INERTE GAS COMPOUNDS

And There are six rare or noble gases: helium He, neon Ne, argon Ar, krypton Kr, xenon Xe, and radon Rn. Argon was the first of these gases to be discovered. The discovery dates back to 1894, it was made by D. Rayleigh (1842–1919) and W. Ramsay (1852–1916). Due to its properties, the new element was named argon from the Greek. - inactive. This gas does not form chemical compounds. So, one of the outstanding experimental chemists of the XIX century. A. Moissan (1852-1907) in 1886 showed that argon does not react with the most active of the oxidizers - fluorine.

Other inert gases discovered after argon formed a group with it, which successfully fit into the table of D.I. Mendeleev, located on the border between the most pronounced metalloids and metals (in the zero group). After the synthesis of compounds of some inert gases and the study of their properties, these elements were placed in group VIII.

The first theories that tried to explain the structure of the atom, its model proposed by N. Bohr (1913), and the Kossel–Lewis (1916) electronic theory of valence based on this model, reinforced the already established opinion about the inertness of noble gases. The electronic configuration of the atoms of inert gases was just such, which, according to these theories, the atoms of other elements sought to acquire by reacting with each other.

The discovery of chemical compounds of inert gases was preceded by obtaining their compounds with water - hydrates. Back in 1896, two years after the discovery of argon, P. Willard, by subjecting the gas to pressure over water at 0 ° C, obtained a crystalline hydrate of the composition Ar 6H 2 O. Following him
R. de Forcran in the first quarter of the 20th century, using the same method, synthesized similar compounds of krypton Kr 5.75H 2 O and xenon Xe 5.75H 2 O. Then he obtained similar hydrates with heavy water.

In our country, interest in compounds of inert gases was shown by Corresponding Member of the Academy of Sciences of the USSR
B.A. Nikitin (1906–1952), who worked at the Radium Institute of the Academy of Sciences. In the 1940s he received not only hydrates of radon, neon and argon, but also the same type of compounds of radon and xenon with phenol and toluene of the general formula R 2C 6 H 5 OH and R 2C 6 H 5 CH 3 (where R is an inert gas). He called these compounds molecular, because he believed that in them the connection is carried out not due to the interaction between atoms, but as a result of intermolecular attraction, and he turned out to be right. In 1948, G. Powell introduced the term "clathrates" for such compounds. Nikitin pointed out that the ability to form molecular compounds increases from argon to radon and can be used to separate inert gases. Research in the field of clathrates contributed to the development of the laboratory experiment technique with inert gases and became a significant step towards the synthesis of their true chemical compounds.

Assumptions about the possibility of synthesizing compounds of inert gases were made already at the beginning of the 20th century. In 1916, W. Kossel, based on the value of their ionization potentials, pointed out that the formation of chemical compounds in xenon and krypton should be expected. Unlike lighter elements, they have larger atomic radii, the outer electrons capable of forming a bond are farther from the nucleus, and therefore less energy is required to remove them.

In 1924, the Austrian researcher A. von Antropov, who even then, contrary to generally accepted opinion, attributed noble gases to group VIII of the periodic system, suggested that they are capable of forming chemical compounds and that the highest valency in these compounds should be equal to the group number, t .e. eight. He considered the synthesis of compounds containing halogens and heavy inert gases to be the most probable.

In 1933, Linus Pauling, a professor at the California Institute of Technology, the future Nobel Prize winner in chemistry (1954), based on the radii of the proposed ions of inert gases, spoke more confidently than his predecessors about the possibility of synthesizing their compounds and even indicated their composition, namely, the krypton hexafluorides KrF 6 and xenon XeF 6 , to which he attributed sufficient stability, and the unstable xenon octafluoride XeF 8 . In addition, Pauling pointed out the possibility of synthesizing an acid of the composition H 4 XeO 6 and its salts. As we shall see below, most of the predictions were subsequently confirmed. Pauling's confidence was so great that he instructed his colleague D. Yost to carry out the synthesis of xenon fluoride. The resulting compound was subjected to spectral analysis. The spectroscopists made a cautious assumption about the formation of a xenon compound, but for chemists it did not seem convincing enough.

In the same 1933, the Italian researcher G. Oddo also attempted to obtain xenon fluoride by passing an electric discharge through a mixture of these gases. The compound was obtained, but the scientist was unable to purify it from the corrosion products of the vessel in which the reaction was carried out, and this pushed back the date of its discovery by thirty years.

In 1935, G. Booth and C. Wilson reported on the interaction of boron trifluoride with argon. An attempt to repeat this reaction in 1948 was unsuccessful. Gradually, interest in the synthesis of compounds of inert gases fell.

H Despite all this, before 1962 it is impossible to speak of any compounds of noble gases. They were still inert. Everything changed in May 1962, when the Englishman Neil Bartlett, who worked at the University of British Columbia (Canada, Vancouver), published his first report on the synthesis of a xenon compound with platinum hexafluoride Xe. At that time, the scientist was interested in the chemistry of inorganic fluorine compounds. It should be noted that already by 1960 achievements in the field of fluorine chemistry were very significant. Thus, the separation of uranium into isotopes U-235 and U-238 was carried out using a gaseous compound of uranium hexafluoride. There has been renewed interest in the synthesis and study of fluorides of other elements, in particular the elements of the platinum group.

Bartlet together with D. Loman in 1961 synthesized platinum hexafluoride PtF 6 . Researchers have long kept platinum hexafluoride crystals in contact with air. Gradually their color changed from red to orange. The analysis showed that a new, hitherto unknown O 2 compound, oxygenyl hexafluoroplatinate, was formed, a compound containing a singly charged positive oxygen ion, the oxygenyl cation O 2 + . Detailed studies of this compound confirmed its composition. From which it followed that platinum hexafluoride has an unusually high oxidizing power and can pull an electron even from oxygen. The value of electron affinity* for this compound turned out to be 6.8 eV, i.e. twice as high as that of fluorine, the strongest oxidizing agent known at that time. Bartlet had the idea to use platinum hexafluoride to oxidize an inert gas. Comparing the first ionization potentials of an oxygen molecule and atoms of inert gases, he came to the conclusion that the ionization potential of xenon comes closest to him. These potentials for oxygen and xenon are 12.20 and 12.129 eV, respectively.

In 1962, Bartlet reported that by reacting xenon with gaseous platinum hexafluoride, he obtained a yellow solid with the composition Xe.

Bartlet's message did not arouse confidence among the majority of chemists, and, consequently, interest. Only for a narrow circle of scientists who studied fluorine and inert gases did it become an incentive for a wide development of research in this area, which began in the same 1962, immediately after the first publication. In the United States, a group of scientists from the Argonne National Laboratory set to work. In the Soviet Union, in Moscow, this problem was taken up by a group at the IV Kurchatov Institute of Atomic Energy; in Leningrad - scientists from the Institute of Nuclear Physics
them. B.P. Konstantinov, as well as the Radium Institute.

Currently, for xenon, compounds of several oxidation states are reliably known: +1, +2, +4, +6 and +8. The compounds corresponding to the oxidation state +1 include, first of all, the first compound synthesized by Bartlett Xe. In addition, he obtained similar compounds with ruthenium, rhodium and palladium. Following Bartlet in the same year, D. Morton and V. Falcone, acting on xenon tetrafluoride XeF 4 rays at low temperature, obtained the XeF radical, where xenon was also in the +1 oxidation state. The same researchers showed that a similar radical is formed during the hydrolysis of xenon difluoride XeF 2 . The existence of the radical was convincingly confirmed in 1968 by our compatriot Academician V.A. Legasov.

XeF 2 difluoride was the first to be obtained from xenon compounds corresponding to the +2 oxidation state. This compound was independently synthesized in 1962 by S.L. Chernik in the United States and R. Hoppe in Germany. In subsequent years, two directions in the synthesis of compounds of inert gases can be distinguished. Firstly, a search was made for new fluorinating agents and energy sources to excite the reaction, and secondly, the technology of the process was improved. A significant contribution was made by our compatriots. V.M. Khutoretsky and V.A. Shpansky found that the formation of xenon difluoride can also proceed under rather mild conditions, if a mixture of gases F 2 and Xe in a volume ratio of 1: 1 is injected at a pressure of 35 atm into a stainless steel vessel.

In a short time, the physicochemical properties of xenon difluoride were studied. It is colorless crystals with a characteristic odor. XeF 2 dissolves slightly in water, while it is hydrolyzed, releasing Xe, HF, and oxygen. By nature, XeF 2 is a very strong oxidizing agent. Hydrolysis was studied in detail by Legasov. In addition, he determined the melting point of XeF 2, which is 140 °C. A.V. Nikolaev determined the solubility of XeF 2 in some solvents. Bartlet and D. Musher independently showed that in xenon difluoride one fluorine atom or both can be replaced at once by such substituents as
OSO 2 F, OTeF 5 , and some others. All such compounds are unstable and easily explode, with the exception of Xe(OTeFS).

In the late 1960s - early 1970s. it was found that Xe(II) in the form of a positive singly charged XeF + cation is able to form ordinary coordination compounds with singly charged anions of the RF 6 – type (where R = As, Sb, Nb, Ta, Ru, V), which are solid, colored substances. In total, more than 20 compounds were studied. Not only Bartlet and his collaborators participated in the research, but also researchers from other countries: groups led by F. Sladka and V. A. Legasov. The scientists came to the conclusion that complex xenon fluorides by their nature are no exception among inorganic covalent halides of other elements.

M We considered xenon compounds in oxidation states +1 and +2. Now let's move on to compounds in the +4 oxidation state. Much less known compounds of this type. Xenon tetrafluoride XeF 4 was first obtained by G. Klassen and co-workers in the United States in 1962, shortly after the synthesis of xenon dichloride by Bartlet became known. The synthesis was carried out in a sealed nickel vessel at 400°C, a pressure of 6 atm, and a fivefold excess of fluorine. It is believed that this is the most easily obtained xenon compound. Subsequently, other, simpler methods have been proposed, but for obtaining large (gram) quantities of this substance, the first method is still considered preferable. True, it has one big drawback - the main product is always contaminated with xenon difluoride. To separate these fluorides, Bartlet in 1968 suggested treating the mixture with arsenic pentafluoride AsF 5 , with which Xe(IV) does not interact, but Xe(II), as we saw above, forms a complex compound. The end products are separated by sublimation under vacuum. In 1973, Legasov and his co-workers took out a patent for the separation of di- and tetrafluorides.

Xenon tetrafluoride is colorless crystals resistant to -radiation. It turned out to be more stable than xenon fluorides in the lower and higher oxidation states. In an aqueous solution, XeF 4 is hydrolyzed, forming XeOF 2 , XeO 3 , HF and O 2 . Some substituents are able to displace fluorine from tetrafluoride completely or partially. Thus, the crystalline compound Xe(OOCCF 3) 4 was obtained, but it was not possible to replace fluorine with chlorine by direct interaction.

Of the xenon compounds corresponding to the +6 oxidation state, XeF 6 hexafluoride was the first to be isolated. The synthesis was carried out in 1963 at once in several laboratories in the United States and Yugoslavia. The synthesis is carried out in a stainless steel vessel at a pressure of 400 atm, a temperature of up to 300°C, and a 20-fold excess of fluorine. Hexafluoride is usually contaminated with impurities of other xenon fluorides, and, most unpleasantly, with an admixture of explosive trioxide XeO 3. To separate from impurities, hexafluoride is absorbed on sodium fluoride, with which it forms a strong compound NaF XeF 6 .

The physical and chemical properties of xenon hexafluoride have been studied quite well. This compound is colorless, relatively stable crystals. Hydrolyzes in water to form xenon trioxide and hydrogen fluoride. In alkaline solutions, the process goes further until the formation of a free perxenate ion H 2 XeO 6 2– , which carries a double negative charge and is an Xe(VIII) compound. Xenon hexafluoride is a strong fluorinating agent. With its help, many fluorocarbons were obtained. Under the action of XeF 6 silicon dioxide is converted into gaseous silicon tetrafluoride. Therefore, it is impossible to work with xenon hexafluoride either in glass or quartz dishes, which are destroyed.

As already mentioned, XeO 3 oxide is known for the +6 oxidation state. This compound is yellow in color, it explodes easily and is quite dangerous to work with.

For Xe(VIII), XeO 4 tetroxide and salts corresponding to the acid H 4 XeO 6, synthesized in 1964, are known. We have already mentioned one of the methods for obtaining anions of this acid in solution. The other consists in passing ozone into a dilute alkaline solution containing Xe(VI). This method was used to obtain stable insoluble perxenate salts of the composition Na 4 XeO 6 8H 2 O, Na 4 XeO 6 6H 2 O, and Ba 2 XeO 6 1.5H 2 O. Xenon tetraoxide is unstable and persists only at a temperature of –196 °C. It is obtained by reacting sodium perxenate with sulfuric acid. Thanks to these compounds, corresponding to the oxidation state Xe +8, inert gases can now rightfully take their place in the eighth group of the periodic table of elements of D.I. Mendeleev.

Relatively recently, in 2003, a very interesting compound of xenon with the organic substance acetylene was synthesized in our country. The synthesis was carried out by employees of the Institute of Synthetic Polymer Materials of the Russian Academy of Sciences and the Physical and Chemical Institute named after L.Ya. Karpov.

To In addition to xenon compounds, compounds of the lighter inert gas krypton are also known. Its interaction with fluorine is much more difficult, and it was possible to obtain only difluoride of the composition KrF 2 . Krypton difluoride was synthesized by A.W. Grosse in the United States in 1963 by passing an electric discharge through a mixture of krypton and fluorine at a temperature of about 200 °C. Other methods for obtaining KrF 2 are currently known, but they are also carried out under harsh conditions.

Krypton difluoride is colorless in both solid and gaseous states. Very unstable and spontaneously decomposes at temperatures well below room temperature. This does not allow to establish its exact physico-chemical characteristics. In water, it decomposes according to the reaction:

KrF 2 + H 2 O \u003d Kr + 0.5O 2 + 2HF.

The instability of xenon difluoride is the reason for its strong fluorinating effect.

In 1965, a stable coordination compound KrF 2 Sb 2 F 5 was obtained. The reaction takes place in a glass vessel at a temperature of –20 °C. The compound is quite stable and has a melting point of 50 °C.

Krypton tetrafluoride KrF 4 was obtained by the same method as KrF 2 , but under somewhat different conditions. Under the action of Ba(OH) 2 on it, barium kryptonate BaKrO 4 was isolated, in which krypton is part of the anion. In this compound, krypton appears to be in the +6 oxidation state.

Convincing evidence in favor of the existence of krypton compounds belonging to other classes is not yet available.

As for radon compounds, due to its rapid decay (3.823 days) and the complexity of working with radioactive gas, its compounds were obtained in very small quantities and their composition is unreliable. In 1962, in the United States, P. Fields and his coworkers reported that they had received radon fluoride, which was attributed to the composition RnF 2 . Scientists based only on the fact that its reaction with water leads to products similar to those that arise during the hydrolysis of XeF 2. Radon difluoride is a white crystalline substance, stable in vacuum up to 250 °C.

The study of radon compounds can be of great importance for medicine, because. will allow you to create concentrated preparations of this gaseous element. In addition, knowledge of the chemistry of radon will help to remove it from the air of uranium mines.

Very little is known about argon compounds. In 2003, it was reported that Finnish researchers from the University of Helsinki had obtained the argon compound HArF by photolysis.

It has not yet been possible to obtain chemical compounds of helium and neon.

BARTLET Neil was born on September 15, 1932 in Newcastle (Great Britain). His father was a ship's carpenter, as were four generations of his Scottish ancestors. In the late 1920s, during the Great Depression, there was practically no work for shipbuilders, so soon after his marriage in 1928 to Anna Wok, he opened a grocery store. The starting capital for the family business was the funds that Anna managed to save up before her marriage, during her work as a salesman.

The family had three children. Neil Bartlet remembers his childhood as a happy time. And although his father died early (during the First World War, he suffered from the effects of poison gases and died in 1944), the family lived well thanks to the grocery store. The children also got their first entrepreneurial experience: Neil and his older brother Ken opened a small ice cream shop with the money saved from Sunday entertainment. Neal spent the proceeds from his small business on books on chemistry, which he was fascinated with at the time, and on equipment for a small home laboratory.

As Bartlet later recalled, his mother was a very knowledgeable and determined woman, well versed in business. She started out as a saleswoman in a shoe store and became one of the prominent figures in the trading industry, although in those days it was not easy for a woman. Having lost her father early and received a very modest education, she ran her grocery business so successfully that the family never lacked funds.

Having passed the special examinations for the course of elementary education, Neil entered the Heaton School for Boys. Subsequently, he considered that he was very lucky with the school: from the very beginning, emphasis was placed on the study of the natural sciences and laboratory experiments. He continued his studies in the home laboratory, replenishing it with equipment bought with pocket money and his share of the profits from the sale of ice cream. From school interest then grew professionalism, which became the basis for obtaining a degree at King's College, Durham.

At first, Neil wanted to become a biochemist and, upon entering college, provided the appropriate recommendations from the teachers necessary in order to study the chemistry of natural compounds. However, as he became more familiar with chemical science, he decided that he was more attracted to inorganic chemistry, and after graduating from college in 1954 (with a bachelor's degree), he began working in Dr. P. L. Robinson's Inorganic Chemistry Research Group. In 1958 he defended his thesis. A year earlier, there were changes in his personal life - he married Christina Cross.

A few months after the defense, Bartlet accepted an invitation to work from the University of British Columbia in Vancouver (Canada) and, together with Dr. G. Clark, launched work on the chemistry of fluorine there. Together with students and interns, he began to investigate the fluorides of platinum metals and germanium.

However, his own scientific research concerned a surprisingly volatile red substance, which he accidentally obtained in his dissertation research by fluorinating platinum salts in a glass vessel. Ultimately, Bartlet, together with his first graduate student Derek Loman, showed that the volatile substance is a dioxygenyl compound O 2 + - dioxygenyl hexafluoroplatinate, which is formed when PtF 6 is mixed with O 2:

O 2 + PtF 6 \u003d O 2.

This fact confirmed that platinum hexafluoride is the strongest oxidizing agent of all known compounds (it even oxidizes oxygen). It was on this and other hexafluorides that Bartlet focused his further research.

At the beginning of 1962, he drew attention to the fact that the ionization potentials of xenon and oxygen are close in magnitude, and that of radon is even somewhat lower. Since experiments with radon were not possible at the time, Bartlet obtained a sample of xenon, prepared some PtF 6 and tried to oxidize the xenon using a glass vessel and quartz apparatus.

Until 1962, helium, neon, argon, krypton, xenon, and radon were considered inert gases that could not form any compounds. On March 23, 1962, Dr. Bartlett removed the baffle separating the gaseous red platinum hexafluoride from the colorless xenon, and the two gases immediately reacted to form an orange-yellow substance. Bartlet later wrote: "I was so surprised by this sight that I ran out of the laboratory to call colleagues or students who could witness this event." However, it was Friday evening, and there was no one in the building ...

In the first article describing this historical experiment, it was recorded that the pressure in the reaction vessel decreased as a result of the interaction, which clearly indicated the formation of a new substance; its composition, as the researchers initially suggested, was + - . In subsequent publications by the Bartlet group, it was shown that in fact its composition should have been presented as follows: + -. After these works, experiments with xenon flooded in.

A huge step in the development of the chemistry of xenon fluoride compounds was made by the Argonne National Laboratory (USA). After a visit there in October 1962, Bartlet decided to make xenon oxide the main focus of his research. And soon he and his graduate student P.Rao obtained oxide. However, due to an accident caused by the explosion of the second sample, both were hospitalized for a month, and the identification of the explosive substance XeO 3 was carried out by other scientists.

In 1964, Bartlet received a professorship, and two years later he was invited to the post of professor of chemistry at Princeton University. He later recalled that he was sorry to leave the west coast, where he felt at home. So three years later, when he received an invitation to the post of professor of chemistry from the University of Berkeley in California, he gladly accepted it, and since then he has not left the west coast. At the same time, Bartlet spent a long time (1969–1999) doing research at the Lawrence Berkeley National Laboratory. In 2000, the scientist received American citizenship.

In 1970, the American Chemical Society awarded him the Inorganic Chemistry Prize for many years of work in the field of transition metal hexafluorides. For his research, he also received the Danni-Heinemann Prize from the German Academy of Sciences in Göttingen (1971), Henry Moissan (France, 1988), medals to them. Linus Pauling (1989), H. Davy (2002) and other awards. He was elected a member of the Royal Society of London (1973), the German Academy of Naturalists "Leopoldina" (1969), a foreign member of the French Academy of Sciences (1989) and the US National Academy of Sciences (1979).

In 1999, Bartlet left active research to spend more time with his family and indulge in his favorite hobbies. He paints with watercolors, carpenters, tends the garden, works on silver. After his departure, there was only one compound that he wanted to get, but did not get it - gold hexafluoride.

N.V. FEDORENKO

* Electron affinity is the energy that is released as a result of the attachment of one electron to an atom. The atom then becomes a negative ion.

Krypton, xenon and radon are characterized by lower ionization potentials than typical elements (He, Ne, Ar), so they are capable of producing compounds of the usual type. Only in 1962, N. Bartlett managed to obtain the first such compound - xenon hexafluoroplatinate Xe + |PtF 6 |. Following this, krypton and radon fluorides and their numerous derivatives were obtained. Information about some xenon compounds is given in table. 17.2.

Table 17.2

Characterization of xenon compounds

oxidation	connections	molecules	Structure	Some properties
		substance	Asymmetrical antiprism	Heat resistant up to 400°C
		Colorless liquid	square pyramid	sustainable
		Colorless crystals	Pyramidal	Explosive, hygroscopic, stable in solutions
		Colorless	tetrahedral	Explosive
		Colorless	Octahedral	Corresponds to 11, XeO g; , there are also acidic anions: Xe0 8 ~, H 2 XeO | "and H 3 XeO with

The conditions for obtaining compounds of noble gases are not quite simple from the point of view of conventional ideas.

Xenon difluoride XeF 2 is obtained by reacting Xe with F 2 at high pressure. The substance is soluble in water. In the presence of acids, the hydrolysis process proceeds slowly, and in the presence of alkalis, hydrolysis intensifies:

XeF 2 is a strong oxidizing agent, for example, when interacting with HC1, the reaction proceeds

Xenon tetrafluoride XeF 4 is formed by prolonged heating and high pressure (400°C and 607 kPa) from xenon and fluorine in a ratio of 1:5. Xenon tetrafluoride is identical in properties to XeFn but is resistant to hydrolysis. In moist air, it undergoes disproportionation:

Xenon hexafluoride XeF 6 can be obtained from XeF 4:

or directly from Xe and F 2 at 250° C. and over 5065 kPa. This compound has a high reactivity, which can be seen from the example of its interaction with quartz:

As a Lewis acid, XeF(i) readily reacts with alkali metal fluorides (except LiF) to form heptafluoro and octafluoroxenate aniopes:

Hydrolysis of XcF 6 may be accompanied by the formation of Xe0 3 and the corresponding unstable Lewis acid XeOF 4:

Xenon fluorides are oxidizing agents:

Xenon hexafluoroplatinate Xe is obtained by the interaction of PtF 6 and Xe at room temperature, i.e. from two gaseous substances an orange-yellow solid is formed, which is stable under normal conditions:

Xe[ PtF c | sublimates without decomposition. Under the action of water hydrolyzes:

Later, several more xenon compounds with ruthenium, rhodium and plutonium hexafluorides were obtained: Xe, Xe, Xe.

Xenon oxotetrafluoride XeOF 4 has amphoteric features, which can be judged by the corresponding cationic complexes, for example.

Xenon oxide (Y1) Xe0 3 is a white, non-volatile compound that forms stable aqueous solutions. The Xe0 3 molecule has the structure of a trigonal pyramid. In an alkaline medium, it forms a xenate (U1) ion:

HXe0 4 due to disproportionation gradually turns into perxenate (USh) ion:

Xenon oxide (XX) Xe0 4 has the shape of a tetrahedron with an Xe atom in the center. Xe0 4 is obtained from barium oxoxenate by the action of H 2 S0 4:

The perxenate ion XeOf forms stable salts - perxenates, among them Na 4 Xe0 8 -6H 2 0, Na 4 Xe0 G -8H 2 0, Ba 2 XeO G -1.5H 2 0 are stable, poorly soluble in water.

Krypton forms compounds that are identical in composition, structure and properties of molecules to xenon compounds. Thus, crystalline krypton difluoride is obtained under the influence of a quiet electric discharge on the reaction mixture at -183°C and a pressure of -2.7 kPa.

Krypton(H) fluoride, or krypton difluoride, KrF 2 is unstable at room temperature, but at -78°C it can be stored for a long time. KrF 2 is a very strong oxidizing agent. When acting on HC1, it displaces chlorine, and oxygen from water. Krypton compounds with transition metals have also been obtained: KrFMeF 6 . Compounds of a similar type were obtained both with arsenic and antimony: Kr 2 F 3 AsF 6 , Kr 2 F 3 SbF G and KrFSb 2 F u .

At present, a significant number of xenon compounds have been described. The chemistry of krypton compounds is also successfully developing. As for radon, due to its high a-radioactivity, obtaining and studying the properties of its derivatives is extremely difficult.

The use of noble gases. Helium, due to such properties as inertness, lightness, mobility and high thermal conductivity, is widely used. For example, it is safe to transfer flammable substances from one vessel to another using helium.

A fundamental contribution to the study of the properties of liquid helium was made by outstanding Russian physicists, Nobel Prize winners L. D. Landau and P. L. Kapitsa.

Biological studies have shown that the helium atmosphere does not affect the human genetic apparatus, since it does not affect the development of cells and the frequency of mutations. Breathing helium air (air in which nitrogen is partially or completely replaced by helium) enhances oxygen exchange in the lungs, prevents nitrogen embolism (caisson sickness).

Neon is often used in technology instead of helium. It is widely used for the manufacture of gas-light neon lamps.

Argon is more accessible than helium and neon. This gas is widely used in metallurgy, it is usually used in the hot processing of titanium, niobium, hafnium, uranium, thorium, alkali metals, where contact with oxygen, nitrogen, water and carbon monoxide (IV) is excluded. The method of electric arc welding in an argon medium has found wide application.

Krypton is mainly used in the manufacture of electric lamps.

Xenon is widely used in the production of xenon lamps, which are characterized by the correct color rendering. Xenon is a radiopaque substance widely used in fluoroscopy of the brain.

In the form of xenon fluorides, it is convenient to store and transport xenon and highly aggressive fluorine, which is of great environmental importance. Xenon oxides can be used as explosives or as strong oxidizers.

Although radon is radioactive, in ultramicrodoses it has a positive effect on the central nervous system, so it is used in balneology and physiotherapy (radon baths).

Summary

The electron shell of helium (it is the only one) has the Is 2 configuration, while for the rest of the elements the outer energy level is completed and contains eight electrons (configuration ...ns 2 np c>), which explains their extremely low activity. These elements are collectively referred to as noble gases. The old name of the elements of this group "inert gases" actually applies only to helium and neon, since the electronic structure of their atoms does not allow any possibility of the formation of covalent compounds, unlike the rest, for which it was possible to obtain chemical compounds.

Questions and tasks

• 1. Give the electronic configurations of noble gases and, based on this, explain why noble gases do not form diatomic molecules.
• 2. Why are helium and neon not capable of forming compounds with other elements?
• 3. What is the proposed mechanism for the formation of bonds involving krypton and xenon?
• 4. Describe the xenon compounds known to you.

Due to the completeness of the outer electronic level, the noble gases are chemically inert. Until 1962, it was believed that they did not form chemical compounds at all. In the Brief Chemical Encyclopedia (M., 1963, v. 2) it is written: "Inert gases do not give compounds with ionic and covalent bonds." By this time, some compounds of the clathrate type had been obtained, in which a noble gas atom is mechanically held in a framework formed by molecules of another substance. For example, under strong compression of argon over supercooled water, the crystal hydrate Ar 6H 2 0 was isolated. At the same time, all attempts to force noble gases to react even with the most energetic oxidizing agents (such as fluorine) ended in vain. And although theorists led by Linus Pauling predicted that the molecules of fluorides and xenon oxides could be stable, the experimenters said: "This cannot be."

Throughout this book, we try to emphasize two important ideas:

• 1) there are no immutable truths in science;
• 2) in chemistry, ABSOLUTELY EVERYTHING is possible, even what seems impossible or ridiculous for decades.

These ideas were perfectly confirmed by the Canadian chemist Neil Bartlett, when in 1962 he received the first chemical compound of xenon. That's how it was.

In one of the experiments with platinum hexafluoride PtF 6 Bartlett obtained red crystals, which, according to the results of chemical analysis, had the formula 0 2 PtF 6 and consisted of ions 0 2 and PtF 6 . This meant that PtF 6 is such a strong oxidizing agent that it takes electrons even from molecular oxygen! Bartlett decided to oxidize another spectacular substance and realized that it was even easier to take electrons from xenon than from oxygen (ionization potentials 0 2 12.2 eV and Xe 12.1 eV). He placed platinum hexafluoride in a vessel, launched a precisely measured amount of xenon into it, and a few hours later received xenon hexafluoroplatinate.

Immediately following this reaction, Bartlett carried out the reaction of xenon with fluorine. It turned out that when heated in a glass vessel, xenon reacts with fluorine, and a mixture of fluorides is formed.

Xenon fluoride^ II) XeF 2 is formed under the action of daylight on a mixture of xenon with fluorine at ordinary temperature

or during the interaction of xenon and F 2 0 2 at -120 ° C.

Colorless XeF 2 crystals are soluble in water. The XeF 2 molecule is linear. A solution of XeF 2 in water is a very strong oxidizing agent, especially in an acidic environment. In an alkaline environment, XeF 2 is hydrolyzed:

Xenon fluoride(H) XeF 4 is formed by heating a mixture of xenon with fluorine to 400 °C.

XeF 4 forms colorless crystals. The XeF 4 molecule is a square with a xenon atom in the center. XeF 4 is a very strong oxidizing agent and is used as a fluorinating agent.

When interacting with water, XeF 4 disproportionates.

Xenon Fluoride(Ch1) XeF 6 is formed from the elements when heated and pressurized with fluorine.

XeF 6 - colorless crystals. The XeF 6 molecule is a distorted octahedron with a xenon atom in the center. Like other xenon fluorides, XeF 6 is a very strong oxidizing agent and can be used as a fluorinating agent.

XeF 6 is partially decomposed by water:

xenon oxide(U I) Xe0 3 is formed during the hydrolysis of XeF 4 (see above). It is a white, non-volatile, highly explosive substance, highly soluble in water, and the solution has a slightly acidic reaction due to the following reactions:

Under the action of ozone on an alkaline solution of Xe0 3, a salt of xenonic acid is formed, in which xenon has an oxidation state of +8.

Xenon oxide (U1N) XeO 4 can be obtained by reacting barium perxenate with anhydrous sulfuric acid at low temperatures.

Xe0 4 is a colorless gas, highly explosive and decomposes at temperatures above 0 °C.

Of the compounds of other noble gases, KrF 2 , KrF 4 , RnF 2 , RnF 4 , RnF 6 , Rn0 3 are known. It is believed that similar compounds of helium, neon and argon are unlikely to ever be obtained in the form of individual substances.

Above we stated that in chemistry "everything is possible". Let us therefore report that compounds of helium, neon and argon exist in the form of so-called excimer molecules, i.e. molecules in which the excited electronic states are stable and the ground state is unstable. For example, upon electrical excitation of a mixture of argon and chlorine, a gas-phase reaction can proceed with the formation of an excimer ArCl molecule.

Similarly, in the reactions of excited noble gas atoms, a whole set of diatomic molecules can be obtained, such as He 2, HeNe, Ne 2, NeCl, NeF, HeCl, ArF, etc. All these molecules are unstable and cannot be isolated as individual substances, however, they can be registered and their structure studied using spectroscopic methods. Moreover, electronic transitions in excimer molecules are used to generate UV radiation in high-power excimer UV lasers.

The main subgroup of the eighth group of the periodic system is the noble gases - helium, neon, argon, krypton, xenon and radon. These elements are characterized by very low chemical activity, which gave reason to call them noble or inert gases. They only with difficulty form compounds with other elements or substances; chemical compounds of helium, neon and argon have not been obtained. Atoms of noble gases are not combined into molecules, in other words, their molecules are monatomic.

The noble gases complete each period of the system of elements. In addition to helium, all of them have eight electrons in the outer electron layer of the atom, forming a very stable system. The electron shell of helium, which consists of two electrons, is also stable. Therefore, noble gas atoms are characterized by high ionization energies and, as a rule, negative electron affinity energies.

In table. 38 shows some of the properties of noble gases, as well as their content in the air. It can be seen that the liquefaction and solidification temperatures of noble gases are the lower, the lower their atomic masses or serial numbers: the lowest liquefaction temperature for helium, the highest for radon.

Table 38. Some properties of noble gases and their content in the air

Until the end of the 19th century, it was believed that air consisted only of oxygen and nitrogen. But in 1894, the English physicist J. Rayleigh found that the density of nitrogen obtained from air (1.2572 ) is somewhat greater than the density of nitrogen obtained from its compounds (1.2505 ). Chemistry professor W. Ramsay suggested that the difference in density is caused by the presence of some heavier gas in atmospheric nitrogen. By binding nitrogen with red-hot magnesium (Ramsay) or by causing it to combine with oxygen (Rayleigh) by an electric discharge, both scientists isolated small amounts of a chemically inert gas from atmospheric nitrogen. Thus, an element unknown until that time, called argon, was discovered. Following argon, helium, neon, krypton and xenon, contained in the air in negligible amounts, were isolated. The last element of the subgroup - radon - was discovered in the study of radioactive transformations.

It should be noted that the existence of noble gases was predicted back in 1883, i.e. 11 years before the discovery of argon, by the Russian scientist II A. Morozov (1854-1946), who was imprisoned in 1882 for participating in the revolutionary movement by the tsarist government to the Shlisselburg fortress. N. A. Morozov correctly determined the place of noble gases in the periodic system, put forward ideas about the complex structure of the atom, about the possibility of synthesizing elements and using intra-atomic energy. N. A. Morozov was released from prison in 1905, and his remarkable predictions became known only in 1907 after the publication of his book Periodic Systems of the Structure of Matter, written in solitary confinement.

In 1926 N. A. Morozov was elected an honorary member of the USSR Academy of Sciences.

For a long time it was believed that atoms of noble gases were generally incapable of forming chemical bonds with atoms of other elements. Only relatively unstable molecular compounds of noble gases were known - for example, hydrates, formed by the action of compressed noble gases on crystallizing supercooled water. These hydrates belong to the clathrate type (see § 72); valence bonds do not arise in the formation of such compounds.

The formation of clathrates with water is favored by the presence of numerous cavities in the crystal structure of ice (see § 70).

However, during the last decades it has been established that krypton, xenon and radon are able to combine with other elements and, above all, with fluorine. So, by direct interaction of noble gases with fluorine (when heated or in an electric discharge), fluorides and are obtained. All of them are crystals that are stable under normal conditions. Xenon derivatives were also obtained in the degree of oxidation - hexafluoride, trioxide, hydroxide. The last two compounds exhibit acidic properties; so, reacting with alkalis, they form salts of xenonic acid, for example:.

For a long time, scientists believed that the noble gases could not form compounds because there was no room for more electrons in their electron shells, which contain valence electrons. This means that they cannot accept more electrons, making chemical bonding impossible. However, in 1933, Linus Pauling suggested that heavy noble gases could react with fluorine or oxygen, since they have atoms with the highest electronegativity. His guess turned out to be correct, and noble gas compounds were later obtained.

The noble gas compound was first obtained by Canadian chemist Neil Bartlett in 1962 by reacting platinum hexafluoride with xenon. The compound was assigned the formula XePtF 6 (as it turned out later - incorrect [ ]). Immediately after Bartlet's report, simple xenon fluorides were obtained in the same year. Since that time, the chemistry of noble gases has been actively developed.

Connection types

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Compounds of noble gases, where the noble gases are included in the crystal or chemical lattice, without the formation of a chemical bond, are called inclusion compounds. These include, for example, hydrates of inert gases, clathrates of inert gases with chloroform, phenols, etc.

Noble gases can also form compounds with endohedral fullerenes, when a noble gas atom is "pushed" into the inside of a fullerene molecule.

Complex compounds

Recently (2000) it has been shown that xenon can complex with gold (eg (Sb 2 F 11) 2) as a ligand. Complex compounds were also obtained, where xenon difluoride acts as a ligand.

Chemical compounds

In recent years, several hundred chemical compounds of noble gases (that is, having at least one noble gas-element bond) have been obtained. These are predominantly xenon compounds, since lighter gases are more inert, and radon has significant radioactivity. For krypton, a little more than a dozen compounds are known (mainly krypton difluoride complexes), for radon, fluoride of unknown composition is known. For gases lighter than krypton, only compounds in a matrix of solid inert gases (for example, HArF) are known that decompose at cryogenic temperatures.

For xenon, compounds are known where there are Xe-F, Xe-O, Xe-N, Xe-B, Xe-C, Xe-Cl bonds. Almost all of them are fluorinated to one degree or another and decompose when heated.

Links

• Khriachtchev, Leonid; Räsänen, Markku; Gerber, R. Benny. Noble-Gas Hydrides: New Chemistry at Low Temperatures // Accounts of Chemical Research (English) Russian: journal. - 2009. - Vol. 42, no. one . - P. 183 . -