The elements included in group VII of the periodic table are divided into two subgroups: the main one - the halogen subgroup - and the secondary one - the manganese subgroup. Hydrogen is placed in this same group, although its atom has a single electron at the outer, valence level and should be placed in group I. However, hydrogen has very little in common with both the elements of the main subgroup - the alkali metals, and the elements of the secondary subgroup - copper, silver and gold. At the same time, like halogens, it adds an electron in reactions with active metals and forms hydrides that have some similarities with halides.

The subgroup of halogens includes fluorine, chlorine, bromine, iodine and astatine. The first four elements are found in nature, the last one is obtained artificially and therefore has been studied much less than the other halogens. The word halogen means salt-forming. The elements of the subgroup received this name due to the ease with which they react with many metals, forming salts.

All halogens have an s2p5 outer shell structure. Therefore, they easily accept an electron, forming a stable noble gas electron shell (s2р6). Fluorine has the smallest atomic radius in the subgroup; for the rest it increases in the series F< Cl < Br < I < Аt и составляет соответственно 133; 181; 196; 220 и 270 пм. В таком же порядке уменьшается сродство атомов элементов к электрону. Галогены - очень активные элементы. Они могут отнимать, электроны не только у атомов, которые их легко отдают, но и у ионов и даже вытеснять другие галогены, менее активные, из их соединений. Например, фтор вытесняет хлор из хлоридов, хлор - бром из бромидов, а бром - иод из иодидов.

Of all the halogens, only fluorine, which is in period II, does not have an unfilled d-level. For this reason, it cannot have more than one unpaired electron and exhibits only a valence of -1. In the atoms of other halogens, the d-level is not filled, which allows them to have a different number of unpaired electrons and exhibit the -1, +1, +3, +5 and +7 valences observed in the oxygen compounds of chlorine, bromine and iodine.

The manganese subgroup includes manganese, technetium and rhenium. Unlike halogens, elements of the manganese subgroup have only two electrons at the outer electronic level and therefore do not exhibit the ability to attach electrons, forming negatively charged ions.

INCOMPLETE SECONDARY EDUCATION in the Russian Federation, a stage in obtaining completed secondary education; knowledge acquired in junior high school. Those who have completed the 8th (9th) grade continue their education in high school or in vocational education institutions.

Evgeniy Petrovich CHELYSHEV (b. 1921), Russian literary critic, academician of the Russian Academy of Sciences (1991; academician of the USSR Academy of Sciences since 1987). Works on problems of Indian literature.

ENTHUSIASM (Greek: enthusiasmos), inspiration, elation in the process of achieving a goal.

9 F 1s 2 2s 2 2p 5

17 Cl 3s 2 3p 5

35 Br 3d 10 4s 2 4p 5

53 I 4d 10 5s 2 5p 5

85 At 4f 14 5d 10 6s 2 6p 5

The 5 elements of the main subgroup of group VII have a common group name “halogens” (Hal), which means “salt-producing”.

The subgroup of halogens includes fluorine, chlorine, bromine, iodine and astatine (astatine is a radioactive element, little studied). These are the p-elements of the group of the periodic system of D.I. Mendeleev. At the outer energy level, their atoms have 7 electrons ns 2 np 5. This explains the commonality of their properties.

Properties of elements of the halogen subgroup

They easily add one electron each, exhibiting an oxidation state of -1. Halogens have this degree of oxidation in compounds with hydrogen and metals.

However, halogen atoms, in addition to fluorine, can also exhibit positive oxidation states: +1, +3, +5, +7. Possible values ​​of oxidation states are explained by the electronic structure, which for fluorine atoms can be represented by the diagram

Being the most electronegative element, fluorine can only accept one electron per 2p sublevel. It has one unpaired electron, so fluorine can only be monovalent, and the oxidation state is always -1.

The electronic structure of the chlorine atom is expressed by the diagram:

The chlorine atom has one unpaired electron in the 3p sublevel and the normal (unexcited) state of chlorine is monovalent. But since chlorine is in the third period, it has five more orbitals of the 3rd sublevel, which can accommodate 10 electrons.

In the excited state of the chlorine atom, electrons move from the 3p and 3s sublevels to the 3d sublevel (shown by arrows in the diagram). The separation (pairing) of electrons located in the same orbital increases the valence by two units. Obviously, chlorine and its analogues (except fluorine) can only exhibit odd variable valency 1, 3, 5, 7 and corresponding positive oxidation states. Fluorine has no free orbitals, which means that during chemical reactions there is no separation of paired electrons in the atom. Therefore, when considering the properties of halogens, it is always necessary to take into account the characteristics of fluorine and compounds.

Aqueous solutions of hydrogen compounds of halogens are acids: HF - hydrofluoric (fluoric), HCl - hydrochloric (hydrochloric), HBr - hydrobromic, HI - hydroiodic.

The identical structure of the outer electronic layer (ns 2 np 5) determines the great similarity of the elements.

Simple substances - non-metals F 2 (gas), Cl 2 (gas), Br 2 (l), l 2 (solid).

When forming covalent bonds, halogens most often use one unpaired p-electron available in an unexcited atom, exhibiting B = I.

Valence states of CI, Br, I atoms.

By forming bonds with atoms of more electronegative elements, atoms of chlorine, bromine and iodine can transition from the ground valence state to excited ones, which is accompanied by the transition of electrons to vacant orbitals of the d-sublevel. In this case, the number of unpaired electrons increases, as a result of which the CI, Br, I atoms can form a larger number of covalent bonds:

Difference between F and other halogens

In the F atom, the valence electrons are in the 2nd energy level, which has only s- and p-sublevels. This excludes the possibility of transition of F atoms to excited states, therefore fluorine in all compounds exhibits a constant B equal to I. In addition, fluorine is the most electronegative element, as a result of which it also has a constant c. O. -1.

The most important halogen compounds

I. Hydrogen halides HHal.

II Metal halides (salts of hydrohalic acids) are the most numerous and stable halogen compounds

III. Organohalogen compounds

IV. Oxygen-containing substances:

Unstable oxides, of which the existence of 6 oxides can be considered reliable (Cl 2 O, ClO 2, Cl 2 O 7, Br 2 O, BrO 2, I 2 O 5);

Unstable oxoacids, of which only 3 acids are isolated as individual substances (HClO 4, HlO 3, HlO 4);

Salts of oxoacids, mainly chlorites, chlorates and perchlorates.

10221 0

Group 17 includes F, Cl, Br, I, At (Tables 1 and 2). The word halogen (“halo” + “gen”) means “salt-forming.” All elements are non-metals. They have 7 electrons in the outer shell. Due to their high electronegativity and reactivity, they are not found in free form in nature. Due to the easy addition of an electron, they form halide ions and therefore exist in the form of diatomic molecules. Atoms in molecules are connected by a covalent bond as a result of sharing a pair of electrons, one from the atom. Halogen molecules are held together by weak van der Waals forces, which explains their high volatility.

Table 1 . Some physical and chemical properties of group 17 metals

	Name		Relates, at. weight	Electronic formula	Radius, pm	Main isotopes (%)
	Fluorine Fluorine [from lat. fluere - to flow]				covalent 58
	Chlorine Chlorine [from Greek. chloros - greenish]				covalent 99
	Bromine Bromine [from Greek. bromos - stench]			3d 10 4s 2 4p 5	Covalent 114.2	79 Vg* (50.69)
	Iodine Iodine [from Greek. iodes - purple]			4d 10 5s 2 5p 5	Covalent 133
	Astatine Astatine [from Greek. Astatos - unstable]			4f 14 5d 10 6s 2 6p 5

All halogens are toxic, have a characteristic pungent odor and color, the intensity of which increases towards the bottom of the group. This group consists of the most reactive elements of the Periodic Table. The atomic and ionic radii of halogens, as well as the bond lengths in molecules, increase towards the bottom of the group in the Periodic Table. On the contrary, the bond dissociation energy and its strength decrease, with the exception of fluorine.

Alkali metal halides (group 1) are compounds of the ionic type. In alkaline earth metal halides (group 2), in addition to ionic ones, there are compounds of a partially covalent type. As you move from left to right along a period, the halides of the elements become more covalent. The covalent nature of the halides also increases as one moves down the group. Moreover, if a metal can exist in several oxidation states, then its bond with the halide in the lowest of them is ionic in nature, and in the highest it is covalent. Both ionic and covalent divalent metal halides tend to crystallize into layered lattices. The exception is CCl 2, having a polymer structure. Halide ions are ligands in many complex ions, displacing less strong ligands such as water.

Silver halides are unstable in sunlight, decomposing into metal and halogen. This property is used in black and white photography. Bromides were the most photosensitive Ag. Hydrogen halides, which are among the most well-known strong acids, are widely used. The acidity of their aqueous solutions increases towards the bottom of the group. The exception is hydrogen fluoride. Its aqueous solution ( hydrofluoric acid) has slight acidity due to bond strength H - F and a small acid dissociation constant.

Table 2. Content in the body, toxic (TD) and lethal doses (LD) of group 17 metals

	In the earth's crust (%)	In the ocean (%)	In the human body
			Average (with body weight 70 kg)			Blood (mg/l)
							TD - 20 mg, LD - 2 g
							Toxic
							TD - 3 g, LD - >35 g
		(0.43-0.58)x10 -5			(0.05-5) x10 -5		TD - 2 mg, LD - 35-350 g
	Traces in some minerals						Toxic due to radioactivity

Fluorine (F) — in terms of prevalence, it ranks 13th among the elements of the earth’s crust, the most reactive element, the most powerful of the industrially produced oxidizing agents. In gaseous form it has a pale yellow color. In industry, its organic compounds, polymers and all salts are used, especially CaF 2 - as a flux in metallurgy, and AlF 3 - during production Al. Large quantities F 2 were produced in the nuclear industry to obtain U.F. 6 in nuclear fuel enrichment processes.

Close arrangement of atoms in a molecule F results in strong repulsion between nonbonding electrons, which explains the weakening of the bond in the molecule. Therefore, fluorine in the elemental state in the form F 2 is not found, but is present as a fluoride ion in cryolite Na 3 AlF 6 and fluorspar (fluorite) CaF 2 .

F always has an oxidation state of -1. The small covalent radius allows it to form compounds with high coordination numbers; For example, SF 6 exists, a S.J. 6 cannot be formed. Metal fluoride ions have a small ion size F- causes high lattice enthalpies and thermodynamic stability.

Due to the high oxidizing ability of fluorine, halogens can react with each other, forming interhalogen compounds(“interhalides”) ClF, ClF 3, BrF 5, IF 7, in which the oxidation state of other halogens varies from +1 to +7.

After incubation of rat liver with NaF absorbed fluorine is concentrated in mitochondria and nuclei of hepatocytes. It is absorbed by bone tissue (teeth, bones, cartilage) 3 times more actively than by blood. F is excreted mainly by the kidneys. The toxic effect of fluoride ions is due to the fact that they bind and thereby inactivate ion-activators of enzyme systems Sa 2+ , Mg 2+ with the formation of poorly soluble fluorides. Complex ions PF - , B.F. 4 - , SiF 6 2-, due to the strength of covalent bonds in their molecules, are biologically inactive. F- inhibits metalloproteins.

Chlorine (Cl) - found in nature mainly in the form of rock salt NaCl. It is obtained from it by electrolysis Cl 2 - heavy yellow-green gas with a pungent odor. In industry, it is used as a bleaching agent and in the production of organochlorine solvents and polymers. In addition, it is widely used for sterilizing water at waterworks in concentrations of (0.6-6)x10 -5 mol/kg. However, when water is polluted with organic nitrogen-containing substances, water chlorination is dangerous because the atoms Cl can replace H atoms in molecules of alkanes and alkenes in photolytic reactions, that is, when irradiated with visible light with a wavelength of 200-800 nm. In this case, toxic organochlorine compounds are formed - derivatives dioxin, in particular, highly toxic 2,3,7,8-tetrachlorodibenzo- n-dioxin (Fig. 1). “Dioxins” generally refer to polychlorinated dibenzo-and-dioxins. All of them, even in very low concentrations, sharply reduce human immunity to viral infections and affect the genetic apparatus.

Rice. 1. Dioxins (2,3,7,8-tetrachlorodibenzo-p-dioxin)

Connections with O 2 (chlorous HClO, “hypochlorite” salts; chloride HClO 2, “chlorite” salts; hypochlorous HClO 3, “chlorate” salts; chlorine HClO 4 acids, perchlorate salts, as well as their anions and oxides) are oxidizing agents; they are used as disinfectants.

The chlorine content in the tissues of mammals is close to its content in sea water. Chloride ions Cl- are almost evenly distributed in the body of living beings in noticeable quantities (from 70 to 103 mmol/l). They are excreted by the kidneys. Liquid chlorine causes serious burns to the skin, and gaseous chlorine severely irritates the eyes and lungs, forming hydrochloric and hypochlorous acids with tissue fluid. Pneumonia may develop in the lungs.

Bromine (Br) - a thick dark red liquid with a pungent odor and heavy brown vapors. It is the only non-metal that is liquid at room temperature. It is used as a fuel additive, as a combustion inhibitor in fire-resistant materials, in paints and pesticides, and in photography. The biological role is poorly studied, although Br 2 poisonous. Ratio Br/Cl in the blood is approximately 0.01, and Br- found mainly in plasma. Accompanies chlorine in metabolic processes and is excreted in the urine.

Iodine (I) - a hard black shiny non-metal. Easily sublimes. It is used as a disinfectant alcohol solution, in food additives, dyes, catalysts, and in photography. It belongs to the biologically necessary (“essential”) elements and is part of the thyroid hormones. Its deficiency is considered a factor predisposing to the development of thyroid and breast cancer.

I selectively accumulates in the thyroid gland (more than 80%). Iodide I- entered the body, quickly concentrates in the gland, where its concentration is 25-500 times higher than in the blood. In the thyroid gland, iodide is oxidized to iodine, which, under the influence of a specific enzyme, iodizes the aromatic rings of tyrosine in thyroglobulin molecules to form lipophilic growth hormones - thyroxine, iodothyronine, triiodothyronine. Iodine in a concentration of 5x10 -5 M uncouples oxidative phosphorylation in mitochondria and easily forms insoluble chelates with doubly charged metal ions, especially with Mg 2+ and MP 2+. The activity of the thyroid gland is activated by iodine-containing thyroid-stimulating hormone of the pituitary gland.

Lack of iodine in human food leads to hypothyroidism and Graves' disease (goiter). Iodine is found in noticeable quantities in seaweed (brown algae of the genus Laminaria) in the form of mono- and diiodotyrosine, as well as mono- and diiodothyronine, which allows these algae to be used for thyroid diseases as a natural source of ready-made growth hormone precursors.

In biogeochemical provinces with a lack of iodine, its salts are added to table salt, but this does not bring positive results. It has been found that iodine deficiency can be combated much more effectively by adding natural iodine-containing products, in particular seaweed, to food products, for example, bread. Previously, iodine was extracted from the ash of brown algae, now - from oil and salt sources. Note that one of the key enzymes of iodine metabolism ( deiodinase), which ensures thyroxine homeostasis, belongs to selenoproteins. Consequently, the fight against iodine deficiency against the background of deficiency Se is meaningless, and taking into account the feedback mechanism, harmful.

Astatine (At) - a radioactive non-metal obtained by neutron bombardment of the isotope 209 Bi. Due to its short half-life, it has not been studied much.

Medical bioinorganics. G.K. Barashkov

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Introduction

Group VII of the periodic table of elements includes manganese, technetium, rhenium, bohrium, as well as, according to the old nomenclature, fluorine, chlorine, bromine, iodine, astatine - which are halogens.

Group 7 elements have 7 valence electrons. All of them are silvery-white refractory metals. In the series Mn - Tc - Re, chemical activity decreases. The electrical conductivity of rhenium is approximately 4 times less than that of tungsten. In air, compact metallic manganese is covered with a thin film of oxide, which protects it from further oxidation even when heated. On the contrary, in a finely crushed state it oxidizes quite easily.

At the outer energy level, halogens have 7 electrons and are strong oxidizing agents. When interacting with metals, an ionic bond occurs and salts are formed. When interacting with more electronegative elements, halogens (except fluorine) can also exhibit reducing properties up to the highest oxidation state of +7.

Technetium and bohrium are radioactive with a fairly short half-life, which is why they do not occur in nature. Manganese is one of the common elements, making up 0.03% of the total number of atoms in the earth's crust.

As for halogens, they are highly reactive and therefore are usually found in nature in the form of compounds. Their prevalence in the earth's crust decreases with increasing atomic radius from fluorine to iodine.

halogen element astatine manganese

1. Seventh1st group of the periodic table

1.1 Glava undergroup seven group. Halogens

The main subgroup of group VII includes the elements fluorine, chlorine, bromine, iodine, and astatine.

Halogens (from the Greek ?lt - salt and gEnpt - birth, origin; sometimes the outdated name halogens is used) - chemical elements of group VII of the periodic table of chemical elements of D. I. Mendeleev

They react with almost all simple substances, except some non-metals. All halogens are energetic oxidizing agents and therefore are found in nature only in the form of compounds. As the atomic number increases, the chemical activity of halogens decreases; the chemical activity of halide ions F ? ,Cl? ,Br? ,I? ,At? decreases.

All halogens are non-metals. At the outer energy level, 7 electrons are strong oxidizing agents. When interacting with metals, an ionic bond occurs and salts are formed. When interacting with more electronegative elements, halogens (except fluorine) can also exhibit reducing properties up to the highest oxidation state of +7.

As mentioned above, halogens have high reactivity, therefore they are usually found in nature in the form of compounds.

Their prevalence in the earth's crust decreases with increasing atomic radius from fluorine to iodine. The amount of astatine in the earth's crust is measured in grams, and ununseptium is absent in nature. Fluorine, chlorine, bromine and iodine are produced on an industrial scale, with production volumes of chlorine being significantly higher than the other three stable halogens.

In nature, these elements occur primarily as halides (with the exception of iodine, which also occurs as sodium or potassium iodate in alkali metal nitrate deposits). Since many chlorides, bromides and iodides are soluble in water, these anions are present in the ocean and natural brines. The main source of fluorine is calcium fluoride, which is very slightly soluble and is found in sedimentary rocks (as fluorite CaF 2).

The main way to obtain simple substances is the oxidation of halides. High positive standard electrode potentials E o (F 2 /F ?) = +2.87 V and E o (Cl 2 /Cl ?) = +1.36 V show that oxidize F ions? and Cl? only possible with strong oxidizing agents. In industry, only electrolytic oxidation is used. When producing fluorine, an aqueous solution cannot be used, since water oxidizes at a much lower potential (+1.32 V) and the resulting fluorine would quickly react with water. Fluorine was first obtained in 1886 by the French chemist Henri Moissan by electrolysis of a solution of potassium hydrofluoride KHF 2 in anhydrous hydrofluoric acid.

In industry, chlorine is mainly produced by electrolysis of an aqueous solution of sodium chloride in special electrolysers. In this case, the following reactions occur:

half-reaction at the anode:

half-reaction at the cathode:

Oxidation of water at the anode is suppressed by using an electrode material that has a higher overvoltage with respect to O 2 than to Cl 2 (such a material is, in particular, RuO 2).

In modern electrolysers, the cathode and anode spaces are separated by a polymer ion-exchange membrane. The membrane allows Na + cations to move from the anode space to the cathode space. The transition of cations maintains electrical neutrality in both parts of the electrolyzer, since during electrolysis negative ions are removed from the anode (conversion of 2Cl ? into Cl 2) and accumulate at the cathode (formation of OH ?). Moving OH ? in the opposite direction could also maintain electrical neutrality, but the OH ion? would react with Cl 2 and negate the whole result.

Bromine is obtained by chemical oxidation of the bromide ion found in seawater. A similar process is used to obtain iodine from natural brines rich in I? . In both cases, chlorine, which has stronger oxidizing properties, is used as an oxidizing agent, and the resulting Br 2 and I 2 are removed from the solution by a stream of air.

Table 1, Some propertieshalogens.

1.2 Fluorine

Fluorine(lat. Fluorum), F, chemical element of group VII of the periodic system of Mendeleev, belongs to the halogens, atomic number 9, atomic mass 18.998403; under normal conditions (0 °C; 0.1 Mn/m2, or 1 kgf/cm2) - a pale yellow gas with a pungent odor.

Natural Fluorine consists of one stable isotope 19 F. A number of isotopes have been artificially obtained, in particular: 16 F with a half-life T ½< 1 сек, 17 F (T Ѕ = 70 сек) , 18 F (T Ѕ = 111 мин) , 20 F (T Ѕ = 11,4 сек) , 21 F (T Ѕ = 5 сек).

Historical reference. The first fluorine compound - fluorite (fluorspar) CaF 2 - was described at the end of the 15th century under the name "fluor" (from the Latin fluo - flow, due to the property of CaF 2 to make viscous slags of metallurgical production liquid-flowing). In 1771, K. Scheele obtained hydrofluoric acid. Free fluorine was isolated by A. Moissan in 1886 by electrolysis of liquid anhydrous hydrogen fluoride containing an admixture of acid potassium fluoride KHF 2.

Fluorine chemistry began to develop in the 1930s, especially quickly during and after the Second World War of 1939-45 in connection with the needs of the nuclear industry and rocket technology. The name "Fluorine" (from the Greek phthoros - destruction, death), proposed by A. Ampere in 1810, is used only in Russian; In many countries the name "fluor" is accepted.

Distribution of fluorine in nature. The average fluorine content in the earth's crust (clarke) is 6.25·10 -2% by mass; in acidic igneous rocks (granites) it is 8·10 -2%, in basic rocks - 3.7·10 -2%, in ultrabasic rocks - 1·10 -2%. Fluorine is present in volcanic gases and thermal waters. The most important fluorine compounds are fluorite, cryolite and topaz. In total, more than 80 fluorine-containing minerals are known. Fluorine compounds are also found in apatites, phosphorites and others. Fluorine is an important biogenic element. In the history of the Earth, the source of fluorine entering the biosphere was the products of volcanic eruptions (gases, etc.).

Physical properties of Fluorine. Gaseous Fluorine has a density of 1.693 g/l (0°C and 0.1 Mn/m2, or 1 kgf/cm2), liquid - 1.5127 g/cm3 (at boiling point); t pl -219.61 °C; boiling point -188.13 °C. The Fluorine molecule consists of two atoms (F 2); at 1000 °C 50% of the molecules dissociate, the dissociation energy is about 155 kJ/mol (37 kcal/mol). Fluorine is poorly soluble in liquid hydrogen fluoride; solubility 2.5·10 -3 g in 100 g HF at -70 °C and 0.4·10 -3 g at -20 °C; in liquid form, unlimitedly soluble in liquid oxygen and ozone.

Chemical properties of Fluorine. The configuration of the outer electrons of the Fluorine atom is 2s 2 2p 5. In compounds it exhibits an oxidation state of -1. The covalent radius of the atom is 0.72E, the ionic radius is 1.3ЗЕ. Electron affinity 3.62 eV, ionization energy (F > F+) 17.418 eV. High values ​​of electron affinity and ionization energy explain the strong electronegativity of the Fluorine atom, the largest among all other elements. The high reactivity of Fluorine determines the exothermic nature of fluorination, which, in turn, is determined by the anomalously low value of the dissociation energy of the Fluorine molecule and the large values ​​of the bond energy of the Fluorine atom with other atoms. Direct fluoridation has a chain mechanism and can easily lead to combustion and explosion. Fluorine reacts with all elements except helium, neon and argon. It interacts with oxygen in a glow discharge, forming oxygen fluorides O 2 F 2, O 3 F 2 and others at low temperatures. Reactions of fluorine with other halogens are exothermic, resulting in the formation of interhalogen compounds. Chlorine interacts with Fluorine when heated to 200-250 "C, giving chlorine monofluoride ClF and chlorine trifluoride ClF 3. ClF 5 is also known, obtained by fluoridation of ClF 3 at high temperature and pressure 25 Mn/m2 (250 kgf/cm2). Bromine and iodine ignites in a fluorine atmosphere at ordinary temperatures, and BrF 3, BrF 5, IF 3, IF 2 can be obtained. Fluorine reacts directly with krypton, xenon and radon, forming the corresponding fluorides (for example, XeF 4, XeF 6, KrF 2 Xenon oxyfluorides are also known.

The interaction of fluorine with sulfur is accompanied by the release of heat and leads to the formation of numerous sulfur fluorides. Selenium and tellurium form higher fluorides SeF 6 and TeF 6 . Fluorine and hydrogen react with combustion; this produces hydrogen fluoride. This is a radical chain branching reaction:

HF* + H 2 = HF + H 2 *; H 2 * + F 2 = HF + H + F

(where HF* and H 2 * are molecules in a vibrationally excited state); the reaction is used in chemical lasers. Fluorine reacts with nitrogen only in an electrical discharge. Charcoal, when interacting with fluorine, ignites at ordinary temperatures; graphite reacts with it under strong heating, and the formation of solid graphite fluoride (CF) X or gaseous perfluorocarbons CF 4, C 2 F 6 and others is possible. Fluorine reacts with boron, silicon, phosphorus, and arsenic in the cold, forming the corresponding fluorides.

Fluorine combines vigorously with most metals; alkali and alkaline earth metals ignite in an atmosphere of fluorine in the cold, Bi, Sn, Ti, Mo, W - with slight heating. Hg, Pb, U, V react with Fluorine at room temperature, Pt - at a dark red heat temperature. When metals react with fluorine, higher fluorides are usually formed, for example UF 6, MoF 6, HgF 2. Some metals (Fe, Cu, Al, Ni, Mg, Zn) react with Fluorine to form a protective film of fluorides, preventing further reaction.

When fluorine interacts with metal oxides in the cold, metal fluorides and oxygen are formed; The formation of metal oxyfluorides (for example, MoO 2 F 2) is also possible. Non-metal oxides either add Fluorine, for example SO 2 + F 2 = SO 2 F 2, or the oxygen in them is replaced by Fluorine, for example SiO 2 + 2F 2 = SiF 4 + O 2. Glass reacts very slowly with Fluorine; in the presence of water the reaction proceeds quickly. Water interacts with Fluorine: 2H 2 O + 2F 2 = 4HF + O 2; in this case, OF 2 and hydrogen peroxide H 2 O 2 are also formed. Nitrogen oxides NO and NO 2 easily add fluorine to form nitrosyl fluoride FNO and nitrile fluoride FNO 2 , respectively. Carbon monoxide adds fluorine when heated to form carbonyl fluoride:

CO + F 2 = COF 2.

Metal hydroxides react with Fluorine to form metal fluoride and oxygen, e.g.

2Ba(OH) 2 + 2F 2 = 2BaF 2 + 2H 2 O + O 2.

Aqueous solutions of NaOH and KOH react with Fluorine at 0°C to form OF 2 .

Metal or nonmetal halides react with fluorine in the cold, with fluorine replacing all halogens.

Sulfides, nitrides and carbides are easily fluorinated. Metal hydrides form metal fluoride and HF with fluorine in the cold; ammonia (in vapor) - N 2 and HF. Fluorine replaces hydrogen in acids or metals in their salts, for example HNO 3 (or NaNO 3) + F 2 = FNO 3 + HF (or NaF); under more severe conditions, Fluorine displaces oxygen from these compounds, forming sulfuryl fluoride, for example

Na 2 SO 4 + 2F 2 = 2NaF + SO 2 F 2 + O 2.

Carbonates of alkali and alkaline earth metals react with fluorine at ordinary temperatures; this produces the corresponding fluoride, CO 2 and O 2 .

Fluorine reacts vigorously with organic substances.

Obtaining Fluorine. The source for the production of fluorine is hydrogen fluoride, which is obtained mainly either by the action of sulfuric acid H 2 SO 4 · on fluorite CaF 2, or by processing apatites and phosphorites. Fluorine production is carried out by electrolysis of the melt of acidic potassium fluoride KF-(1.8-2.0)HF, which is formed when the KF-HF melt is saturated with hydrogen fluoride to a content of 40-41% HF. The material for the electrolyzer is usually steel; electrodes - carbon anode and steel cathode. Electrolysis is carried out at 95-100 °C and a voltage of 9-11 V; Fluorine current output reaches 90-95%. The resulting fluorine contains up to 5% HF, which is removed by freezing followed by absorption with sodium fluoride. Fluorine is stored in a gaseous state (under pressure) and in liquid form (when cooled with liquid nitrogen) in devices made of nickel and alloys based on it (Monel metal), copper, aluminum and its alloys, brass, stainless steel.

Application of Fluorine. Gaseous Fluorine is used for the fluorination of UF 4 into UF 6, used for isotope separation of uranium, as well as for the production of chlorine trifluoride ClF 3 (fluorinating agent), sulfur hexafluoride SF 6 (gaseous insulator in the electrical industry), metal fluorides (for example, W and V ). Liquid Fluorine is an oxidizer for rocket fuels.

Numerous Fluorine compounds are widely used - hydrogen fluoride, aluminum fluoride, silicofluorides, fluorosulfonic acid (solvent, catalyst, reagent for the production of organic compounds containing the group - SO 2 F), BF 3 (catalyst), organofluorine compounds and others.

Safety precautions. Fluorine is toxic, its maximum permissible concentration in the air is approximately 2·10 -4 mg/l, and the maximum permissible concentration with exposure for no more than 1 hour is 1.5·10 -3 mg/l.

Fluoride in the body. Fluorine is constantly included in animal and plant tissues; microelement In the form of inorganic compounds it is found mainly in the bones of animals and humans - 100-300 mg/kg; There is especially a lot of fluoride in teeth. The bones of marine animals are richer in fluorine compared to the bones of land animals. It enters the body of animals and humans mainly with drinking water, the optimal fluorine content in which is 1-1.5 mg/l. With a lack of fluoride, a person develops dental caries, and with an increased intake - fluorosis. High concentrations of fluorine ions are dangerous due to their ability to inhibit a number of enzymatic reactions, as well as to bind biologically important elements. (P, Ca, Mg and others), disrupting their balance in the body. Organic fluoride derivatives are found only in some plants (for example, in the South African Dichapetalum cymosum). The main ones are derivatives of fluoroacetic acid, toxic to both other plants and animals. A connection has been established between fluoride metabolism and the formation of skeletal bone tissue and especially teeth.

Fluorine poisoning is possible among workers in the chemical industry, during the synthesis of fluorine-containing compounds and in the production of phosphate fertilizers. Fluoride irritates the respiratory tract and causes skin burns. In acute poisoning, irritation of the mucous membranes of the larynx and bronchi, eyes, salivation, and nosebleeds occur; in severe cases - pulmonary edema, damage to the central nervous system and others; in chronic cases - conjunctivitis, bronchitis, pneumonia, pneumosclerosis, fluorosis. Skin lesions such as eczema are characteristic. First aid: rinsing the eyes with water, for skin burns - irrigation with 70% alcohol; in case of inhalation poisoning - inhalation of oxygen. Prevention: compliance with safety regulations, wearing special clothing, regular medical examinations, inclusion of calcium and vitamins in the diet.

1.3 Chlorine

Chlorine(lat. Chlorum), Cl, chemical element of group VII of the periodic system of Mendeleev, atomic number 17, atomic mass 35.453; belongs to the halogen family. Under normal conditions (0°C, 0.1 Mn/m2, or 1 kgf/cm2) it is a yellow-green gas with a sharp irritating odor. Natural Chlorine consists of two stable isotopes: 35 Cl (75.77%) and 37 Cl (24.23%). Radioactive isotopes with mass numbers 31-47 have been artificially obtained, in particular: 32, 33, 34, 36, 38, 39, 40 with half-lives (T S) of 0.31, respectively; 2.5; 1.56 sec; 3.1·105 years; 37.3, 55.5 and 1.4 min. 36Cl and 38Cl are used as isotopic tracers.

Historical reference. Chlorine was first obtained in 1774 by K. Scheele by reacting hydrochloric acid with pyrolusite MnO 2 . However, only in 1810 G. Davy established that chlorine is an element and named it chlorine (from the Greek chloros - yellow-green). In 1813, J. L. Gay-Lussac proposed the name Chlorine for this element.

Distribution of Chlorine in nature. Chlorine occurs in nature only in the form of compounds. The average content of Chlorine in the earth's crust (clarke) is 1.7·10 -2% by mass, in acidic igneous rocks - granites and others - 2.4·10 -2, in basic and ultrabasic rocks 5·10 -3. The main role in the history of chlorine in the earth's crust is played by water migration. In the form of Cl ion, it is found in the World Ocean (1.93%), underground brines and salt lakes. The number of its own minerals (mainly natural chlorides) is 97, the main one being halite NaCl (Rock salt). Large deposits of potassium and magnesium chlorides and mixed chlorides are also known: sylvinite KCl, sylvinite (Na,K)Cl, carnalite KCl MgCl 2 6H 2 O, kainite KCl MgSO 4 3H 2 O, bischofite MgCl 2 6H 2 O In the history of the Earth, the supply of HCl contained in volcanic gases to the upper parts of the earth's crust was of great importance.

Physical properties of Chlorine. Chlorine has a boiling point of -34.05°C, a melting point of -101°C. The density of chlorine gas under normal conditions is 3.214 g/l; saturated steam at 0°C 12.21 g/l; liquid Chlorine at a boiling point of 1.557 g/cm3; solid Chlorine at - 102°C 1.9 g/cm 3 . Saturated vapor pressure of Chlorine at 0°C 0.369; at 25°C 0.772; at 100°C 3.814 Mn/m 2 or, respectively, 3.69; 7.72; 38.14 kgf/cm2. Heat of fusion 90.3 kJ/kg (21.5 cal/g); heat of evaporation 288 kJ/kg (68.8 cal/g); The heat capacity of gas at constant pressure is 0.48 kJ/(kg K). Critical constants of Chlorine: temperature 144°C, pressure 7.72 Mn/m2 (77.2 kgf/cm2), density 573 g/l, specific volume 1.745·10 -3 l/g. Solubility (in g/l) of Chlorine at a partial pressure of 0.1 Mn/m2, or 1 kgf/cm2, in water 14.8 (0°C), 5.8 (30°C), 2.8 ( 70°C); in a solution of 300 g/l NaCl 1.42 (30°C), 0.64 (70°C). Below 9.6°C, Chlorine hydrates of variable composition Cl 2 ·nH 2 O (where n = 6-8) are formed in aqueous solutions; These are yellow cubic crystals that decompose with increasing temperature into Chlorine and water. Chlorine is highly soluble in TiCl 4, SiCl 4, SnCl 4 and some organic solvents (especially hexane C 6 H 14 and carbon tetrachloride CCl 4). The Chlorine molecule is diatomic (Cl 2). The degree of thermal dissociation of Cl 2 + 243 kJ = 2Cl at 1000 K is 2.07·10 -4%, at 2500 K 0.909%.

Chemical properties of Chlorine. External electronic configuration of the Cl 3s 2 Sp 5 atom. In accordance with this, Chlorine in compounds exhibits oxidation states of -1, +1, +3, +4, +5, +6 and +7. The covalent radius of the atom is 0.99 E, the ionic radius of Cl is 1.82 E, the electron affinity of the Chlorine atom is 3.65 eV, and the ionization energy is 12.97 eV.

Chemically, Chlorine is very active, directly combines with almost all metals (with some only in the presence of moisture or when heated) and with non-metals (except carbon, nitrogen, oxygen, inert gases), forming the corresponding chlorides, reacts with many compounds, replaces hydrogen in saturated hydrocarbons and joins unsaturated compounds. Chlorine displaces bromine and iodine from their compounds with hydrogen and metals; Of the compounds of chlorine with these elements, it is replaced by fluorine. Alkali metals in the presence of traces of moisture react with Chlorine with ignition; most metals react with dry Chlorine only when heated. Steel, as well as some metals, are resistant in an atmosphere of dry Chlorine at low temperatures, so they are used for the manufacture of equipment and storage facilities for dry Chlorine. Phosphorus ignites in an atmosphere of Chlorine, forming PCl 3, and with further chlorination - PCl 5; sulfur with Chlorine when heated gives S 2 Cl 2, SCl 2 and other S n Cl m. Arsenic, antimony, bismuth, strontium, tellurium interact vigorously with Chlorine. A mixture of chlorine and hydrogen burns with a colorless or yellow-green flame with the formation of hydrogen chloride (this is a chain reaction).

The maximum temperature of the hydrogen-chlorine flame is 2200°C. Mixtures of chlorine with hydrogen containing from 5.8 to 88.5% H 2 are explosive.

With oxygen, Chlorine forms oxides: Cl 2 O, ClO 2, Cl 2 O 6, Cl 2 O 7, Cl 2 O 8, as well as hypochlorites (salts of hypochlorous acid), chlorites, chlorates and perchlorates. All oxygen compounds of chlorine form explosive mixtures with easily oxidized substances. Chlorine oxides are weakly stable and can spontaneously explode; hypochlorites slowly decompose during storage; chlorates and perchlorates can explode under the influence of initiators.

Chlorine in water hydrolyzes, forming hypochlorous and hydrochloric acids: Cl 2 + H 2 O = HClO + HCl. When aqueous solutions of alkalis are chlorinated in the cold, hypochlorites and chlorides are formed: 2NaOH + Cl 2 = NaClO + NaCl + H 2 O, and when heated, chlorates are formed. Chlorination of dry calcium hydroxide produces bleach.

When ammonia reacts with chlorine, nitrogen trichloride is formed. When chlorinating organic compounds, Chlorine either replaces hydrogen or joins multiple bonds, forming various chlorine-containing organic compounds.

Chlorine forms interhalogen compounds with other halogens. Fluorides ClF, ClF 3, ClF 3 are very reactive; for example, in a ClF 3 atmosphere, glass wool spontaneously ignites. Known compounds of chlorine with oxygen and fluorine are Chlorine oxyfluorides: ClO 3 F, ClO 2 F 3, ClOF, ClOF 3 and fluorine perchlorate FClO 4.

Getting Chlorine. Chlorine began to be produced industrially in 1785 by reacting hydrochloric acid with manganese (II) oxide or pyrolusite. In 1867, the English chemist G. Deacon developed a method for producing chlorine by oxidizing HCl with atmospheric oxygen in the presence of a catalyst. Since the late 19th and early 20th centuries, chlorine has been produced by electrolysis of aqueous solutions of alkali metal chlorides. These methods produce 90-95% of Chlorine in the world. Small amounts of Chlorine are obtained by-product in the production of magnesium, calcium, sodium and lithium by electrolysis of molten chlorides. Two main methods of electrolysis of aqueous solutions of NaCl are used: 1) in electrolyzers with a solid cathode and a porous filter diaphragm; 2) in electrolyzers with a mercury cathode. In both methods, Chlorine gas is released on a graphite or oxide titanium-ruthenium anode. According to the first method, hydrogen is released at the cathode and a solution of NaOH and NaCl is formed, from which commercial caustic soda is separated by subsequent processing. According to the second method, sodium amalgam is formed at the cathode; when it is decomposed with pure water in a separate apparatus, a NaOH solution, hydrogen and pure mercury are obtained, which again goes into production. Both methods give 1.125 t of NaOH per 1 ton of Chlorine.

Electrolysis with a diaphragm requires less capital investment to organize the production of Chlorine and produces cheaper NaOH. The mercury cathode method produces very pure NaOH, but the loss of mercury pollutes the environment.

Use of Chlorine. One of the important branches of the chemical industry is the chlorine industry. The main quantities of Chlorine are processed at the site of its production into chlorine-containing compounds. Chlorine is stored and transported in liquid form in cylinders, barrels, railway tanks or in specially equipped vessels. Industrial countries are characterized by the following approximate consumption of Chlorine: for the production of chlorine-containing organic compounds - 60-75%; inorganic compounds containing Chlorine, -10-20%; for bleaching pulp and fabrics - 5-15%; for sanitary needs and water chlorination - 2-6% of total production.

Chlorine is also used to chlorinate some ores to extract titanium, niobium, zirconium and others.

Chlorine in the body. Chlorine is one of the biogenic elements, a constant component of plant and animal tissues. The Chlorine content in plants (a lot of Chlorine in halophytes) ranges from thousandths of a percent to whole percent, in animals - tenths and hundredths of a percent. The daily requirement of an adult for Chlorine (2-4 g) is covered by food products. Chlorine is usually supplied in excess with food in the form of sodium chloride and potassium chloride. Bread, meat and dairy products are especially rich in Chlorine. In the animal body, Chlorine is the main osmotically active substance in blood plasma, lymph, cerebrospinal fluid and some tissues. Plays a role in water-salt metabolism, promoting tissue retention of water. Regulation of acid-base balance in tissues is carried out along with other processes by changing the distribution of Chlorine between the blood and other tissues. Chlorine is involved in energy metabolism in plants, activating both oxidative phosphorylation and photophosphorylation. Chlorine has a positive effect on the absorption of oxygen by roots. Chlorine is necessary for the production of oxygen during photosynthesis by isolated chloroplasts. Most nutrient media for artificial plant cultivation do not contain chlorine. It is possible that very low concentrations of Chlorine are sufficient for plant development.

Chlorine poisoning is possible in the chemical, pulp and paper, textile, pharmaceutical industries and others. Chlorine irritates the mucous membranes of the eyes and respiratory tract. Primary inflammatory changes are usually accompanied by a secondary infection. Acute poisoning develops almost immediately. When inhaling medium and low concentrations of Chlorine, tightness and pain in the chest, dry cough, rapid breathing, pain in the eyes, lacrimation, increased levels of leukocytes in the blood, body temperature, etc. are observed. Bronchopneumonia, toxic pulmonary edema, depression, convulsions are possible . In mild cases, recovery occurs within 3-7 days. As long-term consequences, catarrh of the upper respiratory tract, recurrent bronchitis, pneumosclerosis and others are observed; possible activation of pulmonary tuberculosis. With prolonged inhalation of small concentrations of Chlorine, similar but slowly developing forms of the disease are observed. Prevention of poisoning: sealing production facilities, equipment, effective ventilation, using a gas mask if necessary. The production of chlorine, bleach and other chlorine-containing compounds is classified as production with hazardous working conditions.

1.4 Bromine

Bromine(lat. Bromum), Br, a chemical element of group VII of the periodic system of Mendeleev, belongs to the halogens; atomic number 35, atomic mass 79.904; red-brown liquid with a strong unpleasant odor. Bromine was discovered in 1826 by the French chemist A. J. Balard while studying the brines of the Mediterranean salt fields; named from Greek. bromos - stench. Natural Bromine consists of 2 stable isotopes 79 Br (50.54%) and 81 Br (49.46%). Of the artificially obtained radioactive isotopes, Bromine is the most interesting 80 Br, on the example of which I. V. Kurchatov discovered the phenomenon of isomerism of atomic nuclei.

Distribution of Bromine in nature. The content of Bromine in the earth's crust (1.6·10 -4% by mass) is estimated at 10 15 -10 16 tons. Bromine is found mostly in a dispersed state in igneous rocks, as well as in widespread halides. Bromine is a constant companion of chlorine. Bromide salts (NaBr, KBr, MgBr 2) are found in deposits of chloride salts (in table salt up to 0.03% Br, in potassium salts - sylvite and carnallite - up to 0.3% Br), as well as in sea water (0.065% Br), salt lake brines (up to 0.2% Br) and underground brines commonly associated with salt and oil deposits (up to 0.1% Br). Due to their good solubility in water, bromide salts accumulate in residual brines of sea and lake water bodies. Bromine migrates in the form of easily soluble compounds, very rarely forming solid mineral forms represented by bromyrite AgBr, embolite Ag (Cl, Br) and iodembolite Ag (Cl, Br, I). The formation of minerals occurs in oxidation zones of sulfide silver deposits that form in arid desert areas.

Physical properties of Bromine. At -7.2°C, liquid Bromine solidifies, turning into red-brown needle-shaped crystals with a faint metallic luster. Bromine vapor is yellow-brown in color, boiling point 58.78°C. The density of liquid Bromine (at 20°C) is 3.1 g/cm 3 . Bromine is soluble in water to a limited extent, but better than other halogens (3.58 g of Bromine in 100 g of H 2 O at 20 ° C). Below 5.84°C, garnet-red crystals of Br 2 8H 2 O precipitate from water. Bromine is especially soluble in many organic solvents, which is used to extract it from aqueous solutions. Bromine in solid, liquid and gaseous states consists of 2-atomic molecules. Noticeable dissociation into atoms begins at a temperature of about 800°C; dissociation is also observed under the influence of light.

Chemical properties of Bromine. The configuration of the outer electrons of the Bromine atom is 4s 2 4p 5. The valence of Bromine in compounds is variable, the oxidation state is -1 (in bromides, for example KBr), +1 (in hypobromites, NaBrO), +3 (in bromites, NaBrO 2), +5 (in bromates, KBrOz) and +7 ( in perbromates, NaBrO 4). Chemically, Bromine is very active, occupying a place in reactivity between chlorine and iodine. The interaction of Bromine with sulfur, selenium, tellurium, phosphorus, arsenic and antimony is accompanied by strong heating, sometimes even the appearance of a flame. Bromine also reacts vigorously with some metals, such as potassium and aluminum. However, many metals react with anhydrous Bromine with difficulty due to the formation of a protective film of bromide, which is insoluble in Bromine, on their surface. Of the metals, the most resistant to the action of Bromine, even at elevated temperatures and in the presence of moisture, are silver, lead, platinum and tantalum (gold, unlike platinum, reacts vigorously with Bromine). Bromine does not directly combine with oxygen, nitrogen and carbon, even at elevated temperatures. Bromine compounds with these elements are obtained indirectly. These are the extremely fragile oxides Br 2 O, Br O 2 and Br 3 O 8 (the latter is obtained, for example, by the action of ozone on Bromine at 80°C). Bromine reacts directly with halogens, forming BrF 3, BrF 5, BrCl, IBr and others.

Bromine is a strong oxidizing agent. Thus, it oxidizes sulfites and thiosulfates in aqueous solutions to sulfates, nitrites to nitrates, ammonia to free nitrogen (3Br 2 + 8NH 3 = N 2 + NH 4 Br). Bromine displaces iodine from its compounds, but is itself displaced by chlorine and fluorine. Free Bromine is released from aqueous solutions of bromides also under the influence of strong oxidizing agents (KMnO 4, K 2 Cr 2 O 7) in an acidic environment. When dissolved in water, Bromine partially reacts with it (Br 2 + H 2 O = HBr + HBrO) to form hydrobromic acid HBr and unstable hypobromous acid HBrO. A solution of Bromine in water is called bromine water. When Bromine dissolves in alkali solutions in the cold, bromide and hypobromite are formed (2NaOH + Br 2 = NaBr + NaBrO + H 2 O), and at elevated temperatures (about 100 ° C) - bromide and bromate (6NaOH + 3Br 2 = 5NaBr + NaBrO 3 + 3H 2 O). Of the reactions of Bromine with organic compounds, the most typical are addition at the C=C double bond, as well as substitution of hydrogen (usually under the action of catalysts or light).

Obtaining Bromine. The starting materials for the production of Bromine are sea water, lake and underground brines and potassium production liquors containing Bromine in the form of bromide ion Br - (from 65 g/m 3 in sea water to 3-4 kg/m 3 and higher in potassium liquors production). Bromine is isolated with the help of chlorine (2Br - + Cl 2 = Br 2 + 2Cl -) and distilled from the solution with water vapor or air. Steam stripping is carried out in columns made of granite, ceramics or other material resistant to bromine. Heated brine is fed into the column from above, and chlorine and water vapor are supplied from below. Bromine vapor leaving the column is condensed in ceramic refrigerators. Next, Bromine is separated from water and purified from chlorine impurities by distillation. Air stripping makes it possible to use brines with a low content of bromine to obtain bromine; it is unprofitable to extract bromine from it by steam as a result of high steam consumption. Bromine is removed from the resulting bromine-air mixture using chemical absorbents. For this, solutions of iron bromide (2FeBr 2 + Br 2 = 2FeBr 3) are used, which, in turn, is obtained by reducing FeBr 3 with iron filings, as well as solutions of sodium hydroxides or carbonates or gaseous sulfur dioxide reacting with Bromine in the presence of water vapor with the formation of hydrobromic and sulfuric acids (Br 2 + SO 2 + 2H 2 O = 2HBr + H 2 SO 4). Bromine is isolated from the resulting intermediates by the action of chlorine (from FeBr 3 and HBr) or acid (5NaBr + NaBrO 3 + 3 H 2 SO 4 = 3Br 2 + 3Na 2 SO 4 + 3H 2 O). If necessary, intermediate products are processed into bromide compounds without releasing elemental Bromine.

Inhalation of Bromine vapor when its content in the air is 1 mg/m3 or more causes cough, runny nose, nosebleeds, dizziness, headache; at higher concentrations - suffocation, bronchitis, and sometimes death. The maximum permissible concentration of Bromine vapor in the air is 2 mg/m3. Liquid Bromine acts on the skin, causing poorly healing burns. Work with bromine should be carried out in fume hoods. In case of poisoning with Bromine vapor, it is recommended to inhale ammonia, using for this purpose a highly diluted solution of it in water or ethyl alcohol. Sore throat caused by inhaling Bromine vapor is relieved by ingesting hot milk. Bromine that gets on the skin is washed off with plenty of water or blown off with a strong stream of air. Burnt areas are lubricated with lanolin.

Application of Bromine. Bromine is used quite widely. It is the starting product for the production of a number of bromide salts and organic derivatives. Large quantities of Bromine are used to produce ethyl bromide and dibromoethane - components of the ethyl liquid added to gasoline to increase their detonation resistance. Bromine compounds are used in photography, in the production of a number of dyes, methyl bromide and some other Bromine compounds are used as insecticides. Some organic bromine compounds serve as effective fire extinguishing agents. Bromine and bromine water are used in chemical analyzes to determine many substances. In medicine, sodium, potassium, ammonium bromides are used, as well as organic bromine compounds, which are used for neuroses, hysteria, increased irritability, insomnia, hypertension, epilepsy and chorea.

Bromine in the body. Bromine is a constant constituent of animal and plant tissues. Terrestrial plants contain on average 7·10 -4% Bromine in raw matter, animals ~1·10 -4%. Bromine is found in various secretions (tears, saliva, sweat, milk, bile). In the blood of a healthy person, the bromine content ranges from 0.11 to 2.00 mg%. Using radioactive Bromine (82 Br), its selective absorption by the thyroid gland, the medulla of the kidneys and the pituitary gland was established. Bromides introduced into the body of animals and humans increase the concentration of inhibitory processes in the cerebral cortex and help normalize the state of the nervous system, which has suffered from overstrain of the inhibitory process. At the same time, lingering in the thyroid gland, Bromine enters into a competitive relationship with iodine, which affects the activity of the gland, and in connection with this, the state of metabolism.

1.5 Iodine

Iodine(lat. Iodum), I, a chemical element of group VII of the periodic system of Mendeleev, belongs to the halogens (the outdated name Iodine and the symbol J are also found in the literature); atomic number 53, atomic mass 126.9045; crystals of black-gray color with a metallic sheen. Natural iodine consists of one stable isotope with a mass number of 127. Iodine was discovered in 1811 by the French chemist B. Courtois. By heating the mother brine of seaweed ash with concentrated sulfuric acid, he observed the release of violet vapor (hence the name Iodine - from the Greek iodes, ioides - violet-like in color, violet), which condensed into dark shiny plate-like crystals. In 1813-1814, the French chemist J. L. Gay-Lussac and the English chemist G. Davy proved the elemental nature of iodine.

Distribution of iodine in nature. The average iodine content in the earth's crust is 4·10 -5% by mass. Iodine compounds are scattered in the mantle and magmas and in the rocks formed from them (granites, basalts and others); deep minerals of Iodine are unknown. The history of iodine in the earth's crust is closely related to living matter and biogenic migration. In the biosphere, processes of its concentration are observed, especially by marine organisms (algae, sponges and others). Eight supergene iodine minerals are known to form in the biosphere, but they are very rare. The main reservoir of iodine for the biosphere is the World Ocean (1 liter contains on average 5·10 -5 g of iodine). From the ocean, iodine compounds dissolved in drops of sea water enter the atmosphere and are carried by winds to the continents. (Areas remote from the ocean or fenced off from sea winds by mountains are depleted in iodine.) Iodine is easily adsorbed by organic matter in soils and marine silts. When these silts become compacted and sedimentary rocks form, desorption occurs and some of the iodine compounds pass into groundwater. This is how iodine-bromine waters used for the extraction of iodine are formed, especially characteristic of oil field areas (in some places, 1 liter of these waters contains over 100 mg of iodine).

Physical properties of Iodine. Iodine density is 4.94 g/cm 3, melting point 113.5°C, boiling point 184.35°C. The molecule of liquid and gaseous iodine consists of two atoms (I 2). A noticeable dissociation of I 2 = 2I is observed above 700 °C, as well as under the influence of light. Already at ordinary temperatures, iodine evaporates, forming a sharp-smelling purple vapor. When heated slightly, iodine sublimes, settling in the form of shiny thin plates; this process serves to purify iodine in laboratories and industry. Iodine is poorly soluble in water (0.33 g/l at 25 °C), well soluble in carbon disulfide and organic solvents (benzene, alcohol and others), as well as in aqueous solutions of iodides.

Chemical properties of Iodine. The configuration of the outer electrons of the Iodine atom is 5s 2 5p 5. In accordance with this, iodine exhibits variable valence (oxidation state) in compounds: -1 (in HI, KI), +1 (in HIO, KIO), +3 (in ICl 3), +5 (in HIO 3, KIO 3 ) and +7 (in HIO 4, KIO 4). Chemically, iodine is quite active, although to a lesser extent than chlorine and bromine. Iodine reacts vigorously with metals when slightly heated, forming iodides (Hg + I 2 = HgI 2). Iodine reacts with hydrogen only when heated and not completely, forming hydrogen iodide. Iodine does not combine directly with carbon, nitrogen, or oxygen. Elemental Iodine is an oxidizing agent, less powerful than chlorine and bromine. Hydrogen sulfide H 2 S, sodium thiosulfate Na 2 S 2 O 3 and other reducing agents reduce it to I - (I 2 + H 2 S = S + 2HI). Chlorine and other strong oxidizing agents in aqueous solutions convert it into IO 3 - (5Cl 2 + I 2 + 6H 2 O = 2HIO 3 H + 10HCl). When dissolved in water, iodine partially reacts with it (I 2 + H 2 O = HI + HIO); in hot aqueous solutions of alkalis, iodide and iodate are formed (3I 2 + 6NaOH = 5NaI + NaIO 3 + 3H 2 O). When adsorbed on starch, iodine turns it dark blue; it is used in iodometry and qualitative analysis for the detection of Iodine.

Iodine vapors are poisonous and irritate mucous membranes. Iodine has a cauterizing and disinfecting effect on the skin. Iodine stains are washed off with solutions of soda or sodium thiosulfate.

Obtaining Iodine. The raw material for the industrial production of iodine is oil drilling water; seaweed, as well as mother solutions of Chilean (sodium) nitrate containing up to 0.4% Iodine in the form of sodium iodate. To extract iodine from oil waters (usually containing 20-40 mg/l Iodine in the form of iodides), they are first treated with chlorine (2 NaI + Cl 2 = 2NaCl + I 2) or nitrous acid (2NaI + 2NaNO 2 + 2H 2 SO 4 = 2Na 2 SO 4 + 2NO + I 2 + 2H 2 O). The released iodine is either adsorbed by active carbon or blown out with air. Iodine adsorbed by coal is treated with caustic alkali or sodium sulfite (I 2 + Na 2 SO 3 + H 2 O = Na 2 SO 4 + 2HI). Free Iodine is isolated from the reaction products by the action of chlorine or sulfuric acid and an oxidizing agent, for example, potassium dichromate (K 2 Cr 2 O 7 + 7H 2 SO 4 + 6NaI = K 2 SO 4 + 3Na 2 SO 4 + Cr 2 (SO 4)S + 3I 2). When blown out with air, iodine is absorbed by a mixture of sulfur oxide (IV) with water vapor (2H 2 O + SO 2 + I 2 = H 2 SO 4 + 2HI) and then Iodine is replaced with chlorine (2HI + Cl 2 = 2HCl + I 2). Crude crystalline iodine is purified by sublimation.

Application of Iodine. Iodine and its compounds are used mainly in medicine and analytical chemistry, as well as in organic synthesis and photography.

Iodine in the body. Iodine is a microelement essential for animals and humans. In soils and plants of taiga-forest non-chernozem, dry steppe, desert and mountain biogeochemical zones, iodine is contained in insufficient quantities or is not balanced with some other microelements (Co, Mn, Cu); This is associated with the spread of endemic goiter in these areas. The average iodine content in soils is about 3·10 -4%, in plants about 2·10 -5%. There is little iodine in surface drinking waters (from 10 -7 to 10 -9%). In coastal areas, the amount of iodine in 1 m 3 of air can reach 50 mcg, in continental and mountainous areas it is 1 or even 0.2 mcg.

The absorption of iodine by plants depends on the content of its compounds in the soil and on the type of plant. Some organisms (so-called iodine concentrators), for example, seaweed - fucus, kelp, phyllophora, accumulate up to 1% Iodine, some sponges - up to 8.5% (in the skeletal substance spongin). Algae that concentrate iodine are used for its industrial production. Iodine enters the animal body with food, water, and air. The main source of iodine is plant products and feed. Iodine absorption occurs in the anterior sections of the small intestine. The human body accumulates from 20 to 50 mg of iodine, including about 10-25 mg in the muscles, and 6-15 mg in the thyroid gland. Using radioactive iodine (131 I and 125 I), it was shown that in the thyroid gland Iodine accumulates in the mitochondria of epithelial cells and is part of the diiodo- and monoiodotyrosines formed in them, which condense into the hormone tetraiodothyronine (thyroxine). Iodine is excreted from the body mainly through the kidneys (up to 70-80%), mammary, salivary and sweat glands, partly with bile.

In different biogeochemical provinces, the iodine content in the daily diet varies (for humans from 20 to 240 mcg, for sheep from 20 to 400 mcg). An animal's need for iodine depends on its physiological state, time of year, temperature, and the body's adaptation to the iodine content in the environment. The daily need for Iodine in humans and animals is about 3 mcg per 1 kg of body weight (increases during pregnancy, increased growth, and cooling). The introduction of Iodine into the body increases basal metabolism, enhances oxidative processes, tones muscles, and stimulates sexual function.

Due to a greater or lesser deficiency of Iodine in food and water, iodization of table salt is used, usually containing 10-25 g of potassium iodide per 1 ton of salt. The use of fertilizers containing iodine can double or triple its content in crops.

Iodine in medicine. Preparations containing iodine have antibacterial and antifungal properties, they also have an anti-inflammatory and distracting effect; They are used externally to disinfect wounds and prepare the surgical field. When taken orally, Iodine preparations affect metabolism and enhance thyroid function. Small doses of Iodine (microiodine) inhibit the function of the thyroid gland, affecting the formation of thyroid-stimulating hormone in the anterior pituitary gland. Since iodine affects protein and fat (lipid) metabolism, it has found application in the treatment of atherosclerosis, as it reduces cholesterol in the blood; also increases the fibrinolytic activity of the blood. For diagnostic purposes, radiopaque agents containing iodine are used.

With prolonged use of Iodine preparations and with increased sensitivity to them, iodism may appear - runny nose, urticaria, Quincke's edema, salivation and lacrimation, acne-like rash (iododerma), etc. Iodine preparations should not be taken in case of pulmonary tuberculosis, pregnancy, kidney disease, chronic pyoderma, hemorrhagic diathesis, urticaria.

Iodine is radioactive. Artificially radioactive isotopes of Iodine - 125 I, 131 I, 132 I and others are widely used in biology and especially in medicine to determine the functional state of the thyroid gland and treat a number of its diseases. The use of radioactive iodine in diagnostics is associated with the ability of iodine to selectively accumulate in the thyroid gland; use for medicinal purposes is based on the ability of beta-radiation of iodine radioisotopes to destroy the secretory cells of the gland. When the environment is contaminated with nuclear fission products, radioactive isotopes of iodine are quickly included in the biological cycle, ultimately ending up in milk and, consequently, in the human body. Their penetration into the body of children, whose thyroid gland is 10 times smaller than that of adults and also has greater radiosensitivity, is especially dangerous. In order to reduce the deposition of radioactive isotopes of iodine in the thyroid gland, it is recommended to use stable iodine preparations (100-200 mg per dose). Radioactive iodine is quickly and completely absorbed from the gastrointestinal tract and selectively deposited in the thyroid gland. Its absorption depends on the functional state of the gland. Relatively high concentrations of radioisotopes of Iodine are also found in the salivary and mammary glands and the mucous membrane of the gastrointestinal tract. Radioactive iodine not absorbed by the thyroid gland is almost completely and relatively quickly excreted in the urine.

Similar documents

Study of the concept and basic properties of halogens - chemical elements (fluorine, chlorine, bromine, iodine and astatine), constituting the main subgroup of group VII of the periodic table D.I. Mendeleev. Positive and negative effects of halogens on the human body.

presentation, added 10/20/2011


History of discovery and place in the periodic table of chemical elements D.I. Mendeleev halogens: fluorine, chlorine, bromine, iodine and astatine. Chemical and physical properties of elements, their applications. The prevalence of elements and the production of simple substances.

presentation, added 03/13/2014


Chemical elements related to halogens: fluorine, chlorine, bromine, iodine and astatine. Chemical characteristics, atomic numbers of elements, their physical properties, ionization energy and electronegativity. Oxidation states of halogens, dissociation energy.

presentation, added 12/16/2013


The concept and practical significance of halogens, their physical and chemical properties, distinctive features. Characteristics and methods of producing halogens: iodine, bromine, chlorine, fluorine, astatine. Reactions characteristic of these halogens, areas of their use.

presentation, added 03/11/2011


General characteristics of chemical elements of group IV of the periodic table, their occurrence in nature and compounds with other non-metals. Preparation of germanium, tin and lead. Physico-chemical properties of metals of the titanium subgroup. Areas of application of zirconium.

presentation, added 04/23/2014


Halogens are chemical elements belonging to the main subgroup of group VII of the periodic system of Mendeleev. The halogens include fluorine, chlorine, bromine, iodine and astatine. All halogens are energetic oxidizing agents and therefore are found in nature only in the form of compounds.

abstract, added 03/20/2009


Chlorine is the 17th element of the periodic table of chemical elements of the third period, with atomic number 17. A chemically active non-metal, part of the halogen group. Physical properties of chlorine, interaction with metals and non-metals, oxidative reactions.

presentation, added 12/26/2011


Properties of elements of the nitrogen subgroup, structure and characteristics of atoms. An increase in metallic properties when moving elements from top to bottom in the periodic table. Distribution of nitrogen, phosphorus, arsenic, antimony and bismuth in nature, their application.

abstract, added 06/15/2009


Physical and chemical properties of halogens, their position in Mendeleev’s Periodic Table of Elements. The main sources and biological significance of chlorine, bromine, iodine, fluorine. Finding halogens in nature, their production and industrial use.

presentation, added 12/01/2014


General characteristics of metals. Definition, structure. General physical properties. Methods for obtaining metals. Chemical properties of metals. Metal alloys. Characteristics of the elements of the main subgroups. Characteristics of transition metals.

The elements included in group VII of the periodic table are divided into two subgroups: the main one - the halogen subgroup - and the secondary one - the manganese subgroup. Hydrogen is also placed in this group, although its atom has a single electron on the outer, valence level and should be placed in group I. However, hydrogen has very little in common with both the elements of the main subgroup - the alkali metals, and the elements of the secondary subgroup - copper, silver and gold. At the same time, like halogens, it adds an electron in reactions with active metals and forms hydrides that have some similarities with halides.

The subgroup of halogens includes fluorine, chlorine, bromine, iodine and astatine. The first four elements are found in nature, the last one is obtained artificially and therefore has been studied much less than the other halogens. The word halogen means salt-forming. The elements of the subgroup received this name due to the ease with which they react with many metals, forming salts. All halogens have the structure of the outer electron shell s 2 p 5. Therefore, they easily accept an electron, forming a stable noble gas electron shell (s 2 p 6). Fluorine has the smallest atomic radius in the subgroup; for the rest it increases in the series F< Cl < Br < I < Аt и составляет соответственно 133; 181; 196; 220 и 270 пм. В таком же порядке уменьшается сродство атомов элементов к электрону. Галогены - очень активные элементы. Они могут отнимать, электроны не только у атомов, которые их легко отдают, но и у ионов и даже вытеснять другие галогены, менее активные, из их соединений. Например, фтор вытесняет хлор из хлоридов, хлор - бром из бромидов, а бром - иод из иодидов. Из всех галогенов только фтор, находящийся во II периоде, не имеет незаполненного d-уровня. По этой причине он не может иметь больше одного неспаренного электрона и проявляет валентность только -1. В атомах других галогенов d-уровень не заполнен, что дает им возможность иметь различное количество неспаренных электронов и проявлять валентность -1, +1, +3, +5 и +7, наблюдающуюся в кислородных соединениях хлора, брома и иода К подгруппе марганца принадлежат марганец, технеций и рений. В отличии от галогенов элементы подгруппы марганца имеют на внешнем электронном уровне всего два электрона и поэтому не проявляют способности присоединять электроны, образуя отрицательно заряженные ионы.Марганец распространен в природе и широко используется в промышленности.Технеций радиоактивен, в природе не встречаемся, а получен искусственно (впервые - Э. Сегре и К.Перрье, 1937}. Этот элемент образуется вследствие радиоактивного распада урана. Рений относится к числу рассеянных элементов. Он не образует самостоятельных минералов, а встречается в качестве спутника некоторых минералов, особенно молибденовых. Он был открыт В. и И. Ноддак в 1925 г. Сплавы, имеющие небольшие добавки рения, обладают повышенной устойчивостью против коррозии. Добавка рения к и ее сплавам увеличивает их механическую прочность. Это свойство рения позволяет применять его вместо благородного металла иридия. Платино-платинорениевые термопары работают лучше платино-платиноиридиевых, но их нельзя использовать при очень высоких температурах, так как образуется летучее соединение Re 2 O 7 .

A characteristic feature of nonmetals is a larger (compared to metals) number of electrons in the outer energy level of their atoms. This determines their greater ability to attach additional electrons and exhibit higher oxidative activity than metals. Particularly strong oxidizing properties, i.e. the ability to add electrons, are exhibited by nonmetals located in the 2nd and 3rd periods of groups VI-VII. If we compare the arrangement of electrons in orbitals in the atoms of fluorine, chlorine and other halogens, then we can judge their distinctive properties. The fluorine atom has no free orbitals. Therefore, fluorine atoms can only exhibit valence I and oxidation state 1. The strongest oxidizing agent is fluorine. In the atoms of other halogens, for example in the chlorine atom, there are free d-orbitals at the same energy level. Thanks to this, electron pairing can occur in three different ways. In the first case, chlorine can exhibit an oxidation state of +3 and form chlorous acid HClO2, which corresponds to salts - chlorites, for example potassium chlorite KClO2. In the second case, chlorine can form compounds in which the oxidation state of chlorine is +5. Such compounds include hypochlorous acid HClO3 and its salts - chlorates, for example potassium chlorate KClO3 (Berthollet salt). In the third case, chlorine exhibits an oxidation state of +7, for example in perchloric acid HClO4 and its salts, perchlorates (in potassium perchlorate KClO4).

Particular analytical reactions of Mn 2+ ions

1.5.5. Oxidation with sodium bismuthate NaBiO 3 proceeds according to the equation:

2Mn(NO 3) 2 + 5NaBiO 3 + 16HNO 3 = 2HMnO 4 + 5Bi(NO 3) 3 + 5NaNO 3 + 7H 2 O.

The reaction occurs in the cold. Executing the reaction: add 3-4 drops of a 6 M HNO 3 solution and 5-6 drops of H 2 O to 1-2 drops of a manganese salt solution, after which a little NaBiO 3 powder is added with a spatula. After mixing the contents of the test tube, let it stand for 1-2 minutes, then centrifuge to separate excess sodium bismuthate. In the presence of Mn 2+, the solution turns purple as a result of the formation of manganese acid, which is one of the most powerful oxidizing agents.

1.5.6. Oxidation of PbO 2 with lead dioxide in a nitric acid medium when heated:

2Mn(NO 3) 2 + 5PbO 2 + 6HNO 3 → 2HMnO 4 + 5Pb(NO 3) 2 + 2H 2 O.

Executing the reaction: Take a little PbO 2 powder and place it in a test tube, add 4-5 drops of 6 M HNO 3 there, and heat with stirring. The appearance of a purple color indicates the presence of Mn 2+.

1.5.7. Of importance in the analysis are the reactions of Mn 2+ with alkali metal carbonates, sodium hydrogen phosphate, oxidation reactions with ammonium persulfate, oxidation of benzidine with Mn 4+ compounds, reduction of AgCl to metallic silver with Mn 2+ ions.

88. Elements of group VIII B. Typical properties of the most important compounds. Biological role. Analytical reactions to Fe 3+ and Fe 2+ ions.

Iron subgroup- chemical elements of group 8 of the periodic table of chemical elements (according to the outdated classification - elements of the secondary subgroup of group VIII). The group includes iron Fe, ruthenium Ru and osmium Os. Based on the electronic configuration of the atom, the artificially synthesized element also belongs to the same group Hassiy Hs, which was discovered in 1984 at the Heavy Ion Research Center (German). Gesellschaft für Schwerionenforschung, GSI), Darmstadt, Germany as a result of bombardment of a lead (208 Pb) target with a beam of iron-58 ions from the UNILAC accelerator. As a result of the experiment, 3 265 Hs nuclei were synthesized, which were reliably identified by the parameters of the α-decay chain. Simultaneously and independently, the same reaction was studied at JINR (Dubna, Russia), where, based on the observation of 3 events of α-decay of the 253 Es nucleus, it was also concluded that in this reaction the 265 Hs nucleus, subject to α-decay, was synthesized. All group 8 elements contain 8 electrons in their valence shells. Two elements of the group - ruthenium and osmium - belong to the platinum metal family. As in other groups, members of group 8 elements exhibit patterns of electronic configuration, especially in their outer shells, although, oddly enough, ruthenium does not follow this trend. However, the elements of this group also show similarities in physical properties and chemical behavior: Iron is rarely found in nature in its pure form; it is most often found in iron-nickel meteorites. The prevalence of iron in the earth's crust is 4.65% (4th place after oxygen, silicon and aluminum). Iron is also believed to make up most of the earth's core.

Ruthenium is the only platinum metal found in living organisms. (According to some sources - also platinum). Concentrated mainly in muscle tissue. Higher ruthenium oxide is extremely toxic and, being a strong oxidizing agent, can cause combustion of flammable substances.

Analytical reactions

Potassium hexacyanoferrate(III) K 3 with the Fe 2+ cation forms a blue precipitate of “Turnboole blue”:

3FeSO 4 + 2K 3 → Fe 3 2 ↓+ 3K 2 SO 4 ,

3Fe 2+ + 2Fe(CN) 6 3– → Fe 3 2 ↓.

The precipitate does not dissolve in acids, but decomposes with alkalis to form Fe(OH) 2. If there is an excess of the reagent, the precipitate becomes green. The reaction is interfered with by Fe 3+ ions, which at high concentrations give the reagent a brown color to the solution, and Mn 2+ and Bi 3+ ions, which give the reagent weakly colored precipitates, soluble in acids. Performing reactions. Place 1–2 drops of FeSO 4 solution in a test tube and add 1 drop of the reagent. Divide the resulting precipitate into two parts, add 1-2 drops of 2 M HC1 solution to the first, and 1-2 drops of 2 M alkali solution to the second. The reaction conditions are with dilute solutions in an acidic environment, pH = 3.

1.5.2.> Oxidation of Fe 2+ to Fe 3+. The Fe 2+ ion is a fairly strong reducing agent and can be oxidized under the action of a number of oxidizing agents, for example, H 2 O 2, KMnO 4, K 2 Cr 2 O 7 in an acidic environment, etc.

2Fe 2+ + 4OH – + H 2 O 2 → 2Fe(OH) 3 ↓.

When conducting a systematic analysis, Fe 2+ should be discovered in preliminary tests, because in the process of group separation, Fe 2+ can be oxidized to Fe 3+.

Particular analytical reactions of Fe 3+ ions

1.5.3. Potassium hexacyanoferrate(II) K 4 with Fe 3+ cations forms a dark blue precipitate of “Prussian blue”:

4Fe 3+ + 3Fe(CN) 6 4– → Fe 4 3 ↓.

The precipitate is practically insoluble in acids, but is decomposed by alkalis to form Fe(OH) 3 . In excess of the reagent, the precipitate dissolves noticeably. Executing the reaction. Add 1 drop of reagent to 1–2 drops of FeCl 3 solution. Divide the resulting precipitate into two parts. Add 2-3 drops of 2 M HC1 solution to one part, 1-2 drops of 2 M NaOH solution to the other, mix.

1.5.4. Potassium thiocyanate (rhodanide) KNCS with Fe 3+ ions forms a blood-red complex. Depending on the concentration of thiocyanate, complexes of different compositions can form:

Fe 3+ + NCS – ↔ Fe(NCS) 2+ ,

Fe 3+ + 2NCS – ↔ Fe(NCS) 2+,

etc. to Fe 3+ + 6NCS – ↔ Fe(NCS) 6 3– ,

The reaction is reversible, so the reagent is taken in excess. Determination is interfered with by ions that form stable complexes with Fe 3+, for example, fluoride ions, salts of phosphoric, oxalic and citric acids.

89. Elements of group I B. Typical properties of the most important compounds, biological role. Bactericidal effect of Ag + and Cu 2+ ions. Analytical reactions to silver and copper ions.

n = 4 Cu ns1(n-1)d10, external level - 1 ē,

preexternal - 18 ē

n = 5 Ag Unpaired ē - one(failure, slippage), but

n = 6 Au 18 - electronic layer, stable in the subgroup

zinc, has not yet completely stabilized here and

capable of losing ē, so COs are possible

Only d-elements of the IB group form compounds in which CO exceeds the N group, and it is more stable for Cu2+, Ag+, Au+3

A characteristic property of doubly charged copper ions is their ability to combine with ammonia molecules to form complex ions. Copper is one of the trace elements. Fe, Cu, Mn, Mo, B, Zn, Co received this name due to the fact that small quantities of them are necessary for the normal functioning of plants. Microelements increase the activity of enzymes, promote the synthesis of sugar, starch, proteins, nucleic acids, vitamins and enzymes. Silver is a low-active metal. In the air atmosphere it does not oxidize either at room temperatures or when heated. The often observed blackening of silver objects is the result of the formation of black silver sulfide - AgS 2 - on their surface. This occurs under the influence of hydrogen sulfide contained in the air, as well as when silver objects come into contact with food products containing sulfur compounds. 4Ag + 2H 2 S + O 2 -> 2Ag 2 S + 2H 2 OV In the voltage range, silver is located much further than hydrogen ode. Therefore, hydrochloric and dilute sulfuric acids have no effect on it. Silver is usually dissolved in nitric acid, which interacts with it according to the equation: Ag + 2HNO 3 -> AgNO 3 + NO 2 + H 2 O Silver forms one series of salts, solutions of which contain colorless Ag + cations. When alkalis act on solutions of silver salts, it is possible expect to obtain AgOH, but instead a brown precipitate of silver(I) oxide precipitates: 2AgNO 3 + 2NaOH -> Ag 2 O + 2NaNO 3 + H 2 O In addition to silver(I) oxide, the oxides AgO and Ag 2 O 3 are known. Silver nitrate (lapis ) - AgNO 3 - forms colorless transparent crystals, highly soluble in water. It is used in the production of photographic materials, in the manufacture of mirrors, in electroplating, and in medicine. Like copper, silver has a tendency to form complex compounds. Many water-insoluble silver compounds (for example: silver(I) oxide - Ag 2 O and silver chloride - Ag Cl) , easily dissolve in an aqueous solution of ammonia. Complex silver cyanide compounds are used for galvanic silver, since during electrolysis of solutions of these salts a dense layer of fine-crystalline silver is deposited on the surface of products. All silver compounds are easily reduced with the release of metallic silver pa. If a little glucose or formalin is added as a reducing agent to an ammonia solution of silver(I) oxide in a glass container, then metallic silver is released in the form of a dense shiny mirror layer on the surface of the glass. Silver ions suppress the development of bacteria and, even in very low concentrations, sterilize drinking water. In medicine, for the disinfection of mucous membranes, colloidal solutions of silver stabilized with special additives (protargol, collargol, etc.) are used. Silver (along with other heavy metals such as copper, tin, mercury) is capable of exerting a bactericidal effect in small concentrations (the so-called oligodynamic effect) . A pronounced bactericidal effect (the ability to reliably kill certain bacteria) is observed at concentrations of silver ions above 0.15 mg/l. In an amount of 0.05 - 0.1 mg/l, silver ions have only a bacteriostatic effect (the ability to inhibit the growth and reproduction of bacteria).Although the disinfection rate of silver is not as high as that of ozone or UV rays, silver ions can remain in water for a long time, providing long-term disinfection. The mechanism of action of silver is not yet fully understood. Scientists believe that the disinfecting effect is observed when positively charged silver and copper ions form electrostatic bonds with the negatively charged surface of microorganism cells. These electrostatic bonds create tension that can impair the permeability of cells and reduce the penetration of vital amounts of nutrients into them. Penetrating inside the cells, silver and copper ions interact with amino acids, which are part of proteins and are used in the process of photosynthesis. As a result, the process of converting solar radiation into food and energy for microorganisms is disrupted, which leads to their death. As a result of numerous studies, the effective bactericidal effect of silver ions on most pathogenic microorganisms, as well as viruses, has been confirmed. However, spore-forming varieties of microorganisms are practically insensitive to silver. Enrichment of water with silver ions can be carried out in several ways: direct contact of water with the surface of silver, treatment of water with a solution of silver salts and the electrolytic method.

Qualitative reaction to copper ions
Potassium hexacyanoferrate (2) K 4 forms with a solution of copper salt a red-brown precipitate of Cu 2, insoluble in dilute acids, but soluble in ammonia solution.
Cu 2+ + 4+ ® Cu 2 ¯To 3 drops of CuSO 4 solution add 2 drops of K 4 salt solution. Observe the formation of a red precipitate. Centrifuge the precipitate and add 3-5 drops of ammonia solution to it.

Reactions for detecting copper ions Cu2+

Action of the group reagent H2S. Hydrogen sulfide forms a black precipitate of copper (II) sulfide in acidified solutions of copper salts: CuS: CuSO4 + H2S = CuS + H2SO4,Cu2+ + H2S = CuS + 2H+.

Action of ammonium hydroxide NH4OH. Ammonium hydroxide NH4OH, taken in excess, forms with copper salts a complex cation of tetraammine copper (II) of intense blue color:

CuSO4 + 4NH4OH = SO4 + 4H2O,

Cu2+ + 4NH4OH = + + 4H2O.

Ag+ silver ion detection reactions

Action of group reagent HC1. Hydrochloric acid forms, with solutions of Ag+ salts, a white precipitate of silver chloride AgCl, which is practically insoluble in water:

Ag+ + Cl- = AgCl.

Detection of silver cation. Hydrochloric acid and solutions of its salts (i.e., Cl- chloride ions) form, with solutions of Ag+ salts, a practically water-insoluble white precipitate of silver chloride AgCl, which dissolves well in an excess of NH4OH solution; in this case, a water-soluble complex silver salt, diammine silver chloride, is formed. With the subsequent action of nitric acid, the complex ion is destroyed and silver chloride precipitates again (these properties of silver salts are used to detect it):

AgNO3 + HCl = AgCl + HNO3,

AgCl + 2NH4OH = Cl + 2H2O,

Cl + 2HNO3 = AgCl + 2NH4NO3.

90. Elements of II B group. Typical properties of the most important compounds, biological role. Complex nature of copper- and zinc-containing enzymes. Analytical reactions for Zn 2+ ions.

Enzymes are natural protein catalysts. Some enzymes have a purely protein composition and do not require any other substances to exhibit their activity. However, there is a large group of enzymes whose activity appears only in the presence of certain non-protein compounds. These compounds are called cofactors. Cofactors can be, for example, metal ions or organic compounds of complex structure - they are usually called coenzymes. It has been established that for the normal functioning of the enzyme, both a coenzyme and a metal ion are sometimes required, forming a ternary complex together with the substrate molecule. Thus, metals are part of biological machines as an irreplaceable part. Magnesium ions are needed to work on the transfer of phosphoric acid residues, and potassium ions are also needed for the same purposes; hydrolysis of proteins requires zinc ions, etc. Below we will examine these issues in detail. Enzymes, as a rule, accelerate the same type of reactions, and only a few of them act on only one specific and single reaction. Such enzymes, which have absolute specificity, include, in particular, urease, which decomposes urea. Most enzymes are not as strict in their choice of substrate. The same hydrolase, for example, is capable of catalyzing the hydrolytic decomposition of several different esters. As the chemical side of biological research deepened and chemists increasingly became assistants and collaborators of biologists, the number of newly discovered enzymes steadily increased; soon they had to be counted not in dozens, but in hundreds. This expansion of the range of biological catalysts caused some difficulties in the classification and nomenclature of enzymes. Previously, enzymes were named according to the substrate on which they acted, with the addition of the ending “aza”. So, if an enzyme acts on the sugar maltose, then it was called “maltase”, if on lactose - “lactase”, etc. Currently, a nomenclature has been adopted in which the name also reflects the chemical function of the enzyme. The "aza" particle is reserved for simple enzymes. If a complex of enzymes is involved in the reaction, the term “system” is used.

Enzymes are divided into six classes:

Oxidoreductases. These are enzymes that catalyze redox reactions. Examples of oxidoreductases include pyruvate dehydrogenase, which removes hydrogen from pyruvic acid, catalase, which decomposes hydrogen peroxide, etc.