What are halogens? Chemical elements fluorine, chlorine, iodine and astatine. Halogens and their compounds

GENERAL CHARACTERISTICS

Halogens (from the Greek halos - salt and genes - forming) - elements of the main subgroup of group VII periodic table: fluorine, chlorine, bromine, iodine, astatine.

Table. Electronic structure and some properties of halogen atoms and molecules

1) The general electronic configuration of the outer energy level is nS2nP5.
2) With an increase in the atomic number of elements, the radii of atoms increase, electronegativity decreases, non-metallic properties weaken (metallic properties increase); halogens are strong oxidizing agents; the oxidizing ability of elements decreases with increasing atomic mass.
3) Halogen molecules consist of two atoms.
4) With an increase in atomic mass, the color becomes darker, the melting and boiling points, as well as density, increase.
5) The strength of hydrohalic acids increases with increasing atomic mass.
6) Halogens can form compounds with each other (for example, BrCl)

FLUORINE AND ITS COMPOUNDS

Fluorine F2 - discovered by A. Moissan in 1886.

Physical properties

The gas is light yellow in color; t°melting= -219°C, t°boiling= -183°C.

Receipt

Electrolysis of potassium hydrofluoride melt KHF2:

Chemical properties

F2 is the strongest oxidizing agent of all substances:

1. 2F2 + 2H2O ® 4HF + O2
2. H2 + F2 ® 2HF (with explosion)
3. Cl2 + F2 ® 2ClF

Hydrogen fluoride

Physical properties

Colorless gas, highly soluble in water, mp. = - 83.5°C; t°boil. = 19.5°C;

Receipt

CaF2 + H2SO4(conc.) ® CaSO4 + 2HF

Chemical properties

1) A solution of HF in water - weak acid (hydrofluoric):

HF « H+ + F-

Hydrofluoric acid salts - fluorides

2) Hydrofluoric acid dissolves glass:

SiO2 + 4HF ® SiF4+ 2H2O

SiF4 + 2HF ® H2 hexafluorosilicic acid

CHLORINE AND ITS COMPOUNDS

Chlorine Cl2 - discovered by K. Scheele in 1774.

Physical properties

Gas yellow-green color, mp. = -101°C, t°boil. = -34°C.

Receipt

Oxidation of Cl- ions with strong oxidizing agents or electric current:

MnO2 + 4HCl ® MnCl2 + Cl2 + 2H2O
2KMnO4 + 16HCl ® 2MnCl2 + 5Cl2 + 2KCl + 8H2O
K2Cr2O7 + 14HCl ® 2CrCl3 + 2KCl + 3Cl2 + 7H2O

electrolysis of NaCl solution (industrial method):

2NaCl + 2H2O ® H2 + Cl2 + 2NaOH

Chemical properties

Chlorine is a strong oxidizing agent.

1) Reactions with metals:

2Na + Cl2 ® 2NaCl
Ni + Cl2 ® NiCl2
2Fe + 3Cl2 ® 2FeCl3

2) Reactions with non-metals:

H2 + Cl2 –hn® 2HCl
2P + 3Cl2 ® 2PClЗ

3) Reaction with water:

Cl2 + H2O « HCl + HClO

4) Reactions with alkalis:

Cl2 + 2KOH –5°C® KCl + KClO + H2O
3Cl2 + 6KOH –40°C® 5KCl + KClOЗ + 3H2O
Cl2 + Ca(OH)2 ® CaOCl2(bleach) + H2O

5) Displaces bromine and iodine from hydrohalic acids and their salts.

Cl2 + 2KI ® 2KCl + I2
Cl2 + 2HBr ® 2HCl + Br2

Chlorine compounds
Hydrogen chloride

Physical properties

A colorless gas with a pungent odor, poisonous, heavier than air, highly soluble in water (1: 400).
t°pl. = -114°C, t°boil. = -85°C.

Receipt

1) Synthetic method (industrial):

H2 + Cl2 ® 2HCl

2) Hydrosulfate method (laboratory):

NaCl(solid) + H2SO4(conc.) ® NaHSO4 + HCl

Chemical properties

1) A solution of HCl in water - hydrochloric acid - strong acid:

HCl « H+ + Cl-

2) Reacts with metals in the voltage range up to hydrogen:

2Al + 6HCl ® 2AlCl3 + 3H2

3) with metal oxides:

MgO + 2HCl ® MgCl2 + H2O

4) with bases and ammonia:

HCl + KOH ® KCl + H2O
3HCl + Al(OH)3 ® AlCl3 + 3H2O
HCl + NH3 ® NH4Cl

5) with salts:

CaCO3 + 2HCl ® CaCl2 + H2O + CO2
HCl + AgNO3 ® AgCl¯ + HNO3

The formation of a white precipitate of silver chloride, insoluble in mineral acids, is used as a qualitative reaction for the detection of Cl- anions in solution.
Metal chlorides are salts of hydrochloric acid, they are obtained by the interaction of metals with chlorine or the reactions of hydrochloric acid with metals, their oxides and hydroxides; by exchange with certain salts

2Fe + 3Cl2 ® 2FeCl3
Mg + 2HCl ® MgCl2 + H2
CaO + 2HCl ® CaCl2 + H2O
Ba(OH)2 + 2HCl ® BaCl2 + 2H2O
Pb(NO3)2 + 2HCl ® PbCl2¯ + 2HNO3

Most chlorides are soluble in water (with the exception of silver, lead and monovalent mercury chlorides).

Hypochlorous acid HCl+1O
H–O–Cl

Physical properties

Exists only in the form of dilute aqueous solutions.

Receipt

Cl2 + H2O « HCl + HClO

Chemical properties

HClO is a weak acid and a strong oxidizing agent:

1) Decomposes, releasing atomic oxygen

HClO – in the light® HCl + O

2) With alkalis it gives salts - hypochlorites

HClO + KOH ® KClO + H2O

2HI + HClO ® I2¯ + HCl + H2O

Chlorous acid HCl+3O2
H–O–Cl=O

Physical properties

Exists only in aqueous solutions.

Receipt

It is formed by the interaction of hydrogen peroxide with chlorine oxide (IV), which is obtained from Berthollet salt and oxalic acid in H2SO4:

2KClO3 + H2C2O4 + H2SO4 ® K2SO4 + 2CO2 + 2СlO2 + 2H2O
2ClO2 + H2O2 ® 2HClO2 + O2

Chemical properties

HClO2 is a weak acid and a strong oxidizing agent; salts of chlorous acid - chlorites:

HClO2 + KOH ® KClO2 + H2O

2) Unstable, decomposes during storage

4HClO2 ® HCl + HClO3 + 2ClO2 + H2O

Hypochlorous acid HCl+5O3

Physical properties

Stable only in aqueous solutions.

Receipt

Ba (ClO3)2 + H2SO4 ® 2HClO3 + BaSO4¯

Chemical properties

HClO3 - Strong acid and strong oxidizing agent; salts of perchloric acid - chlorates:

6P + 5HClO3 ® 3P2O5 + 5HCl
HClO3 + KOH ® KClO3 + H2O

KClO3 - Berthollet salt; it is obtained by passing chlorine through a heated (40°C) KOH solution:

3Cl2 + 6KOH ® 5KCl + KClO3 + 3H2O

Berthollet's salt is used as an oxidizing agent; When heated, it decomposes:

4KClO3 – without cat® KCl + 3KClO4
2KClO3 –MnO2 cat® 2KCl + 3O2

Perchloric acid HCl+7O4

Physical properties

Colorless liquid, boiling point. = 25°C, temperature = -101°C.

Receipt

KClO4 + H2SO4 ® KHSO4 + HClO4

Chemical properties

HClO4 is a very strong acid and a very strong oxidizing agent; salts of perchloric acid - perchlorates.

HClO4 + KOH ® KClO4 + H2O

2) When heated, perchloric acid and its salts decompose:

4HClO4 –t°® 4ClO2 + 3O2 + 2H2O
KClO4 –t°® KCl + 2O2

BROMINE AND ITS COMPOUNDS

Bromine Br2 - discovered by J. Balard in 1826.

Physical properties

Brown liquid with heavy toxic fumes; has an unpleasant odor; r= 3.14 g/cm3; t°pl. = -8°C; t°boil. = 58°C.

Receipt

Oxidation of Br ions by strong oxidizing agents:

MnO2 + 4HBr ® MnBr2 + Br2 + 2H2O
Cl2 + 2KBr ® 2KCl + Br2

Chemical properties

In its free state, bromine is a strong oxidizing agent; and its aqueous solution - "bromine water" (containing 3.58% bromine) is usually used as a weak oxidizing agent.

1) Reacts with metals:

2Al + 3Br2 ® 2AlBr3

2) Reacts with non-metals:

H2 + Br2 « 2HBr
2P + 5Br2 ® 2PBr5

3) Reacts with water and alkalis:

Br2 + H2O « HBr + HBrO
Br2 + 2KOH ® KBr + KBrO + H2O

4) Reacts with strong reducing agents:

Br2 + 2HI ® I2 + 2HBr
Br2 + H2S ® S + 2HBr

Hydrogen bromide HBr

Physical properties

Colorless gas, highly soluble in water; t°boil. = -67°C; t°pl. = -87°C.

Receipt

2NaBr + H3PO4 –t°® Na2HPO4 + 2HBr

PBr3 + 3H2O ® H3PO3 + 3HBr

Chemical properties

An aqueous solution of hydrogen bromide is hydrobromic acid, which is even stronger than hydrochloric acid. It undergoes the same reactions as HCl:

1) Dissociation:

HBr « H+ + Br -

2) With metals in the voltage series up to hydrogen:

Mg + 2HBr ® MgBr2 + H2

3) with metal oxides:

CaO + 2HBr ® CaBr2 + H2O

4) with bases and ammonia:

NaOH + HBr ® NaBr + H2O
Fe(OH)3 + 3HBr ® FeBr3 + 3H2O
NH3 + HBr ® NH4Br

5) with salts:

MgCO3 + 2HBr ® MgBr2 + H2O + CO2
AgNO3 + HBr ® AgBr¯ + HNO3

Salts of hydrobromic acid are called bromides. The last reaction - the formation of a yellow, acid-insoluble precipitate of silver bromide - serves to detect the Br - anion in solution.

6) HBr is a strong reducing agent:

2HBr + H2SO4(conc.) ® Br2 + SO2 + 2H2O
2HBr + Cl2 ® 2HCl + Br2

Among the oxygen acids of bromine, the weak brominated acid HBr+1O and the strong brominated acid HBr+5O3 are known.
IODINE AND ITS COMPOUNDS

Iodine I2 - discovered by B. Courtois in 1811.

Physical properties

Crystalline substance of dark purple color with a metallic luster.
r= 4.9 g/cm3; t°pl.= 114°C; boiling point = 185°C. Very soluble in organic solvents (alcohol, CCl4).

Receipt

Oxidation of I-ions by strong oxidizing agents:

Cl2 + 2KI ® 2KCl + I2
2KI + MnO2 + 2H2SO4 ® I2 + K2SO4 + MnSO4 + 2H2O

Chemical properties

1) with metals:

2Al + 3I2 ® 2AlI3

2) with hydrogen:

3) with strong reducing agents:

I2 + SO2 + 2H2O ® H2SO4 + 2HI
I2 + H2S ® S + 2HI

4) with alkalis:

3I2 + 6NaOH ® 5NaI + NaIO3 + 3H2O

Hydrogen iodide

Physical properties

Colorless gas with a pungent odor, highly soluble in water, boiling point. = -35°С; t°pl. = -51°C.

Receipt

I2 + H2S ® S + 2HI

2P + 3I2 + 6H2O ® 2H3PO3 + 6HI

Chemical properties

1) A solution of HI in water - strong hydroiodic acid:

HI « H+ + I-
2HI + Ba(OH)2 ® BaI2 + 2H2O

Salts of hydroiodic acid - iodides (for other HI reactions, see the properties of HCl and HBr)

2) HI is a very strong reducing agent:

2HI + Cl2 ® 2HCl + I2
8HI + H2SO4(conc.) ® 4I2 + H2S + 4H2O
5HI + 6KMnO4 + 9H2SO4 ® 5HIO3 + 6MnSO4 + 3K2SO4 + 9H2O

3) Identification of I- anions in solution:

NaI + AgNO3 ® AgI¯ + NaNO3
HI + AgNO3 ® AgI¯ + HNO3

A dark yellow precipitate of silver iodide is formed, insoluble in acids.

Oxygen acids of iodine

Hydrous acid HI+5O3

Colorless crystalline substance, melting point = 110°C, highly soluble in water.

Receive:

3I2 + 10HNO3 ® 6HIO3 + 10NO + 2H2O

HIO3 is a strong acid (salts - iodates) and a strong oxidizing agent.

Iodic acid H5I+7O6

Crystalline hygroscopic substance, highly soluble in water, melting point = 130°C.
Weak acid (salts - periodates); strong oxidizing agent.

general characteristics

The halogens include the five main non-metallic elements, which are located in group VII of the periodic table. This group includes such chemical elements as fluorine F, chlorine Cl, bromine Br, iodine I, astatine At.

Halogens get their name from Greek word, which in translation means salt-forming or “salt-forming”, since in principle most of the compounds that contain halogens are called salts.

Halogens react with almost all simple substances, with the exception of only a few metals. They are quite energetic oxidizing agents, have a very strong and pungent odor, interact well with water, and also have high volatility and high electronegativity. But in nature they can only be found as compounds.

Physical properties of halogens

1. So simple chemical substances, like halogens, consist of two atoms;
2. If we consider halogens under normal conditions, then you should know that fluorine and chlorine are in gaseous state, while bromine is a liquid substance, and iodine and astatine are classified as solid substances.

3. For halogens, the melting point, boiling point and density increase with increasing atomic mass. Also, at the same time, their color changes, it becomes darker.
4. With each increase in the serial number, chemical reactivity and electronegativity decrease and non-metallic properties become weaker.
5. Halogens have the ability to form compounds with each other, such as BrCl.
6. At room temperature, halogens can exist in all three states of matter.
7. It is also important to remember that halogens are quite toxic chemicals.

Chemical properties of halogens

When reacting chemically with metals, halogens act as oxidizing agents. If, for example, we take fluorine, then even under normal conditions it reacts with most metals. But aluminum and zinc ignite even in the atmosphere: +2-1: ZnF2.

Production of halogens

When producing fluorine and chlorine on an industrial scale, electrolysis or salt solutions are used.

If you look closely at the picture below, you will see how chlorine can be produced in the laboratory using an electrolysis unit:

The first picture shows an installation for molten sodium chloride, and the second one for producing a solution of sodium chloride.

This process of electrolysis of molten sodium chloride can be represented in the form of this equation:

With the help of such electrolysis, in addition to producing chlorine, hydrogen and sodium hydroxide are also formed:

Of course, hydrogen is produced in a simpler and cheaper way, which cannot be said about sodium hydroxide. It, just like chlorine, is almost always obtained only through electrolysis of a solution of table salt.

If you look at the picture above, you will see how chlorine can be produced in the laboratory. And it is obtained by reacting hydrochloric acid with manganese oxide:

In industry, bromine and iodine are obtained by replacing these substances with chlorine from bromides and iodides.

Application of halogens

Fluorine, or it would be more correct to call copper fluoride (CuF2), has quite wide application. It is used in the manufacture of ceramics, enamels and various glazes. The Teflon frying pan found in every home and the refrigerant in refrigerators and air conditioners also appeared thanks to fluorine.

In addition to household needs, Teflon is also used in medical purposes, since it is used in the production of implants. Fluorine is necessary in the manufacture of lenses in optics and toothpastes.

Chlorine is also found literally at every step in our lives. The most widespread and widespread use of chlorine is, of course, salt NaCl. It also acts as a detoxifying agent and is used in the fight against ice.

In addition, chlorine is indispensable in the production of plastic, synthetic rubber and polyvinyl chloride, thanks to which we obtain clothing, shoes and others needed in our Everyday life things. It is used in the production of bleaches, powders, dyes, and other household chemicals.

Bromine is generally needed as a photosensitive substance when printing photographs. In medicine it is used as a sedative. Bromine is also used in the production of insecticides and pesticides, etc.

Well, the well-known iodine, which is in every person’s medicine cabinet, is primarily used as an antiseptic. In addition to its antiseptic properties, iodine is present in light sources and is also an assistant for detecting fingerprints on a paper surface.

The role of halogens and their compounds for the human body

Choosing in the store toothpaste Probably each of you paid attention to the fact that its label indicates the content of fluorine compounds. And this is not without reason, since this component is involved in the construction of tooth enamel and bones, and increases the resistance of teeth to caries. It also plays an important role in metabolic processes, participates in the construction of the bone skeleton and prevents the occurrence of such dangerous disease like osteoporosis.

Chlorine also plays an important role in the human body, as it takes an active part in maintaining water-salt balance and maintaining osmotic pressure. Chlorine is involved in the metabolism of the human body, the construction of tissues, and what is also important - in getting rid of excess weight. Hydrochloric acid found in gastric juice great importance has for digestion, since without it the process of digesting food is impossible.

Chlorine is necessary for our body and must be supplied to it daily in the required doses. But if its intake into the body is exceeded or sharply reduced, then we will immediately feel it in the form of swelling, headaches and other unpleasant symptoms that can not only disrupt metabolism, but also cause intestinal diseases.

In humans, it is present in the brain, kidneys, blood and liver a large number of bromine Bromine is used for medical purposes as depressant. But with its overdose there may be adverse consequences that can lead to a depressed state of the nervous system, and in some cases to mental disorders. And a lack of bromine in the body leads to an imbalance between the processes of excitation and inhibition.

Without iodine our thyroid cannot be avoided as it is capable of killing germs entering our body. If there is a deficiency of iodine in the human body, a disease of the thyroid gland called goiter may begin. With this disease there appear quite unpleasant symptoms. A person who has a goiter feels weakness, drowsiness, fever, irritability and loss of strength.

From all this we can conclude that without halogens a person could not only lose many things necessary in everyday life, but without them our body would not be able to function normally.

A subgroup of halogens consists of the elements fluorine, chlorine, bromine and iodine.

The electronic configurations of the outer valence layer of halogens are those of fluorine, chlorine, bromine and iodine, respectively). Such electronic configurations determine the typical oxidizing properties of halogens - all halogens have the ability to add electrons, although when moving to iodine, the oxidizing ability of halogens is weakened.

Under ordinary conditions, halogens exist in the form of simple substances consisting of diatomic molecules of the type with covalent bonds. The physical properties of halogens differ significantly: for example, when normal conditions fluorine is a gas that is difficult to liquefy, chlorine is also a gas, but it liquefies easily, bromine is a liquid, iodine is a solid.

Chemical properties of halogens.

Unlike all other halogens, fluorine in all its compounds exhibits only one oxidation state, 1-, and does not exhibit variable valency. For other halogens, the most characteristic oxidation state is also 1-, however, due to the presence of free -orbitals on external level they can exhibit other odd oxidation states from to due to partial or complete depairing of valence electrons.

Fluorine has the greatest activity. Most metals, even at room temperature, ignite in its atmosphere, releasing a large amount of heat, for example:

Without heating, fluorine also reacts with many non-metals (hydrogen - see above), while also releasing a large amount of heat:

When heated, fluorine oxidizes all other halogens according to the following scheme:

where , and in the compounds the oxidation states of chlorine, bromine and iodine are equal.

Finally, when irradiated, fluorine reacts even with inert gases:

The interaction of fluorine with complex substances also occurs very vigorously. So, it oxidizes water, and the reaction is explosive:

Free chlorine is also very reactive, although its activity is less than that of fluorine. It reacts directly with all simple substances except oxygen, nitrogen and noble gases, for example:

For these reactions, as for all others, the conditions for their occurrence are very important. Thus, at room temperature, chlorine does not react with hydrogen; when heated, this reaction occurs, but turns out to be highly reversible, and with powerful irradiation it proceeds irreversibly (with an explosion) through a chain mechanism.

Chlorine reacts with many complex substances, for example, substitution and addition with hydrocarbons:

Chlorine is capable of upon heating, displace bromine or iodine from their compounds with hydrogen or metals:

and also reacts reversibly with water:

Chlorine, dissolving in water and partially reacting with it, as shown above, forms an equilibrium mixture of substances called chlorine water.

Note also that chlorine on the left side of the last equation has an oxidation state of 0. As a result of the reaction, the oxidation state of some chlorine atoms became 1- (in), for others (in hypochlorous acid). This reaction is an example of a self-oxidation-self-reduction reaction, or disproportionation.

Let us recall that chlorine can react (disproportionate) with alkalis in the same way (see the section “Bases” in § 8).

The chemical activity of bromine is less than fluorine and chlorine, but is still quite high due to the fact that bromine is usually used in a liquid state and therefore its initial concentrations, other things being equal, are greater than those of chlorine. Being a “softer” reagent, bromine is widely used in organic chemistry.

Note that bromine, like chlorine, dissolves in water and, partially reacting with it, forms the so-called “bromine water”, while iodine is practically insoluble in water and is not capable of oxidizing it even when heated; for this reason there is no “iodine water”.

Production of halogens.

The most common technological method for producing fluorine and chlorine is the electrolysis of molten salts (see § 7). Bromine and iodine in industry are usually obtained chemically.

In the laboratory, chlorine is produced by the action of various oxidizing agents on hydrochloric acid, for example:

Oxidation is carried out even more efficiently with potassium permanganate - see the section “Acids” in § 8.

Hydrogen halides and hydrohalic acids.

All hydrogen halides are gaseous under normal conditions. The chemical bond carried out in their molecules is polar covalent, and the polarity of the bond decreases in the series. The bond strength also decreases in this series. Due to their polarity, all hydrogen halides, unlike halogens, are highly soluble in water. So, at room temperature in 1 volume of water you can dissolve about 400 volumes of volumes and about 400 volumes of

When hydrogen halides are dissolved in water, they dissociate into ions, and solutions of the corresponding hydrohalide acids are formed. Moreover, upon dissolution, HCI dissociates almost completely, so the resulting acids are considered strong. In contrast, hydrofluoric acid is weak. This is explained by the association of HF molecules due to the occurrence of hydrogen bonds between them. Thus, the strength of acids decreases from HI to HF.

Since negative ions of hydrohalic acids can only exhibit reducing properties, when these acids interact with metals, the oxidation of the latter can occur only due to ions. Therefore, acids react only with metals that are in the voltage series to the left of hydrogen.

All metal halides, with the exception of Ag and Pb salts, are highly soluble in water. The low solubility of silver halides allows the use of an exchange reaction like

as qualitative for the detection of the corresponding ions. As a result of the reaction, AgCl precipitates as a white precipitate, AgBr - yellowish-white, Agl - bright yellow.

Unlike other hydrohalic acids, hydrofluoric acid reacts with silicon (IV) oxide:

Since silicon oxide is part of glass, hydrofluoric acid corrodes glass, and therefore in laboratories it is stored in containers made of polyethylene or Teflon.

All halogens, except fluorine, can form compounds in which they have a positive oxidation state. The most important of these compounds are the oxygen-containing acids of the halogen type and their corresponding salts and anhydrides.

SUBGROUP VIIA. HALOGENS
FLUORINE, CHLORINE, BROMINE, IODINE, ASTATE
Halogens, and especially fluorine, chlorine and bromine, are of great importance for industry and laboratory practice, both in the free state and in the form of various organic and inorganic compounds. Fluorine is a pale yellow, highly reactive gas that causes respiratory irritation and corrosion of materials. Chlorine is also a caustic, chemically aggressive gas with a dark greenish-yellow color and is less reactive than fluorine. It is widely used in low concentrations to disinfect water (chlorination), and in high concentrations it is poisonous and causes severe irritation to the respiratory tract (chlorine gas was used as a chemical weapon in the First World War). Bromine is a heavy red-brown liquid under normal conditions, but evaporates easily to become a corrosive gas. Iodine is a dark purple solid that easily sublimes. Astatine is a radioactive element, the only halogen that does not have a stable isotope.
In the family of these elements, compared to other A-subgroups, non-metallic properties are most pronounced. Even heavy iodine is a typical nonmetal. The first member of the family, fluorine, exhibits “supermetallic” properties. All halogens are electron acceptors and have a strong tendency to complete an octet of electrons by accepting one electron. The reactivity of halogens decreases with increasing atomic number, and in general the properties of halogens vary according to their position in the periodic table. In table 8a shows some physical properties, allowing us to understand the differences and patterns of changes in properties in the series of halogens. Fluorine exhibits unusual properties in many ways. For example, it has been established that the electron affinity of fluorine is not as high as that of chlorine, and this property should indicate the ability to accept an electron, i.e. for chemical activity. Fluorine, due to the very small radius and proximity of the valence shell to the nucleus, should have the highest electron affinity. This discrepancy is, at least in part, explained by the unusual low energy FF bond compared to this value for ClCl (see enthalpy of dissociation in Table 8a). For fluorine it is 159 kJ/mol, and for chlorine 243 kJ/mol. Due to the small covalent radius of fluorine, the proximity of lone electron pairs in the structure: F: F: determines the ease of breaking this bond. Indeed, fluorine is chemically more active than chlorine due to the ease of formation of atomic fluorine. The value of hydration energy (see Table 8a) indicates the high reactivity of the fluoride ion: the F ion is hydrated with a greater energetic effect than other halogens. The small radius and correspondingly higher charge density explain the higher hydration energy. Many of the unusual properties of fluorine and fluoride ions become clear when the size and charge of the ion are taken into account.
Receipt. The great industrial importance of halogens places certain demands on the methods of their production. Given the variety and complexity of production methods, the consumption and cost of electricity, raw materials and the need for by-products are significant.
Fluorine. Due to the chemical aggressiveness of fluoride and chloride ions, these elements are obtained electrolytically. Fluorine is obtained from fluorite: CaF2, when treated with sulfuric acid, forms HF (hydrofluoric acid); KHF2 is synthesized from HF and KF, which is subjected to electrolytic oxidation in an electrolyzer with separate anode and cathode spaces, with a steel cathode and a carbon anode; fluorine F2 is released at the anode, and hydrogen is a byproduct at the cathode, which should be isolated from fluorine to avoid an explosion. To synthesize such important compounds as polyfluorocarbons, organic compounds are fluorinated in the electrolyzer with the released fluorine, so that isolation and accumulation of fluorine in separate containers is not required.
Chlorine produced mainly from NaCl brine in electrolysers with a separated anode space to prevent the reaction of chlorine with other electrolysis products: NaOH and H2; Thus, electrolysis produces three important industrial products: chlorine, hydrogen and alkali. To carry out this process, various modifications of electrolyzers are used. Chlorine is also obtained as a by-product during the electrolytic production of magnesium from MgCl2. Most of the chlorine is used to synthesize HCl by reaction with natural gas, and HCl is consumed to produce MgCl2 from MgO. Chlorine is also formed in sodium metallurgy from NaCl, but the electrolysis method from brine is cheaper. In laboratories industrially developed countries produce many thousands of tons of chlorine by the reaction 4HCl + MnO2 = MnCl2 + 2H2O + Cl2.
Bromine obtained from brine wells that contain more bromide ions than seawater, the second most important source of bromine. Bromide ion is more easily converted to bromine than fluoride and chloride ions in similar reactions. Therefore, to obtain bromine, in particular, chlorine is used as an oxidizing agent, since the activity of halogens in a group decreases from top to bottom and each previously standing halogen displaces the next one. In the production of bromine, brines or sea water are pre-acidified with sulfuric acid and then treated with chlorine according to the reaction
2Br+ Cl2 -> Br2 + 2Cl
Bromine is separated from the solution by evaporation or purging, followed by its absorption by various reagents, depending on the further application. For example, when reacting with a heated solution of sodium carbonate, crystalline NaBr and NaBrO3 are obtained; When the mixture of crystals is acidified, the bromine is regenerated, providing an indirect but convenient method of accumulating (storing) this corrosive, foul-smelling, toxic liquid. Bromine can also be absorbed by SO2 solution, which produces HBr. Bromine can be easily isolated from this solution by passing chlorine (for example, to react bromine with ethylene C2H4 to produce dibromoethylene C2H4Br2, which is used as an antiknock agent for gasoline). World bromine production is over 300,000 tons/year.
Iodine obtained from seaweed ash, treating it with a mixture of MnO2 + H2SO4, and purified by sublimation. Iodide is found in significant quantities in underground drilling waters. Iodine is obtained by oxidation of iodide ion (for example, nitrite ion NO2 or chlorine). Iodine can also be precipitated in the form of AgI, from which silver is regenerated by reacting with iron to form FeI2. Iodine is replaced from FeI2 by chlorine. Chilean saltpeter, which contains an admixture of NaIO3, is processed to produce iodine. Iodide ion is an important component of human food, as it is necessary for the formation of the iodine-containing hormone thyroxine, which controls growth and other body functions.
Reactivity and compounds. All halogens react directly with metals to form salts, the ionic character of which depends on both the halogen and the metal. Thus, metal fluorides, especially metals of subgroups IA and IIA, are ionic compounds. The degree of ionicity of the bond decreases with increasing atomic mass of the halogen and decreasing reactivity of the metal. Halides with an ionic bond type crystallize in three-dimensional crystal lattices. For example, NaCl (table salt) has a cubic lattice. With increasing bond covalency, the proportion of layered structures increases (as in CdCl2, CuCl2, CuBr2, PbCl2, PdCl2, FeCl2, etc.). In the gaseous state, covalent halides often form dimers, for example Al2Cl6 (AlCl3 dimer). With non-metals, halogens form compounds with almost purely covalent bonds, for example, carbon, phosphorus and sulfur halides (CCl4, etc.). Nonmetals and metals exhibit maximum oxidation states in reactions with fluorine, for example SF6, PF5, CuF3, CoF3. Attempts to obtain iodides of similar composition fail because of the large atomic radius of iodine (steric factor) and because of the strong tendency of elements in high oxidation states to oxidize I to I2. In addition to direct synthesis, halides can be obtained by other methods. Metal oxides in the presence of carbon react with halogens to form halides (for example, Cr2O3 turns into CrCl3). It is not possible to obtain CrCl3 from CrCl3×6H2O by dehydration, but only basic chloride (or hydroxochloride). Halides are also obtained by treating oxides with HX vapors, for example:

A good chlorinating agent is CCl4, for example for converting BeO to BeCl2. SbF3 is often used for fluoridation of chlorides (see SO2ClF above).
Polyhalides. Halogens react with many metal halides to form polyhalide compounds containing large anionic Xn1 species. For example:

The first reaction provides a convenient method for preparing a highly concentrated solution of I2 by adding iodine to a concentrated solution of KI. Polyiodides retain the properties of I2. It is also possible to obtain mixed polyhalides: RbI + Br2 -> RbIBr2 RbICl2 + Cl2 -> RbICl4
Solubility. Halogens have some solubility in water, but, as would be expected, due to the covalent nature of the XX bond and the small charge, their solubility is low. Fluorine is so active that it pulls an electron pair from the oxygen in water, releasing free O2 and forming OF2 and HF. Chlorine is less active, but reacts with water to produce some HOCl and HCl. Chlorine hydrates (for example, Cl2*8H2O) can be released from solution upon cooling.
Iodine exhibits unusual properties when dissolved in various solvents. When small amounts of iodine are dissolved in water, alcohols, ketones and other oxygen-containing solvents, a brown solution is formed (a 1% solution of I2 in alcohol is a common medical antiseptic). A solution of iodine in CCl4 or other oxygen-free solvents has a purple color. It can be assumed that in such a solvent, iodine molecules behave similar to their state in the gas phase, which has the same color. In oxygen-containing solvents, the electron pair of oxygen is withdrawn to the valence orbitals of iodine.
Oxides. Halogens form oxides. No systematic pattern or periodicity is observed in the properties of these oxides. The similarities and differences, as well as the main methods for producing halogen oxides, are listed in table. 8b.
Oxoacids of halogens. When oxoacids are formed, the systematic nature of halogens becomes more clearly evident. Halogens form halogenated acids HOX, halogenated acids HOXO, halogenated acids HOXO2 and halogenated acids HOXO3, where X is a halogen. But only chlorine forms acids of all specified compositions, and fluorine does not form oxoacids at all, bromine does not form HBrO4. The compositions of acids and the main methods for their preparation are listed in table. 8th century

All halogen acids are unstable, but pure HOClO3 is the most stable (in the absence of any reducing agents). All oxoacids are strong oxidizing agents, but the rate of oxidation does not necessarily depend on the oxidation state of the halogen. Thus, HOCl (ClI) is a fast and effective oxidizing agent, but dilute HOClO3 (ClVII) is not. In general, the higher the oxidation state of the halogen in an oxoacid, the stronger the acid, so HClO4 (ClVII) is the strongest known oxoacid in aqueous solution. The ClO4 ion, formed during the dissociation of an acid in water, is the weakest of the negative ions as an electron pair donor. Na and Ca hypochlorites find industrial use in bleaching and water treatment. Interhalogen compounds are connections of different halogens to each other. A halogen with a large radius always has a positive oxidation state in such a compound (is subject to oxidation), and with a smaller radius it is more negative (subject to reduction). This fact follows from the general trend of changes in activity in the halogen series. In table Figure 8d shows the compositions of known interhalogen compounds (A is a halogen with a more positive oxidation state).
Interhalogen compounds are formed by direct synthesis from elements. The oxidation state 7, which is unusual for iodine, is realized in the compound IF7, and other halogens cannot coordinate 7 fluorine atoms. BrF3 and ClF3 liquid substances, chemically similar to fluorine, but more convenient for fluoridation, are of practical importance. In this case, BrF3 is more effective. Since trifluorides are strong oxidizing agents and are in a liquid state, they are used as oxidizing agents for rocket fuel.
Hydrogen compounds. Halogens react with hydrogen, forming HX, and with fluorine and chlorine the reaction proceeds explosively with slight activation. The interaction with Br2 and I2 occurs more slowly. For a reaction with hydrogen to occur, it is sufficient to activate a small fraction of the reagents using light or heat. Activated particles interact with non-activated ones, forming HX and new activated particles, which continue the process, and the reaction of two activated particles in the main reaction ends with the formation of a product. For example, the formation of HCl from H2 and Cl2:

More convenient methods for producing hydrogen halides than direct synthesis are provided, for example, by the following reactions:

In the gaseous state, HX are covalent compounds, but in aqueous solution they (with the exception of HF) become strong acids. This is explained by the fact that water molecules effectively pull hydrogen away from the halogen. All acids are highly soluble in water due to hydration: HX + H2O -> H3O+ + X
HF is more prone to complex formation than other hydrogen halides. The charges on H and F are so large, and these atoms are so small, that the formation of HX-associates such as polymers of composition (HF)x, where x = 3. In such a solution, dissociation under the influence of a water molecule occurs by no more than a few percent of the total number of hydrogen ions. Unlike other hydrogen halides, hydrogen fluoride reacts actively with SiO2 and silicates, releasing gaseous SiF4. Therefore, an aqueous solution of HF (fluoric acid) is used in glass etching and stored not in glass, but in paraffin or polyethylene containers. Pure HF boils just below room temperature (19.52°C), so it is stored as a liquid in steel cylinders. An aqueous solution of HCl is called hydrochloric acid. A saturated solution containing 36% (wt.) HCl is widely used in chemical industry and laboratories (see also HYDROGEN).
Astatine This chemical element of the halogen family has the symbol At and atomic number 85, it exists only in trace amounts in some minerals. Back in 1869, D.I. Mendeleev predicted its existence and the possibility of discovery in the future. Astatine was discovered by D. Corson, K. Mackenzie and E. Segre in 1940. More than 20 isotopes are known, of which the longest-lived are 210At and 211At. According to some data, when 20983Bi is bombarded with helium nuclei, the isotope astatine-211 is formed; Astatine has been reported to be soluble in covalent solvents, can form At like other halogens, and is likely to produce the AtO4 ion. (These data were obtained using solutions with a concentration of 1010 mol/l.)

The halogens are located to the left of the noble gases in the periodic table. These five toxic non-metallic elements are in group 7 of the periodic table. These include fluorine, chlorine, bromine, iodine and astatine. Although astatine is radioactive and has only short-lived isotopes, it behaves like iodine and is often classified as a halogen. Since halogen elements have seven valence electrons, they only need one extra electron to form a full octet. This characteristic makes them more reactive than other groups of nonmetals.

general characteristics

Halogens form diatomic molecules (type X 2, where X denotes a halogen atom) - a stable form of existence of halogens in the form of free elements. The bonds of these diatomic molecules are non-polar, covalent and single. allow them to easily combine with most elements, so they are never found uncombined in nature. Fluorine is the most active halogen, and astatine is the least.

All halogens form group I salts with similar properties. In these compounds, halogens are present in the form of halide anions with a charge of -1 (for example, Cl -, Br -). The ending -id indicates the presence of halide anions; for example Cl - is called "chloride".

Besides, Chemical properties halogens allow them to act as oxidizing agents - oxidizing metals. Majority chemical reactions, in which halogens participate - redox in an aqueous solution. Halogens form single bonds with carbon or nitrogen where their oxidation number (CO) is -1. When a halogen atom is replaced by a covalently bonded hydrogen atom in organic compound, the prefix halo- can be used in a general sense, or the prefixes fluoro-, chloro-, bromo-, iodo- - for specific halogens. Halogen elements can cross-link to form diatomic molecules with polar covalent single bonds.

Chlorine (Cl2) was the first halogen discovered in 1774, followed by iodine (I2), bromine (Br2), fluorine (F2) and astatine (At, discovered last, in 1940). The name "halogen" comes from the Greek roots hal- ("salt") and -gen ("to form"). Together these words mean “salt-forming,” emphasizing the fact that halogens react with metals to form salts. Halite is the name rock salt, a natural mineral consisting of sodium chloride (NaCl). And finally, halogens are used in everyday life - fluoride is found in toothpaste, chlorine disinfects drinking water, and iodine promotes the production of thyroid hormones.

Chemical elements

Fluorine, an element with atomic number 9, is designated by the symbol F. Elemental fluorine was first discovered in 1886 by isolating it from hydrofluoric acid. In its free state, fluorine exists as a diatomic molecule (F2) and is the most abundant halogen in the earth's crust. Fluorine is the most electronegative element on the periodic table. At room temperature it is a pale yellow gas. Fluorine also has a relatively small atomic radius. Its CO is -1, except in the elemental diatomic state, in which its oxidation state is zero. Fluorine is extremely reactive and reacts directly with all elements except helium (He), neon (Ne) and argon (Ar). In H2O solution, hydrofluoric acid (HF) is a weak acid. Although fluorine is highly electronegative, its electronegativity does not determine acidity; HF is a weak acid due to the fact that the fluoride ion is basic (pH > 7). In addition, fluorine produces very powerful oxidizing agents. For example, fluorine can react with the inert gas xenon to form the strong oxidizing agent xenon difluoride (XeF2). Fluoride has many uses.

Chlorine is an element with atomic number 17 and the chemical symbol Cl. Discovered in 1774 by isolating it from hydrochloric acid. In its elemental state it forms the diatomic molecule Cl 2 . Chlorine has several COs: -1, +1, 3, 5 and 7. At room temperature it is a light green gas. Since the bond that forms between two chlorine atoms is weak, the Cl 2 molecule has a very high ability to form compounds. Chlorine reacts with metals to form salts called chlorides. Chlorine ions are the most common ions found in sea ​​water. Chlorine also has two isotopes: 35 Cl and 37 Cl. Sodium chloride is the most common compound of all the chlorides.

Bromine is a chemical element with atomic number 35 and symbol Br. It was first discovered in 1826. In its elemental form, bromine is a diatomic molecule Br 2 . At room temperature it is a reddish-brown liquid. Its COs are -1, + 1, 3, 4 and 5. Bromine is more active than iodine, but less active than chlorine. In addition, bromine has two isotopes: 79 Br and 81 Br. Bromine is found in bromide dissolved in seawater. Global production of bromide has increased significantly in recent years due to its availability and long shelf life. Like other halogens, bromine is an oxidizing agent and is very toxic.

Iodine is a chemical element with atomic number 53 and symbol I. Iodine has oxidation states: -1, +1, +5 and +7. Exists in the form of a diatomic molecule, I 2. At room temperature it is a purple solid. Iodine has one stable isotope - 127 I. It was first discovered in 1811 using seaweed and sulfuric acid. Currently, iodine ions can be isolated in seawater. Although iodine is not very soluble in water, its solubility can be increased by using individual iodides. Iodine plays an important role in the body, participating in the production of thyroid hormones.

Astatine is a radioactive element with atomic number 85 and the symbol At. Its possible oxidation states are -1, +1, 3, 5 and 7. The only halogen that is not a diatomic molecule. Under normal conditions it is a black metallic solid. Astatine is a very rare element, so little is known about it. In addition, astatine has very short period half-life, no longer than several hours. Obtained in 1940 as a result of synthesis. Astatine is believed to be similar to iodine. Is different

The table below shows the structure of halogen atoms and the structure of the outer layer of electrons.

This structure of the outer layer of electrons means that the physical and chemical properties of halogens are similar. However, when comparing these elements, differences are also observed.

Periodic properties in the halogen group

The physical properties of simple halogen substances change with increasing atomic number of the element. For better understanding and greater clarity, we offer you several tables.

The melting and boiling points of a group increase as the molecular size increases (F

Table 1. Halogens. Physical properties: melting and boiling points

• The atomic radius increases.

Kernel size increases (F< Cl < Br < I < At), так как увеличивается число протонов и нейтронов. Кроме того, с каждым периодом добавляется всё больше уровней энергии. Это приводит к большей орбитали, и, следовательно, к увеличению радиуса атома.

Table 2. Halogens. Physical properties: atomic radii

• The ionization energy decreases.

If the outer valence electrons are not located near the nucleus, then it will not take much energy to remove them from it. Thus, the energy required to eject an outer electron is not as high in the lower part of the element group, since there are more energy levels there. Additionally, high ionization energy causes the element to exhibit non-metallic qualities. Iodine and display astatine exhibit metallic properties because the ionization energy is reduced (At< I < Br < Cl < F).

Table 3. Halogens. Physical properties: ionization energy

• Electronegativity decreases.

The number of valence electrons in an atom increases with increasing energy levels at progressively lower levels. Electrons are progressively further away from the nucleus; Thus, the nucleus and electrons are not attracted to each other. An increase in shielding is observed. Therefore, Electronegativity decreases with increasing period (At< I < Br < Cl < F).

Table 4. Halogens. Physical properties: electronegativity

• Electron affinity decreases.

As atomic size increases with increasing period, electron affinity tends to decrease (B< I < Br < F < Cl). Исключение - фтор, сродство которого меньше, чем у хлора. Это можно объяснить меньшим размером фтора по сравнению с хлором.

Table 5. Electron affinity of halogens

• The reactivity of the elements decreases.

The reactivity of halogens decreases with increasing period (At

Hydrogen + halogens

A halide is formed when a halogen reacts with another, less electronegative element to form a binary compound. Hydrogen reacts with halogens, forming halides of the form HX:

• hydrogen fluoride HF;
• hydrogen chloride HCl;
• hydrogen bromide HBr;
• Hydrogen iodide HI.

Hydrogen halides easily dissolve in water to form hydrohalic acid (hydrofluoric, hydrochloric, hydrobromic, hydroiodic) acid. The properties of these acids are given below.

Acids are formed by the following reaction: HX (aq) + H 2 O (l) → X - (aq) + H 3 O + (aq).

All hydrogen halides form strong acids, with the exception of HF.

The acidity of hydrohalic acids increases: HF

Hydrofluoric acid can etch glass and some inorganic fluorides for a long time.

It may seem counterintuitive that HF ​​is the weakest hydrohalic acid, since fluorine has the highest electronegativity. However, the H-F bond is very strong, resulting in a very weak acid. A strong bond is determined by a short bond length and high dissociation energy. Of all the hydrogen halides, HF has the shortest bond length and the highest bond dissociation energy.

Halogen oxoacids

Halogen oxo acids are acids with hydrogen, oxygen and halogen atoms. Their acidity can be determined by structural analysis. The halogen oxo acids are given below:

• Hypochlorous acid HOCl.
• Chlorous acid HClO 2.
• Hypochlorous acid HClO 3.
• Perchloric acid HClO 4.
• Hypobromous acid HOBr.
• Bromic acid HBrO 3.
• Bromic acid HBrO 4.
• Hydrous acid HOI.
• Hydrous acid HIO 3.
• Metaiodic acid HIO4, H5IO6.

In each of these acids, a proton is bonded to an oxygen atom, so comparing proton bond lengths is not useful here. Electronegativity plays a dominant role here. Acid activity increases with the number of oxygen atoms associated with the central atom.

Appearance and state of the substance

The basic physical properties of halogens can be summarized in the following table.

Explanation of appearance

The color of halogens results from the absorption of visible light by molecules, which causes electrons to be excited. Fluoride absorbs violet light and therefore appears light yellow. Iodine, on the other hand, absorbs yellow light and appears violet (yellow and violet are complementary colors). The color of halogens becomes darker as the period increases.

In closed containers, liquid bromine and solid iodine are in equilibrium with their vapors, which can be observed in the form of a colored gas.

Although the color of astatine is unknown, it is assumed to be darker than iodine (i.e., black) according to the observed pattern.

Now, if you are asked: “Characterize the physical properties of halogens,” you will have something to say.

Oxidation state of halogens in compounds

Oxidation number is often used instead of the concept of halogen valence. Typically, the oxidation state is -1. But if a halogen is bonded to oxygen or another halogen, it can take other states: oxygen CO -2 takes precedence. In the case of two different halogen atoms bonded together, the more electronegative atom prevails and accepts CO -1.

For example, in iodine chloride (ICl), chlorine has CO -1, and iodine +1. Chlorine is more electronegative than iodine, so its CO is -1.

In bromic acid (HBrO 4), oxygen has CO -8 (-2 x 4 atoms = -8). Hydrogen has an overall oxidation state of +1. Adding these values ​​gives an CO of -7. Since the final CO of the compound must be zero, the CO of bromine is +7.

The third exception to the rule is the oxidation state of the halogen in elemental form (X 2), where its CO is zero.

Why is CO fluorine always -1?

Electronegativity increases with increasing period. Fluorine therefore has the highest electronegativity of all the elements, as evidenced by its position on the periodic table. Its electron configuration is 1s 2 2s 2 2p 5. If fluorine gains another electron, the outermost p orbitals are completely filled and form a full octet. Since fluorine has high electronegativity, it can easily take an electron from a neighboring atom. Fluorine in this case is isoelectronic to the inert gas (with eight valence electrons), all its outer orbitals are filled. In this state, fluorine is much more stable.

Production and use of halogens

In nature, halogens are in the state of anions, so free halogens are obtained by oxidation by electrolysis or using oxidizing agents. For example, chlorine is produced by hydrolysis of a solution of table salt. The use of halogens and their compounds is diverse.

• Fluorine. Although fluorine is very reactive, it is used in many industrial applications. For example, it is a key component of polytetrafluoroethylene (Teflon) and some other fluoropolymers. Chlorofluorocarbons are organic compounds that were previously used as refrigerants and propellants in aerosols. Their use has been discontinued due to their possible environmental impact. They have been replaced by hydrochlorofluorocarbons. Fluoride is added to toothpaste (SnF 2) and drinking water (NaF) to prevent tooth decay. This halogen is found in clay used for the production of certain types of ceramics (LiF), used in nuclear energy (UF 6), to produce the antibiotic fluoroquinolone, aluminum (Na 3 AlF 6), and for insulating high-voltage equipment (SF 6).
• Chlorine also found various applications. It is used to disinfect drinking water and swimming pools. (NaClO) is the main component of bleaches. Hydrochloric acid is widely used in industry and laboratories. Chlorine is present in polyvinyl chloride (PVC) and other polymers used to insulate wiring, pipes and electronics. In addition, chlorine has proven useful in the pharmaceutical industry. Medicines containing chlorine are used to treat infections, allergies and diabetes. The neutral form of hydrochloride is a component of many drugs. Chlorine is also used to sterilize hospital equipment and disinfect. In agriculture, chlorine is a component of many commercial pesticides: DDT (dichlorodiphenyltrichloroethane) was used as an agricultural insecticide, but its use has been phased out.

• Bromine, due to its non-flammability, is used to suppress combustion. It is also found in methyl bromide, a pesticide used to preserve crops and kill bacteria. However, overuse has been phased out due to its impact on the ozone layer. Bromine is used in the production of gasoline, photographic film, fire extinguishers, and drugs for the treatment of pneumonia and Alzheimer's disease.
• Iodine plays an important role in the proper functioning of the thyroid gland. If the body does not receive enough iodine, the thyroid gland becomes enlarged. To prevent goiter, this halogen is added to table salt. Iodine is also used as an antiseptic. Iodine is found in solutions used to clean open wounds, as well as in disinfectant sprays. Additionally, silver iodide is important in photography.
• Astatine- radioactive and rare earth halogen, therefore not yet used anywhere. However, it is believed that this element may help iodine regulate thyroid hormones.