Halogens: physical properties, chemical properties. The use of halogens and their compounds

Halogens in the periodic table are located to the left of the noble gases. These five toxic non-metallic elements are included 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 it is often ranked as a halogen. Since halogen elements have seven valence electrons, they only need one additional electron to form a complete octet. This characteristic makes them more active than other groups of non-metals.

general characteristics

Halogens form diatomic molecules (of the form 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 nonpolar, covalent and single. The chemical properties of halogens allow them to easily come into contact with most elements, so they never occur in an unbound form in nature. Fluorine is the most active halogen, and astatine is the least.

All halogens form salts of group I 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 ^- called "chloride."

In addition, the chemical properties of halogens allow them to act as oxidizing agents - to oxidize metals. Most of the chemical reactions in which halogens are involved are redox in an aqueous solution. Halogens form single bonds with carbon or nitrogen in organic compounds, where the degree of oxidation (CO) is -1. When a halogen atom is replaced by a covalently bonded hydrogen atom in an organic compound, the halo prefix can be used in the general sense, or the fluorine, chlorine, bromine, iodine prefixes for specific halogens. Halogen elements can have a cross bond to form diatomic molecules with polar covalent single bonds.

Chlorine (Cl ₂ ) became the first halogen discovered in 1774, then iodine (I ₂ ), bromine (Br ₂ ), fluorine (F ₂ ), and astatine (At, discovered last in 1940) were discovered. The name "halogen" comes from the Greek roots hal- ("salt") and -gen ("form"). Together, these words mean "salt-forming", emphasizing the fact that halogens, when reacted with metals, form salts. Halite is the name for rock salt, a natural mineral made up 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 denoted by the symbol F. Elementary fluorine was first discovered in 1886 by isolating it from hydrofluoric acid. In its free state, fluorine exists in the form of a diatomic molecule (F ₂ ) and is the most common halogen in the earth's crust. Fluorine is the most electronegative element in the periodic table. At room temperature it is a pale yellow gas. Fluorine also has a relatively small atomic radius. Its CO is -1, with the exception of an elementary diatomic state in which its oxidation state is zero. Fluorine is extremely chemically active and directly interacts with all elements except helium (He), neon (Ne) and argon (Ar). In a solution of H ₂ O, 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 fluorine ion is basic (pH> 7). In addition, fluorine produces very powerful oxidizing agents. For example, fluorine can react with an inert gas, xenon, and forms a strong oxidizing agent, xenon difluoride (XeF ₂ ). Fluorine has many uses.

Chlorine is an element with atomic number 17 and the chemical symbol Cl. It was discovered in 1774 by isolating it from hydrochloric acid. In its elementary state, it forms a diatomic molecule of 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 the two chlorine atoms is weak, the Cl ₂ molecule has a very high ability to join compounds. Chlorine reacts with metals to form salts called chlorides. Chlorine ions are the most common ions, they are found in sea water. Chlorine also has two isotopes: ³⁵ Cl and ³⁷ Cl. Sodium chloride is the most common compound of all chlorides.

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

Iodine is a chemical element with atomic number 53 and the symbol I. Iodine has oxidation states: -1, +1, +5 and +7. Exists in the form of a diatomic molecule, I ₂ . It is a violet solid at room temperature. Iodine has one stable isotope - ¹²⁷ I. It was first discovered in 1811 using seaweed and sulfuric acid. Currently, iodine ions can be released in seawater. Despite the fact that iodine is not very soluble in water, its solubility can increase with the use of 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. It is 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 a very short half-life, not longer than a few hours. Received in 1940 as a result of synthesis. Astatine is believed to be similar to iodine. It differs in metallic properties.

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

A similar structure of the outer layer of electrons determines 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 the simple substances of halogens change with increasing sequence number of the element. For better assimilation and greater visibility, we offer you several tables.

The melting and boiling points in the group increase as the size of the molecule grows (F <Cl <Br <I <At). This increase means an increase in the strength of Van der Waals.

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

• The atomic radius increases.

The size of the nucleus increases (F <Cl <Br <I <At), as the number of protons and neutrons increases. In addition, more energy levels are added with each period. This leads to a larger orbital, and therefore to an increase in the radius of the atom.

Table 2. Halogens. Physical properties: atomic radii

• The ionization energy is reduced.

If the external valence electrons are not near the nucleus, then to remove them from it will not require much energy. Thus, the energy required to push the external electron is not so high in the lower part of the group of elements, since there are more energy levels. In addition, the high ionization energy makes the element exhibit non-metallic qualities. Iodine and astatine display exhibit metallic properties because the ionization energy decreases (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 progressively farther from the nucleus; Thus, the nucleus and electrons are not attracted to each other. Increased shielding is observed. Therefore, electronegativity decreases with increasing period (At <I <Br <Cl <F).

Table 4. Halogens. Physical properties: electronegativity

• The electron affinity decreases.

Since the size of an atom increases with increasing period, electron affinity, as a rule, decreases (B <I <Br <F <Cl). The exception is fluorine, whose affinity is less than that of chlorine. This can be explained by the smaller size of fluorine in comparison with chlorine.

Table 5. Halogen affinity for the electron

• The reactivity of the elements decreases.

The reactivity of halogens decreases with increasing period (At <I <Br <Cl <F). This is due to an increase in the radius of the atom with increasing electron energy levels. This reduces the attraction of valence electrons of other atoms, reducing the reactivity. This decrease also occurs due to a decrease in electronegativity with an increase in the period, which also reduces the attraction of electrons. In addition, with an increase in atom size, the oxidizing ability decreases.

Inorganic chemistry. Hydrogen + Halogens

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

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

Hydrogen halides easily dissolve in water with the formation of hydrogen halide (hydrofluoric, hydrochloric, hydrobromic, hydroiodic). 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 <HCl <HBr <HI.

Hydrofluoric acid can engrave 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 very weak acid. A strong bond is determined by the short bond length and high dissociation energy. Of all hydrogen halides, HF has the shortest bond length and the largest bond dissociation energy.

Halogen oxo acids

Halogen oxoacids are acids with atoms of hydrogen, oxygen and halogen. Their acidity can be determined using structural analysis. Halogen oxo acids are listed below:

• Hypochlorous acid HOCl.
• Chloric acid HClO ₂ .
• Chloric acid HClO ₃ .
• Perchloric acid HClO ₄ .
• Hybrid acid HOBr.
• Bromic acid HBrO ₃ .
• Bromic acid HBrO ₄ .
• Hybrid acid HOI.
• Hydrochloric acid HIO ₃ .
• Methiodic acid HIO4, H5IO6.

In each of these acids, a proton is bonded to an oxygen atom, so comparing proton bond lengths is useless here. The dominant role here is played by electronegativity. Acidic activity increases with an increase in the number of oxygen atoms bound to the central atom.

Appearance and condition of the substance

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

Explanation of appearance

The color of halogens is the result of the absorption of visible light by molecules, which causes the excitation of electrons. Fluorine absorbs violet light, and therefore looks light yellow. Iodine, on the contrary, absorbs yellow light and looks purple (yellow and purple are complementary colors). Halogen color becomes darker with increasing period.

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

Although the color of astatine is unknown, it is assumed that it should be darker than iodine (i.e. black) in accordance with the observed pattern.

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

The degree of oxidation of halogens in the compounds

The oxidation state is often used instead of the term “halogen valency”. Typically, the oxidation state is -1. But if the halogen is bound to oxygen or another halogen, it can take other states: CO of oxygen -2 takes precedence. In the case of two different halogen atoms bonded together, a more electronegative atom prevails and takes 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 a CO of -8 (-2 x 4 atoms = -8). Hydrogen has a total oxidation state of +1. Adding these values ​​gives CO -7. Since the final CO compound must be zero, then the bromine CO is +7.

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

Why is fluoride always -1?

Electronegativity increases with increasing period. Therefore, fluorine has the highest electronegativity of all elements, which is confirmed by its position in the periodic table. Its electronic configuration is 1s ² 2s ² 2p ⁵ . If fluorine receives another electron, the extreme p-orbitals are completely filled and constitute a complete octet. Since fluorine has a high electronegativity, it can easily take an electron from a neighboring atom. In this case, fluorine is isoelectronic to an inert gas (with eight valence electrons), all of 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, therefore free halogens are obtained by oxidation by electrolysis or by using oxidizing agents. For example, chlorine is produced by hydrolysis of a solution of sodium chloride. The use of halogens and their compounds is diverse.

• Fluoride . Although fluorine is very reactive, it is used in many areas of industry. For example, it is a key component of polytetrafluoroethylene (Teflon) and some other fluoropolymers. Chlorofluorocarbons are organic chemicals that were previously used as refrigerants and propellants in aerosols. Their use ceased due to their possible impact on the environment. They were replaced by hydrochlorofluorocarbons. Fluoride is added to toothpaste (SnF ₂ ) and drinking water (NaF) to prevent tooth decay. This halogen is contained 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 to isolate high-voltage equipment (SF ₆ ).
• Chlorine has also found a variety of uses. It is used for the disinfection of drinking water and swimming pools. Sodium hypochlorite (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 that are used to insulate wiring, pipes, and electronics. In addition, chlorine proved to be 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 for sterilization of hospital equipment and disinfection. In agriculture, chlorine is a component of many commercial pesticides: DDT (dichlorodiphenyl trichloroethane) was used as an agricultural insecticide, but its use was discontinued.

• Bromine , due to its incombustibility, is used to suppress combustion. It is also found in methyl bromide, a pesticide used to store crops and suppress bacteria. However, overuse of methyl bromide was discontinued due to its effects on the ozone layer. Bromine is used in the production of gasoline, film, fire extinguishers, 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, an enlargement of the thyroid gland occurs. For the prevention of goiter, this halogen is added to sodium chloride. Iodine is also used as an antiseptic. Iodine is found in solutions used to clean open wounds, as well as in disinfectant sprays. In addition, silver iodide is important in photography.
• Astatine is a radioactive and rare earth halogen, therefore it is not used anywhere else. Nevertheless, it is believed that this element may help iodine in the regulation of thyroid hormones.