Properties of elements VII (17) of the group of the main subgroup. General characteristics of groups of elements General characteristics of elements of group VII a Findings in nature Elements of group 7 of the main subgroup

Manganese subgroup- chemical elements of the 7th group of the periodic table chemical elements(according to the outdated classification, elements of a secondary subgroup of group VII). The group includes transition metals manganese Mn, technetium Tc and rhenium Re. Based on the electronic configuration of the atom, the element also belongs to the same group bohrium Bh, artificially synthesized.

As in other groups, members of this family of elements exhibit patterns of electronic configuration, especially outer shells, resulting in similarities in physical properties and chemical behavior:

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. This metal is an excellent material for the manufacture of electric lamp filaments, which are stronger and more durable than conventional tungsten filaments. 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.

Two of the four members of the group, 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. Many rocks contain small amounts of manganese. At the same time, there are also accumulations of its oxygen compounds, mainly in the form of the mineral pyrolusite MnO 2 . The annual world production of manganese ores is about 5 million tons.

Pure manganese can be obtained by electrolysis of solutions of its salts. About 90% of all manganese production is consumed in the manufacture of various iron-based alloys. Therefore, its high-percentage alloy with iron - ferromanganese (60-90% Mn) - is usually smelted directly from ores, which is then used to introduce manganese into other alloys. Ferromanganese is smelted from a mixture of manganese and iron ores in electric furnaces, and manganese is reduced by carbon according to the reaction:

Technetium is not found in the earth's crust. Very small quantities of it were obtained artificially, and it was found that its chemical properties are much closer to rhenium than to manganese. However, a detailed study of the element and its compounds has not yet been carried out.

The content of rhenium in the earth's crust is very small (9·10−9%). This element is extremely dispersed: even the most rhenium-rich minerals (molybdenites) contain it in quantities usually not exceeding 0.002% by weight. Rhenium and its derivatives have not yet found any widespread use. However, in 2007, global production of rhenium was about 45 tons. It is also a chemically active element.

To VII A group of the periodic system D.I. Mendeleev includes Fluor 9F, Chlorine 17Cl, Bromine 35Br, iodine 53I and Astatine 85At (has no stable isotopes). F, Cl, Br, and are called “halogens” (translated from Greek - salts). This name is due to their property of forming salts when directly interacting with metals.
The electronic configuration of the outer layer is ns2nр5. The change in chemical properties in the series F – Cl – Br – I – At is due to a consistent increase in the sizes of ns-, nр-valence orbitals. With an increase in the atomic number of an element, the density increases, the boiling and melting points increase, the strength of halogenated acids increases, and the reactivity decreases.
Halogens are typical nonmetals; under the influence of reducing agents they are easily converted into halide ions G. The affinity of an atom to an electron decreases down the group. Halogens interact vigorously with metals and form ionic compounds with s-metals. The ionic character of halides weakens somewhat with increasing atomic number of the element and is a consequence of decreasing electronegativity. The more electronegative elements, halogens exhibit positive oxidation states.
The properties of fluorine differ markedly from those of other halogens. It has no vacant d-orbitals, 2s22p5 electrons are weakly shielded from the nucleus, which leads to high electron density, ionization energy, and electronegativity. Therefore, for fluorine only the oxidation state is -1.0, and for other halogens 1 (maximum stability of the compounds), 0, +1, +3, +5, +7, +2, +4, +6 are also possible). The binding energy in the F2 molecule is abnormally low, which makes it very reactive (fluorine reacts directly with all elements except HE, Ne, Ar, forming compounds in which the elements are in the highest possible oxidation states). It should also be noted that the enthalpies of formation of ionic and covalent compounds are high compared to other halogens.
2.2 Being in nature

In the earth's crust, the fluorine content is 6 10-2%, chlorine, bromine, iodine, respectively, 2 10-2; 2 10-4; 4 * 10-5%. Fluorine is found in the form of fluoride (about 30 minerals, the most important are CaF2 (fluorite or fluorspar), 3Ca3 (PO4) 2CaF2 (fluorapatite), Na3 - cryolite). Chlorine forms about 70 of its own minerals, mainly chlorides of light metals (rock salt, halite NaCl; sylvite KCl, carnallite KCl MgCl2 6H2O, etc.). The bulk of halogens is concentrated in the water of seas and oceans. Bromine and iodine are also found in drilling waters and seaweed (for example, in seaweed (kelp), the iodine content reaches 0.45%).
2.3 Physical properties

In gaseous, liquid and solid states, halogens are diatomic G2 molecules. Fluorine is a light yellow gas with a very unpleasant, pungent odor. Chlorine is a green-yellow gas with a pungent odor, bromine is a red-brown heavy liquid with a pungent odor, iodine is black, metallic shiny crystals (when heated, it turns into a purple gas (sublimation) - Figure 2.1. Melting and boiling points increase monotonically from fluorine to iodine with an increase in molecular size and increased intermolecular interaction.

A
would
V
a – chlorine; b – bromine; c – iodine
Figure 2.1 – Appearance of chlorine, bromine, iodine

2.4 Extraction methods

Fluorine is obtained by electrolysis of fluoride melts (mainly KHF2, which allows electrolysis to be carried out at 1000C, while KF melts at a temperature of 8570C.
The industrial production of chlorine is based on the electrolysis of aqueous solutions of NaCl. In laboratory conditions it is obtained by reacting concentrated HCl with oxidizing agents:
MnO2 + 4HCl → MnCl2 + Cl2 + 2H2O
2KMnO4 + 16HCl → 2MnCl2 + 2KCl + 5Cl2 + 8H2O
Particularly pure chlorine is obtained by the reaction:
2AuCl3 → 2Au + 3Cl2
Bromine is industrially obtained from sea water, having previously gotten rid of NaСl: 2Br – + Cl2 → Br2 + 2Cl-
Bromine is blown out with a stream of air and absorbed by iron filings or other substances, for example:
Na2CO3 + Br2 → NaBrO + NaBr + CO2
NaBrO + NaBr + H2SO4 → Br2 + Na2SO4 + H2O
In laboratory conditions, bromine is obtained by the reaction:
2KBr + Cl2 → 2KCl + Br2
Industrially, iodine is also extracted from sea water, oil well water, and sea plant ash:
2NaI + Cl2 → 2NaCl + I2
In the laboratory, iodine is obtained by the reaction:
2NaI + MnO2 + 2H2SO4 → I2 + MnSO4 + Na2SO4 + 2H2O
Iodine is adsorbed with activated carbon or extracted with solvents, and purified by sublimation.
2.5 Chemical properties of elements of group VII A

According to their chemical properties, halogens are active nonmetals. Due to the low dissociation energy of the fluorine molecule, the very electronegativity of the atom and the high hydration energy of the ion, fluorine is a strong oxidizing agent (oxidizes other elements to higher positive oxidation states), reacts vigorously with simple substances with the exception of He, E and A r. In the series from fluorine to iodine, the oxidizing properties decrease, and the reducing properties increase.

Interaction with water:
Fluorine interacts extremely energetically with water:
2F2 + 2H2O → 4HF + O2,
The reaction is accompanied by the formation of ozone and ОF2.
When chlorine is dissolved in water, the reaction occurs:
H2O + Сl2 HOСl + HСl – at room temperature in a saturated solution of Сl2 in water, approximately 70% of chlorine is in the form of molecules, while the equilibrium for iodine is almost completely shifted to the left.
Interaction with complex substances:
Fluorine reacts with alkalis to form OF2:

When chlorine acts on cold alkali solutions, salts of hypochlorous acid are formed:
Сl2 + 2KOH → KOСl + KСl + H2O
potassium hypochlorite
When exposed to a hot alkali solution (70-800C), salts of perchloric acid are formed - chlorates:
3Сl2 + 6KOH → KСlО3 + 5KСl + 3H2O
potassium chlorate
Iodine and bromine also preferentially form trioxohalogenates when reacting with alkalis.
Chlorine reacts with soda solution:
2Na2CO3 + Cl2 + H2O → NaClO + NaCl + 2NaHCO3
“Javel water”
Iodine to a small extent exhibits properties characteristic of metals. This way you can get iodine nitrate, which decomposes at temperatures below 0 ° C.
I2 + AgNO3 AgI + INO3; 3INO3 → I2 + I (NO3) 3
2.6 Halogen compounds

Hydrogen halides
Under standard conditions, hydrogen halides are colorless gases with a pungent odor. With increasing mass and size of molecules, intermolecular interaction increases, and, as a result, the melting and boiling points increase. Hydrogen fluoride has abnormally high melting points (-83 ° C) and boiling points (-19.5 ° C), which is explained by the formation of hydrogen bonds between HF molecules.
Due to their high polarity, hydrogen halides dissolve well in water to form acids, the strength of which increases in the series HF-HCl-HBr-HE (due to an increase in radius). The reducing activity of halide ions in the series F- → SI- → Br- → I- also increases. NO is a strong reducing agent, used in organic synthesis. In air, the aqueous solution of NO is gradually oxidized by atmospheric oxygen:
4HI + O2 → 2I2 + 2H2O
HBr behaves similarly. Hydrofluoric acid (HF) and hydrochloric acid (HCl) do not react with concentrated sulfuric acid, but HBr and NO are oxidized by it.
The main amount of hydrochloric acid is obtained by chlorination, dechlorination of organic compounds, pyrolysis (schedule when heated without air access) of organochlorine waste - by-products of various processes. In addition, hydrogen halides are obtained:
direct synthesis from elements: H2 + G2 2NG
This chain reaction, which also underlies the industrial production of HCl, is initiated by light, moisture, and porous solids.
by displacing NG from their salts (laboratory extraction methods):
CaF2 + H2SO4 → CaSO4 ↓ + 2HF;
NaCl + H2SO4 (k) → NaHSO4 + HCl;
NaHSO4 + NaCl → Na2SO4 + HCl.
– HBr acids are not obtained by hydrolysis of phosphorus halides:
PE3 + 3H2O → H3PO3 + 3HE (E – Br or I).
A feature of HF and its aqueous solutions is the destruction of quartz and glass:
SiO2 + 4HF → SiF4 + 2H2O
SiF4 + 2HF → H2
Therefore, HF is stored in plastic or glass containers, but coated inside with wax or paraffin. Rare HF is a highly ionizing solvent. Mixes with water in any ratio. In dilute aqueous solutions there is an equilibrium:
HF + H2O H3O + + F-;
F- + HF HF2-;
By neutralizing HF, potassium bifluoride (potassium hydrogen fluoride) can be obtained:
2HF + KOH → KHF2 + H2O
KHF2 + KOH → 2KF + H2O
Fluorides (salts of hydrofluoric acid) are poorly soluble in water (with the exception of NaF, KF, NH4F, AgF, SnF2), they are divided, similarly to oxides, into acidic (SiF4), basic (NaF) and amphoteric (AlF3). They can react with each other:
2NaF + SiF4 → Na2
KF + SbF5 → K
3KF + AlF3 → K3
Chlorides - salts of hydrochloric acid - dissolve in water, with the exception of AgCl, HgCl2, Hg2Cl2, PbCl2.
Bromides, iodides - dissolve in water, with the exception of AgBr, AgI, PbI2, PbBr2.
Compounds of halogens and oxygen
Binary oxygen compounds of fluorine are called fluoride (Fluor is more electronegative than oxygen). Oxygen difluoride – ОF2, which is formed by the reaction:
2NaOH + 2F2 → 2NaF + OF2 + H2O
ОF2 – light yellow gas, reactive, strong oxidizing agent:
2H2 + OF2 → H2O + 2HF.
Other halogens in compounds with oxygen exhibit positive oxidation states.
Among the oxides, H2O5 (the only thermodynamically stable halide oxide) is a colorless crystalline substance. Medium strength oxidizer, used for quantitative determination of CO:
I2O5 + 5CO → I2 + 5CO2
I2 + 2Na2S2O3 → 2NaI + Na2S4O6
Oxygen chlorine compounds are obtained indirectly. Relatively stable are Cl2O, ClO2, Cl2O7:
Сl2O is a dark yellow gas with a pungent odor, poisonous, unstable, and can explode. This oxide is obtained by the reaction: 2HgO + 2Cl2 → HgCl2 + Cl2O.
Cl2O reacts with water: Cl2O + H2O → 2HOCl or 2HCl – hypochlorous acid. This acid is unstable and exists only in dilute solution.
HOCl and its hypochlorite salts are strong oxidizing agents:
NaOCl + 2KI + H2SO4 → I2 + NaCl + K2SO4 + H2O
ClO2 is a greenish-yellow gas, with a pungent odor, poisonous, can explode when heated, an energetic oxidizer.
ClO2, the only halogen oxide that is produced on an industrial scale by the reactions:
КClO3 + H2SO4 → HClO3 + KHSO4
3HClO3 → 2ClO2 + HClO4 + H2O
In water, ClO2 is disproportionate, as in alkali solutions:
2СlО2 + H2O → HClO3 + HClО2
perchloric acid chloritic acid
2ClO2 + 2KOH → KClO3 + KClO2 + H2O
Cl2O7 is an oily liquid that explodes when heated to 120 ° C, obtained by the reaction: 4HClO4 + P4O10 → 2Cl2O7 + 4HPO3.
Cl2O7 reacts with water: Cl2O7 + H2O → 2HClO4

Hypohalic acids NPO are known only in dilute aqueous solutions. They are obtained by reacting a halogen with a suspension of mercury oxide:
2I2 + HgO + H2O → HgI2 + 2HOI.
These are weak acids; in the series HOCl → HOBr → HOI, the strength of the acids decreases, while the basic properties increase. HOI is already an amphoteric compound.
Hypohalogenites are unstable compounds with strong oxidizing properties, obtained by reacting G2 with a cooled alkali solution. This is how bleach is produced in industry, for a long time Widely used as a disinfectant and bleaching agent:

Of the oxygenic acids of the halogens NGO2, only chlorous acid HClO2 is known; in the free state, the acid is of medium strength (Kd = 10-2). It has no technical significance. Of practical importance is NaClO2 - a strong oxidizing agent, used as a bleaching agent for fabrics, in a small amount (about 0.4%) it is included in washing powder. Obtained by the reaction:
Na2O2 + 2ClO2 → O2 + 2NaClO2
Oxoacids NGO3 are more stable than NGO. HClO3, HBrO3 exist only in solutions whose concentration does not exceed 50%, and HIO3 is isolated as an individual compound.
In the series HClO3 → HBrO3 → HIO3, the strength of acids decreases; they are weaker oxidizing agents than HOX.
HClO3 is obtained through the reactions:
6Ba (OH) 2 + 6Cl2 → 5BaCl2 + Ba (ClO3) 2 + 6H2O
Ba (ClO3) 2 + H2SO4 → BaSO4 ↓ + 2HClO3
HBrO3 is obtained by the reaction:
Br2 + 5Cl2 + 6H2O → 2HBrO3 + 10HCl
HIO3 can be obtained:
3I2 + 10HNO3 → 6HIO3 + 10NO + 2H2O
Salts of these acids, strong oxidizing agents, are obtained by the reaction:
3G2 + 6KON → KEO3 + 5ke + 3H2O
KClO3 - berthollet salt - is widely used in industry - it is used in the manufacture of matches, fireworks, and explosives.
Oxoacids NGO4
HClO4 is a liquid, smokes in air. It is obtained during the reaction:
KClO4 + H2SO4 → HClO4 + KHSO4
Anhydrous HClO4 is a very strong oxidizing agent, one of the strongest acids that is used in inorganic and organic synthesis. Salts are perchlorates, most of which are soluble in water, with the exception of KClO4, RbClO4, CsClO4, Mg (ClO4) 2 (technical name “Anhydron”) - one of the most powerful desiccants.
Bromic acid is known only in aqueous solutions.
Periodate acid H5IO6 is a weak acid, highly soluble in water, forms medium and acid salts. The acid is obtained by the reaction:
Ba5 (IO6) 2 + 5H2SO4 → 5BaSO4 + 2H5IO6.
Salts of periodate acid can be obtained:
KIO3 + Cl2 + 6KOH → K5IO6 + 2KCl + 3H2O
Interhalogen CONNECTIONS
Unlike elements of other groups, halogens interact with each other to form large quantity interhalogens with the general formula XYn (n = 1, 3, 5,7) - table 2.3, where Y is a lighter and more electronegative halogen. They are obtained by direct interaction of simple substances, at different ratios of reagents, temperatures and pressures.
All interhalides, except BrCl, decompose under the influence of water. They have strong oxidizing properties.
2.7 Use

Halogens and their compounds are widely used in industry, agriculture, and everyday life. In terms of the scale of industrial production, the first place among halogens is occupied by chlorine, and the second by fluorine. The main areas of application of halogens and their compounds are given in Table 2.4
In addition, oxygen-halogen compounds are used in pyrotechnics. Fluorine compounds are used to produce glazes and enamel; HF – for glass etching. Chlorine-containing compounds are widely used as chemical warfare agents (phosgene, mustard gas, chloropicrin, etc.). AgBr is used in photography, KBr is used in optics. Iodine and bromine are used in halogen lamps. Sawing AgI and PbI2 aerosols in clouds causes (artificial) rain and is a means of combating hail. Some organoiodine compounds are used to produce high-power gas lasers.
2.8 Biological role and toxicology

Fluorine and its compounds are extremely toxic. F2 has an irritant effect several times more than HF. When HF gets on the skin, it dissolves proteins, penetrates deeply into tissues, and causes severe ulcers. Fluorine in fluorapatite is part of tooth enamel; its deficiency causes caries, and its excess causes increased fragility of bones.
Chlorine belongs to the group of asphyxiating substances, causes severe irritation of the mucous membranes, and can lead to pulmonary edema. High concentrations lead to reflex inhibition of the respiratory center. Chlorine is the most important biogenic element. Chloride ions are part of gastric juice and participate in various intracellular processes - maintaining osmotic pressure and regulating water-salt metabolism.
Bromine vapor also leads to irritation of the mucous membranes, dizziness, and higher concentrations cause spasms of the respiratory tract and damage to the olfactory nerve. When liquid bromine gets on the skin, very painful burns and ulcers form and are difficult to heal. Bromine compounds regulate the processes of excitation and inhibition of the central nervous system.
Inhalation of iodine vapor causes damage to the kidneys and cardiovascular system, respiratory tract, and pulmonary edema is possible. If it comes into contact with the mucous membrane of the eye, pain in the eyes, redness, and tearing appear. Iodine is part of the thyroid hormones of the thyroid gland (thyroxine, triiodothyronine), which play a very important role in metabolism.

1. What oxidation states do halogens exhibit in compounds? What are the features of fluorine valence states? Why do metals exhibit higher oxidation states in compounds with fluorine?
2. Analyze changes in properties in the halogen series.
3. Illustrate industrial and laboratory methods for producing halogens with reactions.
4. Give a comparative description of the redox properties of halogens using the example of various reactions.
5. How do physical and chemical properties change in the series HF-HCl-HBr-NET?
6. Write the reaction equations for the interaction of halogens with water and alkalis.
7. How do the strength and redox properties of halogen oxygen acids change? Give reasons for your answer.
8. What inorganic compounds of fluorine, chlorine, bromine and iodine are used in medicine? What other industries widely use halogens and their compounds?
9. Write the reaction equations that can be used to carry out transformations:
РbBr2 → HBr → Br2 → КBrO3 → НBrO3 → FeBr3;
Сl2 → КClO3 → КClО4 → НClО4 → ClO2 → НClO3;
Cl2 → HCl → KCl → Cl2 → BaCl2 → HCl.
10. What biological role do halogens play in the human body?

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 on the outer energy level 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 by lead dioxide into nitrogen acidic environment 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.

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.

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.

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

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 decomposes with 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 kind. 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 AgCl) , 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 sulfur sconce. 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.

Group VII p-elements include fluorine ( F), chlorine ( Cl), bromine ( Br), iodine ( I) and astatine ( At). These elements are called halogens (producing salts). All elements of this subgroup are non-metals.

The general electronic formula of the valence band of atoms has the form ns 2 np 5, from which it follows that there are seven electrons on the outer electronic layer of the atoms of the elements under consideration and they can exhibit odd valences 1, 3, 5, 7. The fluorine atom does not have a d-sublevel, therefore there are no excited states and the valency of fluorine is only 1.

Fluorine is the most electronegative element in the periodic table and, accordingly, in compounds with other elements exhibits only a negative oxidation state of –1. The remaining halogens can have oxidation states –1, 0, +1, +3, +5, +7. Each halogen in its period is the most powerful oxidizing agent. With an increase in the atomic number of elements in the series F, C1, Br, I and At, the radii of the atoms increase and the oxidative activity of the elements decreases.

Molecules of simple substances are diatomic: F 2, C1 2, Br 2, I 2. Under normal conditions, fluorine is a pale yellow gas, chlorine is a yellow-green gas, bromine is a red-brown liquid, and iodine is a dark purple crystalline substance. All halogens have a very pungent odor. Inhaling them leads to severe poisoning. When heated, iodine sublimates (sublimes), turning into purple vapor; When cooled, iodine vapor crystallizes, bypassing the liquid state.

Halogens are slightly soluble in water, but much better in organic solvents. Fluorine cannot be dissolved in water, as it decomposes it:

2F 2 + 2H 2 O = 4НF + O 2.

When chlorine is dissolved in water, its partial auto-oxidation-self-reduction occurs according to the reaction

C1 2 + H 2 O ↔ HC1+ HC1O.

The resulting solution is called chlorine water. It has strong acidic and oxidizing properties and is used to disinfect drinking water.

Halogens interact with many simple substances, exhibiting the properties of oxidizing agents. Fluorine reacts explosively with many non-metals:

H 2 + F 2 → 2HF,

Si + 2F 2 → SiF 4,

S + 3F 2 → SF 6.

In a fluorine atmosphere, stable substances such as glass in the form of cotton wool and water burn:

SiO 2 + 2F 2 → SiF 4 + O 2,

2H 2 O + 2F 2 → 4HF + O 2.

Fluorine does not directly interact only with oxygen, nitrogen, helium, neon and argon.

In an atmosphere of chlorine, many metals burn, forming chlorides:

2Na + С1 2 → 2NaCl (bright flash);

Сu + С1 2 → СuС1 2,

2Fe + 3Cl 2 → 2FeCl 3.

Chlorine does not directly interact with N2, O2 and inert gases.

The oxidative activity of halogens decreases from fluorine to astatine, and the reducing activity of halide ions increases in this direction. It follows from this that the more active halogen displaces the less active one from solutions of its salts:

F 2 + 2NaCl → Cl 2 + 2NaF,

Cl 2 + 2NaBr → Br 2 + 2NaCl,

Br 2 + 2NaI → I 2 + 2NaBr.

Hydrogen compounds of halogens are highly soluble in water. Their aqueous solutions are acids:

HF – hydrofluoric (fluoric) acid,

HC1 – hydrochloric acid (aqueous solution – hydrochloric),

НВг – hydrobromic acid,

HI – hydroiodic acid.

HF should be one of the strongest acids, but due to the formation of a hydrogen bond (H–F···H–F) it is a weak acid. Confirmation of the presence of a hydrogen bond between H–F molecules, as in the case of water, is the anomalously high boiling point of H–F.

Hydrofluoric acid reacts with SiO 2, so HF cannot be prepared and stored in glass containers

SiO 2 + 4HF = SiF 4 + 2H 2 O.

The remaining hydrogen halides are strong acids.

Chlorine, bromine and iodine form oxygen-containing acids and their corresponding salts. Below, using chlorine as an example, are the formulas

acids and their corresponding salts:

HClO, HClO 2, HClO 3, HClO 4;

hypochlorous chloride hypochlorous chlorine

strengthening acid properties

KClO, KClO 2, KClO 3, KClO 4.

potassium hypochlorite potassium chlorite potassium chlorate potassium perchlorate

Perchloric and hypochlorous acids are strong, while chloric and hypochlorous acids are weak. Among the salts we can note:

CaOC1 2 - “bleach” is a mixed salt of hydrochloric and hypochlorous acids.

KClO 3 – potassium chlorate, technical name – Berthollet salt.

Fluorine and its compounds are used to produce heat-resistant plastics (Teflon) and refrigerants (freon) for refrigeration machines.

Chlorine is used in large quantities for the production of hydrochloric acid by the synthetic method, organochlorine insecticides, plastics, synthetic fibers, bleach, bleaching fabrics and paper, chlorinating water for disinfection purposes, and chlorinating ores for the extraction of metals.

Bromine and iodine compounds are used to produce medicines, photographic materials.