked the idea, but felt that the title neon, derived from a Greek synonym, will sound better.

Neon, like helium, has a very high ionization potential (21.57 eV), so it does not form valence-type compounds. Its main difference from helium is due to the relatively greater polarizability of the atom, i.e. a slightly greater tendency to form intermolecular bonds.

Neon has very low boiling points (-245.9 o C) and melting points (-248.6 o C), second only to helium and hydrogen. Compared to helium, neon has a slightly higher solubility and ability to be adsorbed.

Like helium, neon, when strongly excited by atoms, forms molecular ions of the Ne 2 + type.

Neon is produced together with helium as a by-product during the process of liquefying and separating air. The separation of helium and neon is carried out by adsorption or condensation. The adsorbed method is based on the ability of neon, unlike helium, to be adsorbed by activated carbon cooled with liquid nitrogen. The condensation method is based on freezing out neon while cooling the mixture with liquid hydrogen.

Neon is used in electric vacuum technology to fill voltage stabilizers, photocells and other devices. Various types of neon lamps with a characteristic red glow are used in lighthouses and other lighting devices, in illuminated advertising, etc.

Natural neon consists of three stable isotopes: 21 Ne and 22 Ne.

In world matter neon It is distributed unevenly, but in general it ranks fifth in abundance in the Universe among all elements - about 0.13% by mass. The highest concentration of neon is observed on the Sun and other hot stars, in gaseous nebulae, in the atmosphere of outer planets of the solar system- Jupiter, Saturn, Uranus, Neptune. In the atmosphere of many stars, neon ranks third after hydrogen and helium. Of all the elements of the second period neon- the smallest population on Earth. Within the eighth group neon It ranks third in terms of content in the earth's crust - after argon and helium. Gas nebulae and some stars contain many times more neon than is found on Earth.

On Earth, the highest concentration of neon is observed in the atmosphere - 1.82 10 −3% by volume, and its total reserves are estimated at 7.8 10 14 m³. 1 m³ of air contains about 18.2 cm³ of neon (for comparison: the same volume of air contains only 5.2 cm³ of helium). The average neon content in the earth's crust is low - 7·10−9% by mass. In total, there are about 6.6 10 10 tons of neon on our planet. Igneous rocks contain about 10 9 tons of this element. As rocks break down, gas escapes into the atmosphere. To a lesser extent, the atmosphere is supplied with neon and natural waters.

Scientists see the reason for the neon poverty of our planet in the fact that the Earth once lost its primary atmosphere, which took with it the bulk of inert gases that could not, like oxygen and other gases, chemically bond with other elements into minerals and thereby gain a foothold on planet.

In 1892, the British scientist John Strett, better known to us as Lord Rayleigh ( cm. Rayleigh criterion), was engaged in one of those monotonous and not very exciting works, without which experimental science nevertheless cannot exist. He studied the optical and chemical properties of the atmosphere, setting himself the goal of measuring the mass of a liter of nitrogen with an accuracy that no one before him had been able to achieve.

However, the results of these measurements seemed paradoxical. The mass of a liter of nitrogen obtained by removing all other then known substances (such as oxygen) from the air and the mass of a liter of nitrogen obtained through a chemical reaction (by passing ammonia over copper heated to red heat) turned out to be different. It turned out that nitrogen from the air is 0.5% heavier than nitrogen obtained chemically. This discrepancy haunted Rayleigh. Having made sure that no errors were made in the experiment, Rayleigh published in the journal Nature letter asking if anyone could explain the reason for these discrepancies.

Sir William Ramsay (1852–1916), then working at University College London, responded to Rayleigh's letter. Ramsay suggested that there may be an undiscovered gas in the atmosphere, and he proposed using the latest equipment to isolate this gas. In the experiment, oxygen-enriched air mixed with water was exposed to an electrical discharge, which caused atmospheric nitrogen to combine with oxygen and dissolve the resulting nitrogen oxides in water. At the end of the experiment, after all the nitrogen and oxygen from the air had been exhausted, there was still a small bubble of gas left in the vessel. When an electric spark was passed through this gas and subjected to spectroscopy, scientists saw previously unknown spectral lines ( cm. Spectroscopy). This meant that a new element had been discovered. Rayleigh and Ramsay published their results in 1894, naming the new gas argon, from the Greek “lazy”, “indifferent”. And in 1904, both of them received the Nobel Prize for this work. However, it was not divided between scientists, as is customary in our time, but each received a prize in their field - Rayleigh in physics, and Ramsay in chemistry.

There was even some kind of conflict. At the time, many scientists believed that they "mastered" certain areas of research, and it was not entirely clear whether Rayleigh gave Ramsay permission to work on this problem. Fortunately, both scientists were wise enough to realize the benefits of working together, and by publishing their results together, they eliminated the possibility of an unpleasant battle for primacy.

Argon is a monatomic gas. Having a relatively larger atomic size, argon is more prone to forming intermolecular bonds than helium and neon. Therefore, argon in the form of a millet substance is characterized by slightly higher boiling points (at normal pressure) -185.9 °C (slightly lower than oxygen, but slightly higher than nitrogen) and melting points (-184.3 °C). 3.3 ml of argon dissolves in 100 ml of water at 20 °C; argon dissolves in some organic solvents much better than in water.

Argon forms intermolecular inclusion compounds - clathrates of the approximate composition Ar*6H 2 0 is a crystalline substance that decomposes at atmospheric pressure and a temperature of -42.8 °C. It can be obtained directly by the interaction of argon with water at 0°C and a pressure of the order of 1.5 * 10 7 Pa. With compounds H 2 S, SO 2, CO 2, HCl, argon gives double hydrates, i.e. mixed clathrates.

Argon is obtained by separating liquid air, as well as from waste gases of ammonia synthesis. Argon is used in metallurgical and chemical processes that require an inert atmosphere, in lighting engineering, electrical engineering, nuclear energy, etc.

Argon (along with neon) is observed on some stars and in planetary nebulae. In general, there is more of it in space than calcium, phosphorus, and chlorine, while on Earth the opposite relationships exist.

Argon is the third most abundant component of air after nitrogen and oxygen, its average content in the Earth's atmosphere is 0.934% by volume and 1.288% by mass, its reserves in the atmosphere are estimated at 4 10 14 tons. Argon is the most common inert gas in the Earth's atmosphere , 1 m³ of air contains 9.34 liters of argon (for comparison: the same volume of air contains 18.2 cm³ of neon, 5.2 cm³ of helium, 1.1 cm³ of krypton, 0.09 cm³ of xenon).

§4. Krypton

In 1898, the English scientist W. Ramsay isolated from liquid air (having previously removed oxygen, nitrogen and argon) a mixture in which two gases were discovered by the spectral method: krypton (“hidden”, “secret”) and xenon (“alien”, “ unusual").

From Greek κρυπτός - hidden.

Located in atmospheric air. It is formed during nuclear fission, including as a result of natural processes occurring in ores of radioactive metals. Krypton is obtained as a by-product from air separation.

Gaseous oxygen containing Kr and Xe from the condenser of the installation for producing O 2 is supplied for rectification in the so-called. a krypton column, in which Kr and Xe are extracted from gaseous O 2 when it is washed with reflux formed at the top of the condenser of the krypton column. The bottom liquid is enriched in Kr and Xe; it is then almost completely evaporated, the non-evaporated part is the so-called. called lean iron-xenon concentrate (less than 0.2% Kr and Xe) - continuously flows through the evaporator into the gas tank. With an optimal reflux ratio of 0.13, the degree of extraction of Kr and Xe is 0.90. The separated concentrate is compressed to 0.5-0.6 MPa and fed through a heat exchanger into a contact apparatus with CuO heated to ~1000 K to burn off the hydrocarbons contained in it. After cooling in a water refrigerator, the gas mixture is purified from impurities of CO 2 and water using KOH, first in scrubbers and then in cylinders. Burning and cleaning are repeated several times. once. The purified concentrate is cooled and continuously fed into the rectifier. column under pressure 0.2-0.25 MPa. In this case, Kr and Xe accumulate in the bottom liquid to a content of 95-98%. This so-called The raw krypton-xenon mixture is sent through a gasifier, an apparatus for burning hydrocarbons and a purification system into gas tanks. From the gas holder, the gas mixture enters the gasifier, where it is condensed at 77 K. Part of this mixture is subjected to fractional evaporation. As a result, the last purification from O 2 in a contact apparatus with CuO produces pure krypton. The remaining gas mixture is subjected to adsorption in devices with activator. coal at 200-210 K; in this case, pure krypton is released, and Xe and part of the krypton are absorbed by coal. Adsorbed Kr and Xe are separated by fractionated desorption. With a capacity of 20,000 m 3 /h of processed air (273 K, 0.1 MPa), 105 m 3 of krypton are obtained per year. It is also extracted from the methane fraction of purge gases in NH 3 production. They produce pure krypton (more than 98.9% by volume of krypton), technical. (more than 99.5% mixture of Kr and Xe) and krypton-xenon mixture (less than 94.5% krypton). Krypton is used to fill incandescent lamps, gas-discharge and X-ray tubes. The radioactive isotope 85 Kr is used as a source of b-radiation in medicine, to detect leaks in vacuum installations, as isotope tracer during corrosion studies, to monitor wear of parts. Krypton and its mixtures with Xe are stored and transported under a pressure of 5-10 MPa at 20°C in sealed steel cylinders black resp. with one yellow stripe and the inscription "Krypton" and two yellow stripes and the inscription "Krypton-xenon". Krypton was discovered in 1898 by W. Ramsay and M. Travers. Lit.

§5. Xenon

Discovered in 1898 by English scientists W. Ramsay and W. Rayleigh as a small admixture of krypton.

From Greek ξένος - stranger.

Melting point −112 °C, boiling point −108 °C, violet glow in the discharge.

The first inert gas for which true chemical compounds were prepared. Examples of connections could be xenon difluoride, xenon tetrafluoride, xenon hexafluoride, xenon trioxide.

Xenon is produced as a by-product when air separation. It is isolated from krypton-xenon concentrate (see Krypton). They produce pure (99.4% by volume) and high purity (99.9%) xenon. Xenon is obtained as a by-product of the production of liquid oxygen at metallurgical enterprises.

In industry, xenon is produced as a by-product of the separation of air into oxygen and nitrogen. After this separation, which is usually carried out by rectification, the resulting liquid oxygen contains small amounts of krypton and xenon. Further rectification enriches liquid oxygen to a content of 0.1-0.2% krypton-xenon mixture, which is separated adsorption on silica gel or by distillation. Finally, the xenon-krypton concentrate can be separated by distillation into krypton and xenon.

Due to its low prevalence, xenon is much more expensive than lighter inert gases.

Despite its high cost, xenon is indispensable in a number of cases:

• 
  Xenon is used to fill incandescent lamps, powerful gas-discharge and pulsed light sources (the high atomic mass of the gas in lamp bulbs prevents the evaporation of tungsten from the surface of the filament).

• 
  Radioactive isotopes (127 Xe, 133 Xe, 137 Xe, etc.) are used as radiation sources in radiography and for diagnostics in medicine, for detecting leaks in vacuum installations.

• 
  Xenon fluorides are used for passivation of metals.

• 
  Xenon, both in its pure form and with a small addition of cesium-133 vapor, is a highly efficient working fluid for electric propulsion (mainly ion and plasma) engines of spacecraft.

• 
  Since the end of the 20th century, xenon began to be used as a means for general anesthesia (quite expensive, but absolutely non-toxic, or rather, like an inert gas, it does not cause chemical consequences). The first dissertations on the technique of xenon anesthesia in Russia - 1993, as a therapeutic anesthesia, it is effectively used to relieve acute withdrawal states and treat drug addiction, as well as mental and somatic disorders.

• 
  Liquid xenon is sometimes used as a working medium for lasers.

• 
  Xenon fluorides and oxides are proposed as powerful oxidizers of rocket fuel, as well as components of gas mixtures for lasers.

• 
  In the 129 Xe isotope, it is possible to polarize a significant portion of the nuclear spins to create a state with co-directed spins - a state called hyperpolarization.

• 
  Xenon is used in the design of the Golay cell.

• 
  As chemical catalysts.

• 
  For transportation of fluorine, which exhibits strong oxidizing properties.

Xenon is relatively rare in the solar atmosphere, on Earth, and in asteroids and comets. The concentration of xenon in the atmosphere of Mars is similar to that on Earth: 0.08 ppm, although the abundance of 129 Xe on Mars is higher than on Earth or the Sun. Since this isotope is formed through radioactive decay, the findings may indicate that Mars lost its primary atmosphere, perhaps within the first 100 million years after the planet formed. Jupiter, by contrast, has an unusually high concentration of xenon in its atmosphere—almost twice that of the Sun.

Xenon is in earth's atmosphere in extremely small quantities, 0.087±0.001 parts per million (μL/L), and is also found in gases emitted by some mineral springs. Some radioactive isotopes of xenon, such as 133 Xe and 135 Xe, are produced by neutron irradiation of nuclear fuel in reactors.

The English scientist E. Rutherford noted in 1899 that thorium preparations emit, in addition to α-particles, some previously unknown substance, so that the air around the thorium preparations gradually becomes radioactive. He proposed to call this substance an emanation (from the Latin emanatio - outflow) of thorium and give it the symbol Em. Subsequent observations showed that radium preparations also emit a certain emanation, which has radioactive properties and behaves like an inert gas.

Initially, the emanation of thorium was called thoron, and the emanation of radium was called radon. It was proven that all emanations are actually radionuclides of a new element - an inert gas, which corresponds to atomic number 86. It was first isolated in its pure form by Ramsay and Gray in 1908, they also proposed to call the gas niton (from the Latin nitens, luminous ). In 1923, the gas was finally named radon and the symbol Em was changed to Rn.

Radon is a radioactive monatomic gas, colorless and odorless. Solubility in water 460 ml/l; in organic solvents and in human adipose tissue, the solubility of radon is tens of times higher than in water. Gas penetrates well through polymer films. Easily adsorbed by activated carbon and silica gel.

Radon's own radioactivity causes it to fluoresce. Gaseous and liquid radon fluoresces with blue light, while solid radon when cooled to nitrogen temperatures The fluorescence color becomes first yellow, then red-orange.

Radon forms clathrates, which, although they have a constant composition, do not contain chemical bonds involving radon atoms. With fluorine, radon at high temperatures forms compounds of the composition RnF n, where n = 4, 6, 2. Thus, radon difluoride RnF 2 is a white non-volatile crystalline substance. Radon fluorides can also be produced by the action of fluorinating agents (for example, halogen fluorides). At hydrolysis of tetrafluoride RnF 4 and hexafluoride RnF 6 form radon oxide RnO 3 . Compounds with the RnF + cation were also obtained.

To obtain radon, air is blown through an aqueous solution of any radium salt, which carries with it the radon formed during the radioactive decay of radium. Next, the air is carefully filtered to separate microdroplets of the solution containing the radium salt, which can be captured by the air current. To obtain radon itself, chemically active substances (oxygen, hydrogen, water vapor, etc.) are removed from a mixture of gases, the residue is condensed with liquid nitrogen, then nitrogen and other inert gases (argon, neon, etc.) are distilled from the condensate.

Radon is used in medicine to prepare radon baths. Radon is used in agriculture to activate animal feed [ source not specified 272 days ] , in metallurgy as an indicator when determining the speed of gas flows in blast furnaces and gas pipelines. In geology, measuring radon content in air and water is used to search for deposits of uranium and thorium, in hydrology - to study the interaction of groundwater and river waters. The dynamics of radon concentration in groundwater can be used to predict earthquakes.

It is part of the radioactive series 238 U, 235 U and 232 Th. Radon nuclei constantly arise in nature during the radioactive decay of parent nuclei. The equilibrium content in the earth's crust is 7·10−16% by mass. Due to its chemical inertness, radon relatively easily leaves the crystal lattice of the “parent” mineral and enters groundwater, natural gases and air. Since the longest-lived of the four natural isotopes of radon is 222 Rn, it is its content in these environments that is maximum.

The concentration of radon in the air depends primarily on the geological situation (for example, granites, which contain a lot of uranium, are active sources of radon, while at the same time there is little radon above the surface of the seas), as well as on the weather (during rain, microcracks through which radon comes from the soil, are filled with water; snow cover also prevents radon from entering the air). Before the earthquakes, an increase in radon concentration in the air was observed, probably due to a more active exchange of air in the ground due to an increase in microseismic activity.

(Galina Afanasyevna – HELP with krypton, xenon, argon! can I add something else? And what should I write next?)

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Noble gases (inert or rare gases) is a group of chemical elements with similar properties: under normal conditions they are monatomic gases. These are chemical elements that form the main subgroup of the 8th group of the Mendeleev periodic system.

Under normal conditions, these are colorless, tasteless and odorless gases, poorly soluble in water, do not ignite under normal conditions, with very low chemical reactivity. Their melting and boiling points naturally increase with increasing atomic number.

Among all the noble gases, only Rn There are no stable isotopes and it is the only radioactive chemical element.

Rare (inert) gases are:

• helium ( He) (atomic number 2),
• neon ( Ne) (10),
• argon ( Ar) (18),
• krypton ( Kr) (36),
• xenon ( Xe) (54)
• radioactive radon ( Rn) (86).

Recently, ununoctium ( Uuo) (118).

All inert gases complete the corresponding period in the Periodic Table and have a completely completed, stable external electronic level.

Noble gases have an electron configuration ns 2 n.p. 6 (for helium 1s 2) and they form group VIIIA. With increasing atomic number, the radii of atoms and their ability to polarizability increase, which leads to an increase in intermolecular interactions, to an increase T pl And T bale, to improve the solubility of gases in water and other solvents. For inert gases, there are such well-known groups of compounds: molecular ions, inclusion compounds, valence compounds.

Noble gases belong to the latter; they occupy the first 6 periods and belong to the 18th group in the periodic table of chemical elements. Flerovium, an element of group 14, shows some properties of noble gases, so it can replace ununoctium in the periodic table. Noble gases are chemically inactive and can only take part in chemical reactions under extreme conditions.

Colors and spectra of noble gases.

Colors and spectra of noble gases. The first line of the table shows noble gases in flasks through which a current is passed, the second - the gas itself in a tube, the third - in tubes that depict the designation of an element in the periodic table.

Helium	Neon	Argon	Krypton	Xenon

Prevalence of inert (rare) gases in nature.

Due to the fact that noble gases are chemically inert, they could not be detected for quite a long time, and their discovery took place only in the 2nd half of the 19th century.

Helium- is the second (after hydrogen) most abundant element in the Universe; in the earth’s crust, the content of helium is only 1 × 10-6 mass. %. Helium is a product of radioactive decay and is found in rock cavities and natural gas.

All noble gases are components of air. In 1 m 3 of air there are 9.3 argon, 18 ml of neon, 5 ml of helium, 1 ml of krypton and 0.09 ml of xenon. The sun is approximately 10% helium, formed from hydrogen by nuclear fusion reaction:

(β + - positron, - antineutrino). The helium lines, which were first discovered in 1868, appear quite intensely in the solar radiation spectrum. On Earth, helium was found only in 1895 during a spectral analysis of gases released when the mineral kleveite is dissolved in acids U 2 O 3. Uranium, which is part of the mineral, spontaneously decays according to the equation:

238 U→ 234 Th + 4 He.

They are present in small quantities in the air and some rocks, as well as in the atmospheres of some giant planets.

The industrial use of inert gases is based on their low chemical activity or specific physical properties.

Some characteristics of elements of subgroup VIIIA (inert gases).

Element	Radius avolumes, nm
Helium Not
Neon Ne
Argon Ar
Krypton TOr	3d 10 4s 2 4р 6
Xenon Heh	[Kr]4d 10 5s 2 5p 6
Radon Rn	[He]4f 1 4 5d 10 6s 2 6p 6

Plan.

1. Physical properties.
2. Chemical properties.
3. History of the discovery of inert gases.
4. Application area.
5. Effect on the human body.

1. Physical properties of inert gases.

Inert gases are colorless and odorless. And they are monatomic. Noble gases are considered noble gases. They have higher electrical conductivity (compared to others) and glow brightly when current passes through them.

Neon is a fiery red light, as its brightest lines are in the red region of the spectrum.

Helium has a bright yellow light, this is explained by the fact that in its relatively simple spectrum, the double yellow line predominates over all others.

Noble gases have lower liquefaction and freezing points than other gases of the same molecular weight. This is due to the saturated nature of the atomic molecules of noble gases.

2. Chemical properties of inert gases.

Inert gases have very low chemical activity, which is explained by the rigid eight-electron configuration of the outer electron layer. As is known, with an increase in the number of electronic layers, the polarizability of atoms increases. Therefore, it should increase when going from helium to radon.

For a long time, scientists did not find conditions at all under which noble gases could interact chemically or form true chemical compounds. Their valency was zero. And they decided to consider the new group of chemicals zero.

But in 1924, the idea was expressed that some compounds of heavy inert gases (in particular, xenon fluorides and chlorides) are thermodynamically quite stable and can exist under normal conditions. In theory, when studying the electronic structure of the shells of krypton and xenon from the standpoint of quantum mechanics, it turned out that these gases are able to form stable compounds with fluorine.

But time passed, and in practice all experiments in this area ended in failure. Xenon fluoride did not work. Gradually they came to the conclusion that this was not possible and the experiments stopped.

Only in 1961, Bartlett, an employee of one of the universities in Canada, studying the properties of platinum hexafluoride, a compound more active than fluorine itself, established that the ionization potential of xenon is lower than that of oxygen (12, 13 and 12, 20 eV, respectively), and oxygen formed a compound with the composition O2PtF6... with platinum hexafluoride...

At room temperature, Bartlett conducted an experiment and from gaseous platinum hexafluoride and gaseous xenon he obtained a solid orange-yellow substance called xenon hexafluoroplatinate XePtF6..

When heated in a vacuum, hexafluoroplatinate XePtF6 sublimes without decomposition. Hydrolyzes in water, releasing xenon:

2XePtF6 + 6H2O = 2Xe + O2 + 2PtO2 + 12HF

While studying the new substance, Bartlett came to the conclusion that the behavior of hexafluoroplatinate is no different from the behavior of ordinary chemical compounds.

Bartlett's work made it possible to establish that xenon, depending on the reaction conditions, is capable of forming two different compounds with platinum hexafluoride: XePtF6 and Xe(PtF6)2. But when these compounds are hydrolyzed, the same end products are obtained.

In 1962, Bartlett gives a presentation.

And just three weeks after his experiments, the experiment was repeated by a group of American researchers at the Argonne National Laboratory, led by Chernik. Scientists were the first to succeed in synthesizing similar xenon compounds with ruthenium, rhodium and plutonium hexafluorides.

So, the first five xenon compounds: XePtF6, Xe (PtF6)2, XeRuF6, XeRhF6, XePuF6

The myth about the absolute inertness of gases has not been confirmed.

They decided to test the existing hypothesis about the possibility of direct interaction of xenon with Fluorine.

For this purpose, a mixture of gases (1 part xenon and 5 parts fluorine) was placed in a nickel vessel, as the most resistant to the action of fluorine, and heated under relatively low pressure.

An hour later, the vessel was sharply cooled and the gas was pumped out. The remaining gas turned out to be nothing more than fluorine. All xenon reacted!

Afterwards, colorless crystals of xenon tetrafluoride XeF4 were found in the opened vessel.

This is a stable compound, its molecule has the shape of a square with fluorine ions at the corners and xenon in the center.

Xenon tetrafluoride XeF4 fluorides mercury, platinum (but only dissolved in hydrogen fluoride): XeF4 + 2Hg = Xe + 2HgF2

The remarkable thing is that by changing the reaction conditions, it is possible to obtain not only XeF4, but also other fluorides, for example XeF2, XeF6.

V. M. Khutoretsky and V. A. Shpansky, Soviet chemists, showed that harsh conditions are not at all necessary for the synthesis of xenon difluoride.

They proposed a method where a mixture of xenon and fluorine (in a molecular ratio of 1:1) is fed into a vessel made of nickel or stainless steel, and when the pressure increases to 35 atm, a spontaneous reaction begins.

XeF2 is the only xenon fluoride produced by applying an electric discharge to a mixture of xenon and carbon tetrafluoride, without the use of elemental fluorine.

Pure XeF2 is obtained by irradiating xenon and fluorine with ultraviolet light.

XeF2 difluoride has a sharp, specific odor.

The solubility of difluoride in water is low. Its solution is a strong oxidizing agent. Gradually it self-decomposes into xenon, oxygen and hydrogen fluoride. In an alkaline environment, decomposition occurs especially quickly.

The method for the synthesis of xenon difluoride, based on the effect of ultraviolet radiation (wavelength of the order of 2500-3500 A) on a mixture of gases, is of great theoretical interest.

Radiation causes fluorine molecules to split into free atoms. And this is precisely the reason for the formation of difluoride, because atomic fluorine is unusually active.

To obtain xenon hexafluoride XeF6, more stringent conditions are required: 700 ° C and 200 atm. Under such conditions, in a mixture of xenon and fluorine, in a ratio of 1:4 to 1:20, almost all xenon is converted into XeF6.

Xenon hexafluoride is extremely active and decomposes explosively.

Reacts easily with alkali metal fluorides (except LiF): XeF6 + RbF = RbXeF7

Already at 50°C this salt decomposes: 2RbXeF7 = XeF6 + Rb2XeF8

Higher fluoride XeF8 is stable only at temperatures below minus 196° C.

If previously the noble gases were separated into a separate zero group, which fully corresponded to the idea of ​​their valency, then the synthesis of the first xenon compounds raised the question of the place of noble gases in the periodic table. It was decided to transfer the inert gases to group VIII when its higher fluoride became known, in which the valency of xenon is eight, which is quite consistent with the structure of its electron shell.

All currently known xenon compounds are obtained from its fluorides. It has not yet been possible to force xenon to react without the participation of fluorine (or some of its compounds).

The interaction of xenon fluorides with water has been well studied.

When XeF4 is hydrolyzed in an acidic environment, xenon oxide XeO3 is formed - colorless crystals that diffuse in air.

The XeO3 molecule has the structure of a flattened triangular pyramid with a xenon atom at the top.

This is an extremely unstable compound; when decomposed, the power of the explosion approaches the power of a TNT explosion. Therefore, a few hundred milligrams of XeO3 are enough for the desiccator to be blown into pieces.

In the future, it is planned to use xenon trioxide as an explosive. Such explosives would be very convenient, because all the products of an explosive reaction are gases. In the meantime, using xenon trioxide for this purpose is too expensive due to small reserves in the atmosphere and technical difficulties.

To obtain 1 m3 of xenon, 11 million m3 of air must be processed.

The unstable acid of hexavalent xenon H6XeO6 corresponding to the trioxide is formed as a result of the hydrolysis of XeF6 at a temperature of 0 ° C:

XeF6 + 6H2O = 6HF + H6XeO6

If Ba(OH)2 is quickly added to the products of this reaction, a white amorphous

precipitate Ba3XeO6. At 125°C it decomposes into barium oxide, xenon and oxygen.

Similar sodium and potassium xenonate salts were obtained.

Under the influence of ozone, a salt of the higher acid xenon, Na4XeO6, is formed from a solution of XeO3 in one-molar sodium hydroxide. Sodium perxenonate can be isolated as a colorless crystalline hydrate Na4XeO6 6H2O. The hydrolysis of XeF6 in sodium and potassium hydroxides also leads to the formation of perxenonates.

It is enough to treat the solid salt Na4XeO6 with a solution of lead, silver or uranyl nitrate and the corresponding perxenonates are obtained:

Ag4XeO6 - black, bXeO6 and (UO2) 2XeO6 - yellow.

Similar salts are produced by potassium, calcium, lithium, cesium. By reacting Na4XeO6 with anhydrous cooled sulfuric acid, an oxide corresponding to the higher acid of xenon is obtained - XeO4 tetroxide.

As in octafluoride, xenon has a valency of eight.

Solid tetroxide at temperatures above 0 ° C decomposes into xenon and oxygen, and gaseous (at room temperature) - into xenon trioxide, xenon and oxygen.

The XeO4 molecule has the shape of a tetrahedron with a xenon atom in the center. Depending on the conditions, the hydrolysis of xenon hexafluoride can proceed in two ways:

1. tetraoxyfluoride XeOF4 is obtained,
2. dioxyfluoride XeO2F2 is obtained.

Direct synthesis from elements leads to the formation of oxyfluoride XeOF2.

The reaction of xenon difluoride with anhydrous HC1O4 has recently been studied.

A new xenon compound, XeClO4, was obtained - a powerful oxidizing agent, as a result of this reaction, the most powerful of all perchlorates. Xenon compounds containing no oxygen have been synthesized.

These are double salts, products of the interaction of xenon fluorides with fluorides of antimony, arsenic, boron, tantalum: XeF2 SbF5, XeF6 AsF3, XeF6 BF3 and XeF2 2TaF5.

Finally, substances of the XeSbF6 type, stable at room temperature, and XeSiF6, an unstable complex, were obtained. To date, it has been established that radon also interacts with fluorine, forming non-volatile fluorides.

Difluoride KrF2 and tetrafluoride for krypton KrF4 were isolated and studied for properties reminiscent of xenon compounds. 4. History of the discovery of noble gases. The noble gases include helium, neon, argon, krypton, xenon and radon. In terms of their properties, they are unlike any other elements and in the periodic table they are located between typical metals and non-metals.

The history of the discovery of inert gases is of great interest: firstly, as a triumph of the quantitative methods of chemistry introduced by Lomonosov (the discovery of argon), and secondly, as a triumph of theoretical foresight (the discovery of other inert gases), based on the greatest generalization of chemistry - Mendeleev’s periodic law.

The discovery of the first noble gas, argon, by physicist Rayleigh and chemist Ramsay occurred at a time when the construction of the periodic system seemed complete and only a few empty cells remained in it.

Back in 1785, the English chemist and physicist G. Cavendish discovered some new gas in the air, unusually chemically stable. This gas accounted for approximately one hundred and twentieth of the volume of air. But Cavendish was unable to find out what kind of gas it was.

This experiment was remembered 107 years later, when John William Strutt (Lord Rayleigh) came across the same impurity, noting that the nitrogen in the air was heavier than the nitrogen isolated from the compounds. Having not found a reliable explanation for the anomaly, Rayleigh, through the journal Nature, turned to his fellow natural scientists with a proposal to think together and work on unraveling its causes...

Two years later, Rayleigh and W. Ramsay established that the nitrogen in the air actually contains an admixture of an unknown gas, heavier than nitrogen and extremely inert chemically.

The air was stripped of its oxygen using hot copper and then heated with pieces of magnesium in a tube. After a significant amount of nitrogen had been absorbed by magnesium, the density of the residue was determined.

The density turned out to be 15 times greater than the density of hydrogen, while the density of nitrogen was only 14 times greater. This density increased further as nitrogen was absorbed further, until it reached 18.

Thus, it was proven that air contains a gas whose density is greater than the density of nitrogen... We obtained 100 cm3 of this substance with a density equal to 19.9. It turned out to be a monatomic gas.

When they went public with their discovery, it was stunning. It seemed incredible to many that several generations of scientists, who performed thousands of air tests, overlooked its component, and even such a noticeable one - almost a percentage! By the way, it was on this day and hour, August 13, 1894, that argon received its name, which translated from Greek means “inactive.”

Helium was first identified as a chemical element in 1868 by P. Jansen while studying a solar eclipse in India. During a spectral analysis of the solar chromosphere, a bright yellow line was discovered, initially attributed to the spectrum of sodium, but in 1871 J. Lockyer and P. Jansen proved that this line does not belong to any of the elements known on earth. Lockyer and E. Frankland named the new element helium from the Greek. “helios”, which means sun.

At that time they did not know that helium was an inert gas, and assumed that it was a metal. And only almost a quarter of a century later, helium was discovered on earth. In 1895, a few months after the discovery of argon, W. Ramsay and almost simultaneously the Swedish chemists P. Kleve and N. Lenglet established that helium is released when the mineral kleveite is heated.

A year later, G. Keyser discovered an admixture of helium in the atmosphere, and in 1906 helium was discovered in the natural gas of oil wells in Kansas. In the same year, E. Rutherford and T. Royds established that a-particles emitted by radioactive elements are helium nuclei.

After this discovery, Ramsay came to the conclusion that there is a whole group of chemical elements that is located in the periodic table between the alkali metals and halogens. Using the periodic law and Mendeleev's method, the amount of unknown noble gases and their properties, in particular their atomic masses, were determined. This made it possible to carry out targeted searches for noble gases.

Ramsay and his collaborators looked for minerals, natural waters, and even meteorites in search of inert gases. However, everything was to no avail; the tests invariably turned out to be negative.

Meanwhile, there was new gas in them, but the methods used were not sensitive enough and these “microtraces” were not detected.

Having begun to explore the air, in just four subsequent years four new elements were discovered, and gases such as neon, krypton and xenon were even isolated from the air.

To do this, the air, previously purified from carbon dioxide and moisture, was liquefied and then began to slowly evaporate. During this procedure, lighter gases evaporate and the heavy inert gases remaining after evaporation are sorted.

The resulting fractions were subjected to various studies.

Let's consider spectral analysis as one of the methods of determining:

This simple procedure allows you to accurately identify noble gases by spectral lines.

To do this, the gas is placed in a discharge tube to which a current is connected.

When the first, lightest and lowest-boiling fraction of air was placed in the discharge tube, new lines were discovered in the spectrum, along with the known lines of nitrogen, helium and argon, of which red and orange were especially bright. They gave the light in the tube a fiery color. The history of the name of this gas is interesting:

When Ramsay observed, in another experiment, the spectrum of the newly obtained gas, his twelve-year-old son, who had already become a “fan” of his father’s work, entered the laboratory. Seeing the unusual glow, he exclaimed: “new one!” , which means “new” in ancient Greek.

This is how the name of the gas “neon” arose.

It was not immediately possible to find the inert gases that complete the fourth, fifth and sixth periods of the periodic table, although after helium, neon and argon, completing the first three periods of the periodic table, were discovered, there was no doubt about their existence.

But by that time they had learned to obtain significant quantities of liquid air, largely thanks to the efforts of the English scientist Travers.

Even liquid hydrogen became available.

And Ramsay, together with Travers, were able to study the most difficult-to-volatile fraction of air, resulting after the distillation of helium, hydrogen, neon, oxygen, nitrogen and argon.

The remainder was dominated by raw (unrefined) krypton. And after pumping it out, a gas bubble invariably remained in the vessel. This gas gave a peculiar spectrum with lines in the regions from orange to violet and had a bluish glow in the electrical discharge. As is known, an element can be accurately identified by spectral lines. Both Ramsay and Travers had every reason to believe that a new inert gas had been discovered.

It was named xenon, which translated from Greek means “alien”. Indeed, in the krypton fraction of air he looked like a stranger.

In search of a new element and to study its properties, Ramsay and Travers processed about one hundred tons of liquid air. The xenon content in the atmosphere is extremely low, but air is practically the only and inexhaustible source of xenon (almost all xenon returns to the atmosphere).

The identity of xenon as a new chemical element was established by operating with only 0.2 cm3 of this gas.

Ramsay also deserves the credit for the discovery of the highest representative of inert gases. Using subtle technical techniques, he proved that the radioactive outflow from radium - the emanation of radium - is a gas that obeys all the laws of ordinary gases, is chemically inert and has a characteristic spectrum. Ramsay measured the rate of diffusion, which allowed the molecular weight of the gas to be determined to be approximately 220:

Based on the assumption that the nucleus of an atom of radium emanation is the remainder of the radium nucleus after the nucleus of a helium atom (a-particle) is ejected from it, it turns out that its charge should be equal to 88-2 = 86. So the new element must really be an inert gas. And its atomic weight is 226-4=222. It was officially decided to include a new group of chemical elements in the periodic table on March 16, 1900, after Ramsay’s meeting with Mendeleev.

1. Scope of application of inert gases.

Helium is a source of low temperatures.

Liquid helium is used in the study of many phenomena, such as superconductivity in the solid state. The thermal movement of atoms and free electrons in solids is practically absent at the temperature of liquid helium.

In addition, liquid helium is beneficial for cooling magnetic superconductors, particle accelerators and other devices. A rather unusual application of helium as a refrigerant is the process of continuously mixing 3He and 4He to create and maintain temperatures below 0.005 K

Helium gas is used as a light gas to fill balloons.

Since it is not flammable, it is used to fill the shell of an airship, adding it to hydrogen.

Helium is used as an inert medium for arc welding, especially magnesium and its alloys, in the production of Si, Ge, Ti and Zr, for cooling nuclear reactors.

Other uses of helium are for gas lubrication of bearings, in neutron counters (helium-3), gas thermometers, X-ray spectroscopy, food storage, and high voltage switches. Mixed with other noble gases, helium is used in outdoor neon advertising (in gas discharge tubes).

Large quantities of helium are used in breathing mixtures for work under pressure, since helium is less soluble in the blood than nitrogen. For example, during sea diving, when creating underwater tunnels and structures.

When using helium, the release of dissolved gas from the blood, decompression, is less painful for the diver, decompression sickness is less likely. The phenomenon of nitrogen narcosis, a constant and dangerous companion to the diver’s work, is completely eliminated.

He–O2 mixtures are used, due to their low viscosity, to relieve asthma attacks and to treat various respiratory diseases.

Argon is widely used in production.

Electric arc welding in an argon environment is very convenient, because In an argon jet, it is possible to weld thin-walled products and metals that were previously considered difficult to weld. It is believed that the electric arc in an argon atmosphere revolutionized metal cutting technology. The process was much faster, and it became possible to cut thick sheets of the most refractory metals.

By blowing argon through liquid steel, gas inclusions are removed from it. This improves the properties of the metal. Argon blown along the arc column (mixed with hydrogen) protects the cut edges and the tungsten electrode from the formation of oxide, nitride and other films. At the same time, it compresses and concentrates the arc on a small surface, causing the temperature in the cutting zone to reach 4000-6000 ° C.

In addition, the gas jet blows out the cutting products.

And when welding in an argon jet, there is no need for fluxes and electrode coatings, and therefore, there is no need to clean the seam from slag and flux residues.

The use of xenon is often based on its ability to react with fluorine.

In medicine, xenon has become widespread in fluoroscopic examinations of the brain. Used for intestinal candling (xenon strongly absorbs x-rays and helps to find lesions). However, it is completely harmless.

And the active isotope of xenon, xenon-133, is used in studying the functional activity of the lungs and heart.

High-pressure xenon lamps are widely used in lighting technology. The principle of operation is based on the fact that in such lamps an arc discharge shines in xenon, which is under a pressure of several tens of atmospheres.

The light in such lamps is bright and has a continuous spectrum - from ultraviolet to near-infrared, and it appears immediately after switching on.

6. Effect on the human body.

It would be natural to believe that noble gases should not affect living organisms, because they are chemically inert. However, this is not quite true. When mixed with oxygen, inhalation of higher inert gases leads a person to a state similar to alcohol intoxication. This narcotic effect of inert gases is caused by their dissolution in nerve tissues. And the higher the atomic weight of an inert gas, the higher its solubility, and the greater the narcotic effect it can have.

Bibliography.

1. Guzey L.S. Lectures on general chemistry
2. Akhmetov N.S. “General and inorganic chemistry”
3. Petrov M.M., Mikhilev L.A., Kukushkin Yu.N. "Inorganic chemistry"
4. Nekrasov B.V. “Textbook of General Chemistry”
5. Glinka N.L. "General chemistry"