Halogen gas. What are halogens? Chemical properties and significance of halogens

Chemistry of Elements

Nonmetals of VIIA subgroup

Elements of the VIIA subgroup are typical nonmetals with high

electronegativity, they have a group name - “halogens”.

Main issues covered in the lecture

General characteristics of non-metals of the VIIA subgroup. Electronic structure, the most important characteristics of atoms. The most characteristic ste-

oxidation penalties. Features of the chemistry of halogens.

Simple substances.

Natural compounds.

Halogen compounds

Hydrohalic acids and their salts. Salt and hydrofluoric acid

slots, receipt and application.

Halide complexes.

Binary oxygen compounds of halogens. Instability approx.

Redox properties of simple substances and co-

unities. Disproportionation reactions. Latimer diagrams.

Chemistry of elements of the VIIA subgroup

general characteristics

VIIA-group is formed by p-elements: fluorine F, chlorine

Cl, bromine Br, iodine I and astatine At.

The general formula for valence electrons is ns 2 np 5.

All elements of group VIIA are typical non-metals.

to form a stable eight-electron shell

boxes, that's why they have there is a strong tendency towards

addition of an electron.

All elements easily form simple single-charge

ny anions G – .

In the form of simple anions, elements of group VIIA are found in natural water and in crystals of natural salts, for example, halite NaCl, sylvite KCl, fluorite

CaF2.

General group name of elements VIIA-

group “halogens”, i.e. “giving birth to salts”, is due to the fact that most of their compounds with metals are pre-

is a typical salt (CaF2, NaCl, MgBr2, KI), which

which can be obtained through direct interaction

interaction of metal with halogen. Free halogens are obtained from natural salts, so the name “halogens” is also translated as “born from salts.”

The minimum oxidation state (–1) is the most stable

for all halogens.

Some characteristics of the atoms of Group VIIA elements are given in

The most important characteristics of atoms of elements of group VIIA

Halogens have a high electron affinity (maximum at

Cl) and very high ionization energy (maximum at F) and maximum

possible electronegativity in each period. Fluorine is the most

electronegative of all chemical elements.

The presence of one unpaired electron in halogen atoms determines

represents the union of atoms in simple substances into diatomic molecules Г2.

For simple substances, halogens, the most characteristic oxidizing agents are

properties, which are strongest in F2 and weaken when moving to I2.

Halogens are characterized by the greatest reactivity of all non-metallic elements. Fluorine, even among halogens, stands out

has extremely high activity.

The element of the second period, fluorine, differs most strongly from the other

other elements of the subgroup. This general pattern for all non-metals.

Fluorine, as the most electronegative element, does not show sex

resident oxidation states. In any connection, including with ki-

oxygen, fluorine is in the oxidation state (-1).

All other halogens exhibit positive oxidation degrees

leniya up to a maximum of +7.

The most characteristic oxidation states of halogens:

F: -1, 0;

Cl, Br, I: -1, 0, +1, +3, +5, +7.

Cl has known oxides in which it is found in oxidation states: +4 and +6.

The most important halogen compounds, in positive states,

Penalties of oxidation are oxygen-containing acids and their salts.

All halogen compounds in positive oxidation states are

are strong oxidizing agents.

terrible degree of oxidation. Disproportionation is promoted by an alkaline environment.

Practical application of simple substances and oxygen compounds

The reduction of halogens is mainly due to their oxidizing effect.

Widest practical use find simple substances Cl2

and F2. The largest amount of chlorine and fluorine is consumed in industrial

organic synthesis: in the production of plastics, refrigerants, solvents,

pesticides, drugs. Significant amounts of chlorine and iodine are used to obtain metals and for their refining. Chlorine is also used

for bleaching cellulose, for disinfecting drinking water and in production

water of bleach and hydrochloric acid. Salts of oxoacids are used in the production of explosives.

Acids—hydrochloric and molten acids—are widely used in practice.

Fluorine and chlorine are among the twenty most common elements

there, there is significantly less bromine and iodine in nature. All halogens occur in nature in oxidation states(-1). Only iodine occurs in the form of the salt KIO3,

which is included as an impurity in Chilean saltpeter (KNO3).

Astatine is an artificially produced radioactive element (it does not exist in nature). The instability of At is reflected in the name, which comes from the Greek. "astatos" - "unstable". Astatine is a convenient emitter for radiotherapy of cancer tumors.

Simple substances

Simple substances of halogens are formed by diatomic molecules G2.

In simple substances, during the transition from F2 to I2 with an increase in the number of electrons

throne layers and an increase in the polarizability of atoms, there is an increase

intermolecular interaction, leading to a change in aggregate co-

standing under standard conditions.

Fluorine (under normal conditions) is a yellow gas, at –181o C it turns into

liquid state.

Chlorine is a yellow-green gas that turns into liquid at –34o C. With the color of ha-

The name Cl is associated with it, it comes from the Greek “chloros” - “yellow-

green". A sharp increase in the boiling point of Cl2 compared to F2,

indicates increased intermolecular interaction.

Bromine is a dark red, very volatile liquid, boils at 58.8o C.

the title of the element is associated with a sharp unpleasant smell gas and formed from

"bromos" - "smelly".

Iodine – dark purple crystals, with a faint “metallic”

lumps, which when heated easily sublimate, forming violet vapors;

The boiling point of iodine is 183 ° C. Its name comes from the color of iodine vapor -

"iodos" - "purple".

All simple substances have a pungent odor and are poisonous.

Inhalation of their vapors causes irritation of the mucous membranes and respiratory organs, and at high concentrations - suffocation. During the First World War, chlorine was used as a poisonous agent.

Fluorine gas and liquid bromine cause skin burns. Working with ha-

logens, precautions should be taken.

Since simple substances of halogens are formed by non-polar molecules

cools, they dissolve well in non-polar organic solvents:

alcohol, benzene, carbon tetrachloride, etc. Chlorine, bromine and iodine are sparingly soluble in water; their aqueous solutions are called chlorine, bromine and iodine water. Br2 dissolves better than others, bromine concentration in sat.

The solution reaches 0.2 mol/l, and chlorine – 0.1 mol/l.

Fluoride decomposes water:

2F2 + 2H2 O = O2 + 4HF

Halogens exhibit high oxidative activity and transition

into halide anions.

Г2 + 2e–  2Г–

Fluorine has especially high oxidative activity. Fluorine oxidizes precious metals(Au, Pt).

Pt + 3F2 = PtF6

It even interacts with some inert gases (krypton,

xenon and radon), for example,

Xe + 2F2 = XeF4

Many very stable compounds burn in an F2 atmosphere, e.g.

water, quartz (SiO2).

SiO2 + 2F2 = SiF4 + O2

In reactions with fluorine, even such strong oxidizing agents as nitrogen and sulfur

nic acid, act as reducing agents, while fluorine oxidizes the input

containing O(–2) in their composition.

2HNO3 + 4F2 = 2NF3 + 2HF + 3O2 H2 SO4 + 4F2 = SF6 + 2HF + 2O2

The high reactivity of F2 creates difficulties with the choice of con-

structural materials for working with it. Usually for these purposes we use

There are nickel and copper, which, when oxidized, form dense protective films of fluorides on their surface. The name F is due to its aggressive action.

I eat, it comes from the Greek. “fluoros” – “destructive”.

In the series F2, Cl2, Br2, I2, the oxidizing ability weakens due to an increase

increasing the size of atoms and decreasing electronegativity.

In aqueous solutions, the oxidative and reductive properties of matter

Substances are usually characterized using electrode potentials. The table shows standard electrode potentials (Eo, V) for reduction half-reactions

formation of halogens. For comparison, the Eo value for ki-

carbon is the most common oxidizing agent.

Standard electrode potentials for simple halogen substances

Reduced oxidative activity

As can be seen from the table, F2 is a much stronger oxidizing agent,

than O2, therefore F2 does not exist in aqueous solutions , it oxidizes water,

recovering to F–. Judging by the Eо value, the oxidizing ability of Cl2

also higher than that of O2. Indeed, during long-term storage of chlorine water, it decomposes with the release of oxygen and the formation of HCl. But the reaction is slow (the Cl2 molecule is noticeably stronger than the F2 molecule and

activation energy for reactions with chlorine is higher), dispro-

portioning:

Cl2 + H2 O HCl + HOCl

In water it does not reach the end (K = 3.9 . 10–4), therefore Cl2 exists in aqueous solutions. Br2 and I2 are characterized by even greater stability in water.

Disproportionation is a very characteristic oxidative

reduction reaction for halogens. Disproportionation of the amplification

pours in an alkaline environment.

Disproportionation of Cl2 in alkali leads to the formation of anions

Cl– and ClO–. The disproportionation constant is 7.5. 1015.

Cl2 + 2NaOH = NaCl + NaClO + H2O

When iodine is disproportioned in alkali, I– and IO3– are formed. Ana-

Logically, Br2 disproportionates iodine. Product change is disproportionate

nation is due to the fact that the anions GO– and GO2– in Br and I are unstable.

The chlorine disproportionation reaction is used in industrial

ability to obtain a strong and fast-acting hypochlorite oxidizer,

bleaching lime, bertholet salt.

3Cl2 + 6 KOH = 5KCl + KClO3 + 3H2 O

Interaction of halogens with metals

Halogens react vigorously with many metals, for example:

Mg + Cl2 = MgCl2 Ti + 2I2  TiI4

Na + halides, in which the metal has a low oxidation state (+1, +2),

- These are salt-like compounds with predominantly ionic bonds. How to

lo, ionic halides are solids with high temperature floating

Metal halides, in which the metal has high degree oxidation

tions are compounds with predominantly covalent bonds.

Many of them are gases, liquids or fusible solids under normal conditions. For example, WF6 is a gas, MoF6 is a liquid,

TiCl4 is liquid.

Interaction of halogens with non-metals

Halogens interact directly with many nonmetals:

hydrogen, phosphorus, sulfur, etc. For example:

H2 + Cl2 = 2HCl 2P + 3Br2 = 2PBr3 S + 3F2 = SF6

The bonding in nonmetal halides is predominantly covalent.

Typically these compounds have low melting and boiling points.

When passing from fluorine to iodine, the covalent nature of the halides increases.

The covalent halides of typical nonmetals are acidic compounds; when interacting with water, they hydrolyze to form acids. For example:

PBr3 + 3H2 O = 3HBr + H3 PO3

PI3 + 3H2 O = 3HI + H3 PO3

PCl5 + 4H2 O = 5HCl + H3 POinterga-

leads. In these compounds, the lighter and more electronegative halogen is in the (–1) oxidation state, and the heavier one is in the positive state.

oxidation penalties.

Due to the direct interaction of halogens upon heating, the following are obtained: ClF, BrF, BrCl, ICl. There are also more complex interhalides:

ClF3, BrF3, BrF5, IF5, IF7, ICl3.

All interhalides under normal conditions – liquid substances With low temperatures boiling. Interhalides have a high oxidative activity

activity. For example, such chemically stable substances as SiO2, Al2 O3, MgO, etc. burn in ClF3 vapors.

2Al2 O3 + 4ClF3 = 4 AlF3 + 3O2 + 2Cl2

Fluoride ClF 3 is an aggressive fluorinating reagent that acts quickly

yard F2. It is used in organic syntheses and to obtain protective films on the surface of nickel equipment for working with fluorine.

In water, interhalides hydrolyze to form acids. For example,

ClF5 + 3H2 O = HClO3 + 5HF

Halogens in nature. Obtaining simple substances

In industry, halogens are obtained from their natural compounds. All

processes for obtaining free halogens are based on the oxidation of halogen

Nid ions.

2Г –  Г2 + 2e–

A significant amount of halogens is found in natural waters in the form of anions: Cl–, F–, Br–, I–. IN sea ​​water may contain up to 2.5% NaCl.

Bromine and iodine are obtained from oil well water and sea water.

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's 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; It has bad smell; 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

Of 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

Dark crystalline substance purple with a metallic sheen.
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°C; 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, t°melt.= 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.

Nuclear Research Dubna. Fluoride is a poisonous and reactive pale yellow gas. Chlorine is a heavy, poisonous, light green gas with an unpleasant chlorine odor. Bromine, a toxic red-brown liquid that can damage the olfactory nerve, is contained in ampoules, because. has the property of volatility. Iodine is an easily sublimated poisonous violet-black crystal. Astatine is a radioactive blue-black crystal, the period of the longest isotope is 8.1 hours. All halogens react with almost all simple substances, with the exception of a few. They are energetic oxidizing agents, so they can only be found in the form of compounds. The chemical activity of halogens decreases with increasing atomic number. Halogens have high oxidation activity, which decreases when moving from fluorine to iodine. The most active is fluorine, which reacts with all metals. Many of the metals in the atmosphere of this element spontaneously ignite and release a large number of warmth. Without heating, fluorine can react with many non-metals, and all reactions are . Fluorine reacts with noble () gases upon irradiation. Free chlorine, despite the fact that its activity is less than that of fluorine, is also very reactive. Chlorine can react with all simple substances except oxygen, nitrogen and inert gases. This element also reacts with many complex substances, substituting and joining with hydrocarbons. When heated, chlorine displaces bromine, as well as iodine, from their compounds with metals or hydrogen. The chemical activity is also quite high, although less than that of fluorine or chlorine, so bromine is mainly used in the liquid state and its initial concentrations, other conditions being equal more than chlorine. This element, similarly, dissolves in water and, partially reacting with it, creates “bromine water.” Iodine differs in chemical activity from other halogens. It cannot react with most non-metals, and only reacts with metals when heated and very slowly. The reaction is highly reversible and endothermic. Iodine is insoluble in water and even when heated cannot oxidize it, so “iodine water” does not exist. Iodine can dissolve in solutions of iodides to form complex anions. Astatine reacts with hydrogen and metals. The chemical activity of halogens decreases successively from fluorine to iodine. Each halogen displaces the next one from its compounds with metals or hydrogen, i.e. each halogen in the form of a simple substance can oxidize the halogen ion of any of the following halogens.

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 complete octet. This characteristic makes them more reactive than other groups of nonmetals.

general characteristics

Halogens form diatomic molecules (type X2, 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. The chemical properties of halogens allow them to easily combine with most elements, which is why 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 as 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 aqueous solution. Halogens form single bonds with carbon or nitrogen in organic compounds, where their oxidation state (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-, iodine- 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 for rock salt, a naturally occurring mineral composed 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 is an element with atomic number 9 and is designated by the symbol F. Elemental fluorine was first discovered in 1886 by isolating it from hydrofluoric acid. In the free state, fluorine exists as a diatomic molecule (F2) and is the most abundant halogen in earth's crust. Fluorine is the most electronegative element on the periodic table. At room temperature 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 Cl2. 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 Cl2 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 seawater. Chlorine also has two isotopes: 35Cl and 37Cl. Sodium chloride is the most common compound of all the chlorides.

Bromine – chemical element with atomic number 35 and symbol Br. It was first discovered in 1826. In its elemental form, bromine is a diatomic molecule Br2. 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: 79Br and 81Br. Bromine occurs as bromide salts dissolved in seawater. Behind last years The world's bromide production has increased significantly 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 as a diatomic molecule, I2. At room temperature it is a purple solid. Iodine has one stable isotope - 127I. 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. Differs in metallic properties.

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

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

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

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

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 halogens decreases with increasing period (At

Inorganic chemistry. 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 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) + H2O (l) → X- (aq) + H3O+ (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:

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 valency. 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 (HBrO4), 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 (X2), 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 1s2 2s2 2p5. 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.