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.

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Chemistry of elements of the VIIA subgroup

general characteristics

							Manganese
							Technetium

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.

							As can be seen from the distribution
							valence electrons
according to orbitals of atoms						only one electron missing

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.”

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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

		Relative-		Affinity
		electric
		negative	ionization,
		ness (according to
		Polling)
						increase in number
						electronic layers;
						increase in size
						reduction of electrical
						triple negativity

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 is a general pattern for all non-metals.

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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.

The simplest substances, Cl2, find the widest practical application.

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.

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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 their oxidation state(-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 name of the element is associated with the sharp unpleasant odor of gas and is derived from

"bromos" - "smelly".

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

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

with rapid cooling

vapors up to 114o C

liquid is formed. Temperature

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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 noble 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

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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

Eo, B, for reaction

O2 + 4e– + 4H+  2H2 O

Eo, V

for electrode

2Г– +2е – = Г2

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

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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

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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 a high melting point

Metal halides in which the metal has a high degree of 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 PO4

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The first two reactions are used to produce bromine and hydrogen iodide.

noic acid.

Interhalides. Halogens, combining with each other, form interg-

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 are liquid substances with low boiling points. 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–. Sea water can contain up to 2.5% NaCl.

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

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The hydrogen atom has the electronic formula of the outer (and only) electron level 1 s 1 . On the one hand, in terms of the presence of one electron on the outer electronic level, the hydrogen atom is similar to alkali metal atoms. However, just like halogens, it only needs one electron to fill the outer electronic level, since the first electronic level can contain no more than 2 electrons. It turns out that hydrogen can be placed simultaneously in both the first and the penultimate (seventh) group of the periodic table, which is sometimes done in various versions of the periodic system:

From the point of view of the properties of hydrogen as a simple substance, it still has more in common with halogens. Hydrogen, like halogens, is a non-metal and forms diatomic molecules (H 2) like them.

Under normal conditions, hydrogen is a gaseous, low-active substance. The low activity of hydrogen is explained by the high strength of the bonds between the hydrogen atoms in the molecule, the breaking of which requires either strong heating, or the use of catalysts, or both.

Interaction of hydrogen with simple substances

with metals

Of the metals, hydrogen reacts only with alkali and alkaline earth metals! Alkali metals include metals of the main subgroup of group I (Li, Na, K, Rb, Cs, Fr), and alkaline earth metals include metals of the main subgroup of group II, except beryllium and magnesium (Ca, Sr, Ba, Ra)

When interacting with active metals, hydrogen exhibits oxidizing properties, i.e. lowers its oxidation state. In this case, hydrides of alkali and alkaline earth metals are formed, which have an ionic structure. The reaction occurs when heated:

It should be noted that interaction with active metals is the only case when molecular hydrogen H2 is an oxidizing agent.

with non-metals

Of the non-metals, hydrogen reacts only with carbon, nitrogen, oxygen, sulfur, selenium and halogens!

Carbon should be understood as graphite or amorphous carbon, since diamond is an extremely inert allotropic modification of carbon.

When interacting with non-metals, hydrogen can only perform the function of a reducing agent, that is, only increase its oxidation state:

Interaction of hydrogen with complex substances

with metal oxides

Hydrogen does not react with metal oxides that are in the activity series of metals up to aluminum (inclusive), however, it is capable of reducing many metal oxides to the right of aluminum when heated:

with non-metal oxides

Of the non-metal oxides, hydrogen reacts when heated with the oxides of nitrogen, halogens and carbon. Of all the interactions of hydrogen with non-metal oxides, especially noteworthy is its reaction with carbon monoxide CO.

The mixture of CO and H2 even has its own name - “synthesis gas”, since, depending on the conditions, such popular industrial products as methanol, formaldehyde and even synthetic hydrocarbons can be obtained from it:

with acids

Hydrogen does not react with inorganic acids!

Of organic acids, hydrogen reacts only with unsaturated acids, as well as with acids containing functional groups capable of reduction with hydrogen, in particular aldehyde, keto or nitro groups.

with salts

In the case of aqueous solutions of salts, their interaction with hydrogen does not occur. However, when hydrogen is passed over solid salts of some metals of medium and low activity, their partial or complete reduction is possible, for example:

Chemical properties of halogens

Halogens are the chemical elements of group VIIA (F, Cl, Br, I, At), as well as the simple substances they form. Here and further in the text, unless otherwise stated, halogens will be understood as simple substances.

All halogens have a molecular structure, which determines the low melting and boiling points of these substances. Halogen molecules are diatomic, i.e. their formula can be written in general form as Hal 2.

It should be noted such a specific physical property of iodine as its ability to sublimation or, in other words, sublimation. Sublimation, is a phenomenon in which a substance in a solid state does not melt when heated, but, bypassing the liquid phase, immediately passes into the gaseous state.

The electronic structure of the external energy level of an atom of any halogen has the form ns 2 np 5, where n is the number of the periodic table period in which the halogen is located. As you can see, the halogen atoms only need one electron to reach the eight-electron outer shell. From this it is logical to assume the predominantly oxidizing properties of free halogens, which is confirmed in practice. As is known, the electronegativity of nonmetals decreases when moving down a subgroup, and therefore the activity of halogens decreases in the series:

F 2 > Cl 2 > Br 2 > I 2

Interaction of halogens with simple substances

All halogens are highly reactive substances and react with most simple substances. However, it should be noted that fluorine, due to its extremely high reactivity, can react even with those simple substances with which other halogens cannot react. Such simple substances include oxygen, carbon (diamond), nitrogen, platinum, gold and some noble gases (xenon and krypton). Those. actually, fluorine does not react only with some noble gases.

The remaining halogens, i.e. chlorine, bromine and iodine are also active substances, but less active than fluorine. They react with almost all simple substances except oxygen, nitrogen, carbon in the form of diamond, platinum, gold and noble gases.

Interaction of halogens with non-metals

hydrogen

When all halogens interact with hydrogen, they form hydrogen halides with the general formula HHal. In this case, the reaction of fluorine with hydrogen begins spontaneously even in the dark and proceeds with an explosion in accordance with the equation:

The reaction of chlorine with hydrogen can be initiated by intense ultraviolet irradiation or heat. Also proceeds with explosion:

Bromine and iodine react with hydrogen only when heated, and at the same time, the reaction with iodine is reversible:

phosphorus

The interaction of fluorine with phosphorus leads to the oxidation of phosphorus to the highest oxidation state (+5). In this case, phosphorus pentafluoride is formed:

When chlorine and bromine interact with phosphorus, it is possible to obtain phosphorus halides both in the oxidation state + 3 and in the oxidation state +5, which depends on the proportions of the reacting substances:

Moreover, in the case of white phosphorus in an atmosphere of fluorine, chlorine or liquid bromine, the reaction begins spontaneously.

The interaction of phosphorus with iodine can lead to the formation of only phosphorus triodide due to its significantly lower oxidizing ability than that of other halogens:

gray

Fluorine oxidizes sulfur to the highest oxidation state +6, forming sulfur hexafluoride:

Chlorine and bromine react with sulfur, forming compounds containing sulfur in the oxidation states +1 and +2, which are extremely unusual for it. These interactions are very specific, and to pass the Unified State Exam in chemistry, the ability to write equations for these interactions is not necessary. Therefore, the following three equations are given rather for reference:

Interaction of halogens with metals

As mentioned above, fluorine is capable of reacting with all metals, even such inactive ones as platinum and gold:

The remaining halogens react with all metals except platinum and gold:

Reactions of halogens with complex substances

Substitution reactions with halogens

More active halogens, i.e. the chemical elements of which are located higher in the periodic table are capable of displacing less active halogens from the hydrohalic acids and metal halides they form:

Similarly, bromine and iodine displace sulfur from solutions of sulfides and or hydrogen sulfide:

Chlorine is a stronger oxidizing agent and oxidizes hydrogen sulfide in its aqueous solution not to sulfur, but to sulfuric acid:

Reaction of halogens with water

Water burns in fluorine with a blue flame in accordance with the reaction equation:

Bromine and chlorine react differently with water than fluorine. If fluorine acted as an oxidizing agent, then chlorine and bromine are disproportionate in water, forming a mixture of acids. In this case, the reactions are reversible:

The interaction of iodine with water occurs to such an insignificant degree that it can be neglected and it can be assumed that the reaction does not occur at all.

Interaction of halogens with alkali solutions

Fluorine, when interacting with an aqueous alkali solution, again acts as an oxidizing agent:

The ability to write this equation is not required to pass the Unified State Exam. It is enough to know the fact about the possibility of such an interaction and the oxidative role of fluorine in this reaction.

Unlike fluorine, other halogens in alkali solutions are disproportionate, that is, they simultaneously increase and decrease their oxidation state. Moreover, in the case of chlorine and bromine, depending on the temperature, flow in two different directions is possible. In particular, in the cold the reactions proceed as follows:

and when heated:

Iodine reacts with alkalis exclusively according to the second option, i.e. with the formation of iodate, because hypoiodite is not stable not only when heated, but also at ordinary temperatures and even in the cold.

1. General characteristics of halogens . Atomic structure and oxidation states of halogens in compounds. The nature of changes in atomic radii, ionization energies, electron affinities and electronegativity in the series F - At. The nature of chemical bonds of halogens with metals and non-metals. Stability of higher valence states of halogens. Features of fluorine.

1. With. 367-371; 2. With. 338-347; 3. With. 415-416; 4. With. 270-271; 7. With. 340-345.

2. Molecular structure and physical properties of simple halogen substances . The nature of chemical bonds in halogen molecules. Physical properties of halogens: state of aggregation, melting and boiling points in the fluorine - astatine series, solubility in water and organic solvents.

1. With. 370-372; 2. With. 340-347; 3. With. 415-416; 4. With. 271-287; 8. With. 367-370.

3. Chemical properties of halogens . Reasons for the high chemical activity of halogens and its change by group. Relation to water, alkali solutions, metals and non-metals. The influence of temperature on the composition of halogen disproportionation products in alkali solutions. Features of fluorine chemistry. Natural halogen compounds. Principles of industrial and laboratory methods for producing halogens. Use of halogens. Physiological and pharmacological effects of halogens and their compounds on living organisms. Toxicity of halogens and precautions when working with them.

1. With. 372-374, p. 387-388; 2. With. 342-347; 3. With. 416-419; 4. With. 276-287; 7. pp.340-345, p. 355; 8. With. 380-382.

Simple substances, halogens, unlike hydrogen, are very active. They are most characterized by oxidizing properties, which gradually weaken in the series F 2 – At 2. The most active of the halogens is fluorine: even water and sand spontaneously ignite in its atmosphere! Halogens react vigorously with most metals, non-metals, and complex substances.

4. Production and use of halogens .

1. With. 371-372; 2. With. 345-347; 3. With. 416-419; 4. With. 275-287; 7. pp.340-345; 8. With. 380-382.

All methods for producing halogens are based on the oxidation reactions of halide anions with various oxidizing agents: 2Gal -1 -2e - = Gal

In industry, halogens are obtained by electrolysis of melts (F 2 and Cl 2) or aqueous solutions (Cl 2) of halides; displacement of less active halogens by more active ones from the corresponding halides (I 2 - bromine; I 2 or Br 2 - chlorine)

Halogens in the laboratory are obtained by oxidation of hydrogen halides (HCl, HBr) in solutions with strong oxidizing agents (KMnO 4, K 2 Cr 2 O 7, PbO 2, MnO 2, KClO 3); oxidation of halides (NaBr, KI) with the indicated oxidizing agents in an acidic environment (H 2 SO 4).

Binary halogen compounds

1. Hydrogen compounds (hydrogen halides) . The nature of chemical bonds in molecules. Polarity of molecules. Physical properties, state of aggregation, solubility in water. The nature of changes in melting and boiling temperatures in the HF – HI series. Association of hydrogen fluoride molecules. Thermal stability of hydrogen halides. Reactivity. Acid properties, features of hydrofluoric acid. Restorative properties. General principles for the production of hydrogen halides: synthesis from simple substances and from halides. Hydrogen chloride and hydrochloric acid. Physical and chemical properties. Methods of obtaining. The use of hydrochloric acid. The role of hydrochloric acid and chlorides in life processes. Halides.

1. With. 375-382; 2. With. 347-353; 3. With. 419-420; 4. With. 272-275, p. 289-292; 7. pp.354-545; 8. With. 370-373, p. 374-375.

2 . Compounds of halogens with oxygen.

1. With. 377-380; 2. With. 353-359; 3. With. 420-423; 4. With. 292-296; 7. pp.350-354; 8. With. 375-376, p. 379.

3. Compounds with other non-metals.

1. With. 375-381; 2. With. 342-345; 4. With. 292-296; 7. p.350-355.

4 . Connections to metals .

2. With. 342; 4. With. 292-296; 7. p.350-355.

Multi-element halogen compounds

1. Oxygen-containing chlorine acids and their salts. Hypochlorous, chlorous, perchloric and perchloric acids. Changes in acid properties, stability and oxidizing properties in the series HClO – HClO 4 . Principles for obtaining these acids. Hypochlorites, chlorites, chlorates and perchlorates. Thermal stability and oxidative properties. General principles for obtaining salts. Use of salts. Bleaching powder. Berthollet's salt. Ammonium perchlorate.

1. With. 382-387; 2. With. 353-359; 3. With. 423; 4. With. 292-296; 7. pp.350-354; 8. With. 375-378.

2 . Oxygen-containing acids of bromine and iodine and their salts .

1. With. 382-387; 2. With. 353-359; 3. With. 423; 4. With. 292-296; 7. pp.350-354; 8. With. 379-380.

3 . Application of halogens and their most important compounds

1. With. 387-388; 2. With. 345-347; 3. With. 419-423; 4. With. 272-296; 8. With. 380-382.

4 . Biological role of halogen compounds

1. With. 387-388; 2. With. 340-347; 3. With. 419-423; 4. With. 272-296; 8. With. 380-382.

Relationshipthe most important chlorine compounds:

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 gain 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, under normal conditions, fluorine is a gas that is difficult to liquefy, chlorine is also a gas, but 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 at the outer level, they can also exhibit other odd oxidation states from to due to partial or complete pairing 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.

Halogens– Group VII elements – fluorine, chlorine, bromine, iodine, astatine (astatine has been little studied due to its radioactivity). Halogens are distinct non-metals. Only iodine in rare cases exhibits some properties similar to metals.

In the unexcited state, the halogen atoms have a common electronic configuration: ns2np5. This means that halogens have 7 valence electrons, except for fluorine.

Physical properties of halogens: F2 – colorless, difficult to liquefy gas; Cl2 is a yellow-green, easily liquefied gas with a pungent suffocating odor; Br2 – red-brown liquid; I2 is a violet crystalline substance.

Aqueous solutions of hydrogen halides form acids. HF – hydrogen fluoride (fluoride); HCl – hydrochloric (salt); НBr—hydrogen bromide; HI – hydrogen iodide. The strength of acids decreases from top to bottom. Hydrofluoric acid is the weakest in the series of halogenated acids, and hydroiodic acid is the strongest. This is explained by the fact that the binding energy of Hg decreases from above. The strength of the NG molecule decreases in the same direction, which is associated with an increase in the internuclear distance. The solubility of slightly soluble salts in water also decreases:

From left to right, the solubility of halides decreases. AgF is highly soluble in water. All halogens in the free state are oxidizing agents. Their strength as oxidizing agents decreases from fluorine to iodine. In crystalline, liquid and gaseous states, all halogens exist in the form of individual molecules. Atomic radii increase in the same direction, which leads to an increase in melting and boiling points. Fluorine dissociates into atoms better than iodine. Electrode potentials decrease when moving down the halogen subgroup. Fluorine has the highest electrode potential. Fluorine is the strongest oxidizing agent. Any higher free halogen will displace the lower one, which is in the state of a negative singly charged ion in solution.

20. Chlorine. Hydrogen chloride and hydrochloric acid

Chlorine (Cl) – stands in the 3rd period, in group VII of the main subgroup of the periodic system, serial number 17, atomic mass 35.453; refers to halogens.

Physical properties: yellow-green gas with a pungent odor. Density 3.214 g/l; melting point -101 °C; boiling point -33.97 °C, At ordinary temperature it easily liquefies under a pressure of 0.6 MPa. Dissolving in water, it forms yellowish chlorine water. It is highly soluble in organic solvents, especially hexane (C6H14), and carbon tetrachloride.

Chemical properties of chlorine: electronic configuration: 1s22s22p63s22p5. There are 7 electrons in the outer level. To complete the level, you need 1 electron, which chlorine accepts, exhibiting an oxidation state of -1. There are also positive oxidation states of chlorine up to + 7. The following chlorine oxides are known: Cl2O, ClO2, Cl2O6 and Cl2O7. They are all unstable. Chlorine is a strong oxidizing agent. It reacts directly with metals and non-metals:

Reacts with hydrogen. Under normal conditions, the reaction proceeds slowly, with strong heating or lighting - with an explosion, according to a chain mechanism:

Chlorine interacts with alkali solutions, forming salts - hypochlorites and chlorides:

When chlorine is passed into an alkali solution, a mixture of chloride and hypochlorite solutions is formed:

Chlorine is a reducing agent: Cl2 + 3F2 = 2ClF3.

Interaction with water:

Chlorine does not react directly with carbon, nitrogen and oxygen.

Receipt: 2NaCl + F2 = 2NaF + Cl2.

Electrolysis: 2NaCl + 2H2O = Cl2 + H2 + 2NaOH.

Finding in nature: contained in the following minerals: halite (rock salt), sylvite, bischofite; sea ​​water contains chlorides of sodium, potassium, magnesium and other elements.

Hydrogen chloride HCl. Physical properties: a colorless gas, heavier than air, highly soluble in water to form hydrochloric acid.

Receipt: in the laboratory:

In industry: hydrogen is burned in a stream of chlorine. Next, hydrogen chloride is dissolved in water to form hydrochloric acid (see above).

Chemical properties: hydrochloric acid is strong, monobasic, interacts with metals in the voltage series up to hydrogen: Zn + 2HCl = ZnCl2 + H2.

As a reducing agent it reacts with oxides and hydroxides of many metals.