Cl 2 at vol. T - yellow-green gas with a sharp suffocating odor, heavier than air - 2.5 times, slightly soluble in water (~ 6.5 g / l); X. R. in nonpolar organic solvents. It is found free only in volcanic gases.
How to get
Based on the process of oxidation of anions Cl -
2Cl - - 2e - = Cl 2 0
Industrial
Electrolysis of aqueous solutions of chlorides, more often - NaCl:
2NaCl + 2H 2 O \u003d Cl 2 + 2NaOH + H 2
Laboratory
Oxidation conc. HCI various oxidizing agents:
4HCI + MnO 2 \u003d Cl 2 + MpCl 2 + 2H 2 O
16HCl + 2KMnO 4 \u003d 5Cl 2 + 2MnCl 2 + 2KCl + 8H 2 O
6HCl + KClO 3 \u003d ZCl 2 + KCl + 3H 2 O
14HCl + K 2 Cr 2 O 7 \u003d 3Cl 2 + 2CrCl 3 + 2KCl + 7H 2 O
Chemical properties
Chlorine is a very strong oxidizing agent. Oxidizes metals, non-metals and complex substances, while turning into very stable anions Cl -:
Cl 2 0 + 2e - \u003d 2Cl -
Reactions with metals
Active metals in an atmosphere of dry chlorine gas ignite and burn; in this case, metal chlorides are formed.
Cl 2 + 2Na = 2NaCl
3Cl 2 + 2Fe = 2FeCl 3
Inactive metals are more easily oxidized by wet chlorine or its aqueous solutions:
Cl 2 + Cu \u003d CuCl 2
3Cl 2 + 2Au = 2AuCl 3
Reactions with non-metals
Chlorine does not directly interact only with O 2, N 2, C. Reactions proceed with other non-metals under various conditions.
Non-metal halides are formed. The most important is the reaction of interaction with hydrogen.
Cl 2 + H 2 \u003d 2HC1
Cl 2 + 2S (melt) = S 2 Cl 2
ЗCl 2 + 2Р = 2РCl 3 (or РCl 5 - in excess of Cl 2)
2Cl 2 + Si = SiCl 4
3Cl 2 + I 2 \u003d 2ICl 3
Displacement of free non-metals (Br 2, I 2, N 2, S) from their compounds
Cl 2 + 2KBr = Br 2 + 2KCl
Cl 2 + 2KI \u003d I 2 + 2KCl
Cl 2 + 2HI \u003d I 2 + 2HCl
Cl 2 + H 2 S \u003d S + 2HCl
ZCl 2 + 2NH 3 \u003d N 2 + 6HCl
Disproportionation of chlorine in water and aqueous solutions of alkalis
As a result of self-oxidation-self-healing, some chlorine atoms are converted into Cl - anions, while others in a positive oxidation state are part of the ClO - or ClO 3 - anions.
Cl 2 + H 2 O \u003d HCl + HClO hypochlorous to-ta
Cl 2 + 2KOH \u003d KCl + KClO + H 2 O
3Cl 2 + 6KOH = 5KCl + KClO 3 + 3H 2 O
3Cl 2 + 2Ca (OH) 2 \u003d CaCl 2 + Ca (ClO) 2 + 2H 2 O
These reactions are important because they lead to the production of oxygen compounds of chlorine:
KClO 3 and Ca (ClO) 2 - hypochlorites; KClO 3 - potassium chlorate (bertolet salt).
Interaction of chlorine with organic substances
a) substitution of hydrogen atoms in OB molecules
b) attachment of Cl 2 molecules at the point of breaking of multiple carbon-carbon bonds
H 2 C \u003d CH 2 + Cl 2 → ClH 2 C-CH 2 Cl 1,2-dichloroethane
HC≡CH + 2Cl 2 → Cl 2 HC-CHCl 2 1,1,2,2-tetrachloroethane
Hydrogen chloride and hydrochloric acid
Hydrogen chloride gas
Physical and chemical properties
HCl is hydrogen chloride. At rev. T - colorless. gas with a pungent odor, liquefies quite easily (mp. -114°С, bp. -85°С). Anhydrous HCl, both in gaseous and liquid states, is non-conductive, chemically inert with respect to metals, metal oxides and hydroxides, and also to many other substances. This means that in the absence of water, hydrogen chloride does not exhibit acidic properties. Only at very high temperatures does gaseous HCl react with metals, even such inactive ones as Cu and Ag.
The reducing properties of the chloride anion in HCl also manifest themselves to a small extent: it is oxidized by fluorine at vol. T, and also at high T (600°C) in the presence of catalysts, it reversibly reacts with oxygen:
2HCl + F 2 \u003d Cl 2 + 2HF
4HCl + O 2 \u003d 2Cl 2 + 2H 2 O
Gaseous HCl is widely used in organic synthesis (hydrochlorination reactions).
How to get
1. Synthesis from simple substances:
H 2 + Cl 2 \u003d 2HCl
2. Formed as a by-product during hydrocarbon chlorination:
R-H + Cl 2 = R-Cl + HCl
3. In the laboratory, they receive the action of conc. H 2 SO 4 for chlorides:
H 2 SO 4 (conc.) + NaCl \u003d 2HCl + NaHSO 4 (with low heating)
H 2 SO 4 (conc.) + 2NaCl \u003d 2HCl + Na 2 SO 4 (with very strong heating)
An aqueous solution of HCl is a strong acid (hydrochloric, or hydrochloric)
HCl is very soluble in water: at vol. T in 1 l of H 2 O dissolves ~ 450 l of gas (dissolution is accompanied by the release of a significant amount of heat). A saturated solution has a mass fraction of HCl equal to 36-37%. This solution has a very pungent, suffocating odor.
HCl molecules in water almost completely decompose into ions, i.e., an aqueous solution of HCl is a strong acid.
Chemical properties of hydrochloric acid
1. HCl dissolved in water exhibits all the general properties of acids due to the presence of H + ions
HCl → H + + Cl -
Interaction:
a) with metals (up to H):
2HCl 2 + Zn \u003d ZnCl 2 + H 2
b) with basic and amphoteric oxides:
2HCl + CuO \u003d CuCl 2 + H 2 O
6HCl + Al 2 O 3 \u003d 2AlCl 3 + ZN 2 O
c) with bases and amphoteric hydroxides:
2HCl + Ca (OH) 2 \u003d CaCl 2 + 2H 2 O
3HCl + Al(OH) 3 \u003d AlCl 3 + ZN 2 O
d) with salts of weaker acids:
2HCl + CaCO 3 \u003d CaCl 2 + CO 2 + H 3 O
HCl + C 6 H 5 ONa \u003d C 6 H 5 OH + NaCl
e) with ammonia:
HCl + NH 3 \u003d NH 4 Cl
Reactions with strong oxidizing agents F 2 , MnO 2 , KMnO 4 , KClO 3 , K 2 Cr 2 O 7 . Anion Cl - is oxidized to free halogen:
2Cl - - 2e - = Cl 2 0
For reaction equations, see "Getting Chlorine". OVR between hydrochloric and nitric acids is of particular importance:
Reactions with organic compounds
Interaction:
a) with amines (as organic bases)
R-NH 2 + HCl → + Cl -
b) with amino acids (as amphoteric compounds)
Oxides and oxoacids of chlorine
Acid oxides
acids
salt
Chemical properties
1. All oxoacids of chlorine and their salts are strong oxidizers.
2. Almost all compounds decompose when heated due to intramolecular oxidation-reduction or disproportionation.
Bleaching powder
Chlorine (whitewash) lime - a mixture of hypochlorite and calcium chloride, has a bleaching and disinfecting effect. Sometimes it is considered as an example of a mixed salt, which simultaneously contains anions of two acids:
Javel water
Aqueous solution of chloride and potassium hapochlorite KCl + KClO + H 2 O
15.1. General characteristics of halogens and chalcogens
Halogens ("giving birth to salts") are elements of group VIIA. These include fluorine, chlorine, bromine and iodine. This group also includes unstable, and therefore not naturally occurring, astatine. Sometimes hydrogen is also included in this group.
Chalcogens ("copper-producing") are elements of the VIA group. These include oxygen, sulfur, selenium, tellurium and the almost non-natural polonium.
Of the eight naturally occurring atoms elements of these two groups, the most common oxygen atoms ( w= 49.5%), followed by chlorine atoms in abundance ( w= 0.19%), then - sulfur ( w= 0.048%), then - fluorine ( w= 0.028%). The atoms of other elements are hundreds and thousands of times smaller. You already studied oxygen in the eighth grade (chap. 10), of the other elements, chlorine and sulfur are the most important - you will get to know them in this chapter.
The orbital radii of the atoms of halogens and chalcogens are small and only for the fourth atoms of each group approach one angstrom. This leads to the fact that all these elements are elements that form non-metals and only tellurium and iodine show some signs of amphoterism.
The general valence electronic formula of halogens is ns 2 np 5 , and chalcogens - ns 2 np four . The small size of atoms does not allow them to donate electrons; on the contrary, the atoms of these elements tend to accept them, forming singly charged (for halogens) and doubly charged (for chalcogens) anions. Connecting with small atoms, the atoms of these elements form covalent bonds. Seven valence electrons enable halogen atoms (except fluorine) to form up to seven covalent bonds, and six valence electrons of chalcogen atoms - up to six covalent bonds.
In compounds of fluorine, the most electronegative element, only one oxidation state is possible, namely -I. Oxygen, as you know, has a maximum oxidation state of +II. For atoms of other elements, the highest oxidation state is equal to the group number.
Simple substances of elements of group VIIA are of the same type in structure. They are made up of diatomic molecules. Under normal conditions, fluorine and chlorine are gases, bromine is a liquid, and iodine is a solid. According to their chemical properties, these substances are strong oxidizing agents. Due to the growth in the size of atoms with an increase in the atomic number, their oxidative activity decreases.
Of the simple substances of group VIA elements, under normal conditions, only oxygen and ozone, consisting of diatomic and triatomic molecules, are gaseous, respectively; the rest are solids. Sulfur consists of eight-atomic cyclic molecules S 8 , selenium and tellurium from polymer molecules Se n and Te n. In terms of their oxidizing activity, chalcogens are inferior to halogens: only oxygen is a strong oxidizing agent of them, while the rest exhibit oxidizing properties to a much lesser extent.
Compound hydrogen compounds halogens (NE) fully complies with the general rule, and chalcogens, in addition to the usual hydrogen compounds of the composition H 2 E, can also form more complex hydrogen compounds of the composition H 2 E n chain structure. In aqueous solutions, both hydrogen halides and other hydrogen chalcogens exhibit acidic properties. Their molecules are acid particles. Of these, only HCl, HBr and HI are strong acids.
For halogens formation oxides uncharacteristic, most of them are unstable, however, higher oxides of the composition E 2 O 7 are known for all halogens (except for fluorine, whose oxygen compounds are not oxides). All halogen oxides are molecular substances, chemically they are acidic oxides.
In accordance with their valence capabilities, chalcogens form two series of oxides: EO 2 and EO 3. All these oxides are acidic.
Hydroxides of halogens and chalcogens are oxoacids.
Make abbreviated electronic formulas and energy diagrams of atoms of elements of VIA and VIIA groups. Indicate the outer and valence electrons.
Chlorine is the most common and therefore the most important of the halogens.
In the earth's crust, chlorine is found in the composition of minerals: halite (rock salt) NaCl, sylvin KCl, carnallite KCl MgCl 2 6H 2 O and many others. The main industrial production method is the electrolysis of sodium or potassium chlorides.
The simple substance chlorine is a greenish gas with a pungent, suffocating odor. At -101 °C, it condenses into a yellow-green liquid. Chlorine is very poisonous, during the First World War they even tried to use it as a chemical warfare agent.
Chlorine is one of the strongest oxidizing agents. It reacts with most simple substances (exception: noble gases, oxygen, nitrogen, graphite, diamond and some others). As a result, halides are formed:
Cl 2 + H 2 \u003d 2HCl (when heated or in the light);
5Cl 2 + 2P = 2PCl 5 (when burned in excess chlorine);
Cl 2 + 2Na = 2NaCl (at room temperature);
3Cl 2 + 2Sb = 2SbCl 3 (at room temperature);
3Cl 2 + 2Fe \u003d 2FeCl 3 (when heated).
In addition, chlorine can also oxidize many complex substances, for example:
Cl 2 + 2HBr = Br 2 + 2HCl (in the gas phase and in solution);
Cl 2 + 2HI \u003d I 2 + 2HCl (in the gas phase and in solution);
Cl 2 + H 2 S = 2HCl + S (in solution);
Cl 2 + 2KBr = Br 2 + 2KCl (in solution);
Cl 2 + 3H 2 O 2 = 2HCl + 2H 2 O + O 2 (in concentrated solution);
Cl 2 + CO \u003d CCl 2 O (in the gas phase);
Cl 2 + C 2 H 4 \u003d C 2 H 4 Cl 2 (in the gas phase).
In water, chlorine partially dissolves (physically), and partially reacts reversibly with it (see § 11.4 c). With a cold solution of potassium hydroxide (and any other alkali), a similar reaction proceeds irreversibly:
Cl 2 + 2OH \u003d Cl + ClO + H 2 O.
As a result, a solution of chloride and potassium hypochlorite is formed. In the case of reaction with calcium hydroxide, a mixture of CaCl 2 and Ca(ClO) 2 is formed, called bleach.
With hot concentrated solutions of alkalis, the reaction proceeds differently:
3Cl 2 + 6OH = 5Cl + ClO 3 + 3H 2 O.
In the case of reaction with KOH, potassium chlorate, called Berthollet salt, is obtained in this way.
Hydrogen chloride is the only hydrogen bond chlorine. This colorless gas with a suffocating odor is highly soluble in water (it completely reacts with it, forming oxonium ions and chloride ions (see § 11.4). Its solution in water is called hydrochloric or hydrochloric acid. This is one of the most important products of chemical technology, since hydrochloric acid is consumed in many industries.It is also of great importance for humans, in particular because it is contained in gastric juice, contributing to the digestion of food.
Hydrogen chloride used to be produced industrially by burning chlorine in hydrogen. At present, the need for hydrochloric acid is almost completely satisfied through the use of hydrogen chloride, which is formed as a by-product during the chlorination of various organic substances, for example, methane:
CH 4 + Cl 2 \u003d CH 3 + HCl
And laboratories produce hydrogen chloride from sodium chloride by treating it with concentrated sulfuric acid:
NaCl + H 2 SO 4 = HCl + NaHSO 4 (at room temperature);
2NaCl + 2H 2 SO 4 \u003d 2HCl + Na 2 S 2 O 7 + H 2 O (when heated).
Higher oxide chlorine Cl 2 O 7 - colorless oily liquid, molecular substance, acid oxide. As a result of reaction with water, it forms perchloric acid HClO 4 , the only oxoacid of chlorine that exists as an individual substance; the remaining oxoacids of chlorine are known only in aqueous solutions. Information about these chlorine acids is given in table 35.
Table 35
C/O |
Formula |
Name |
Strength |
Name |
hydrochloric |
||||
hypochlorous |
hypochlorites |
|||
chloride |
||||
chlorine |
||||
perchlorates |
Most chlorides are soluble in water. The exceptions are AgCl, PbCl 2 , TlCl and Hg 2 Cl 2 . The formation of a colorless precipitate of silver chloride when a solution of silver nitrate is added to the test solution - qualitative reaction for chloride ion:
Ag + Cl = AgCl
Chlorine can be obtained from sodium or potassium chlorides in the laboratory:
2NaCl + 3H 2 SO 4 + MnO 2 = 2NaHSO 4 + MnSO 4 + 2H 2 O + Cl 2
As an oxidizing agent in the production of chlorine by this method, you can use not only manganese dioxide, but also KMnO 4 , K 2 Cr 2 O 7 , KClO 3 .
Sodium and potassium hypochlorites are found in various household and industrial bleaches. Bleach is also used as a bleach and is also used as a disinfectant.
Potassium chlorate is used in the manufacture of matches, explosives and pyrotechnic compositions. When heated, it decomposes:
4KClO 3 \u003d KCl + 3KClO 4;
2KClO 3 = 2KCl + O 2 (in the presence of MnO 2).
Potassium perchlorate also decomposes, but at a higher temperature: KClO 4 \u003d KCl + 2O 2.
1.Compose molecular reaction equations for which ionic equations are given in the text of the paragraph.
2. Make the equations of the reactions given in the text of the paragraph descriptively.
3. Make up equations of reactions that characterize the chemical properties of a) chlorine, b) hydrogen chloride (and hydrochloric acid), c) potassium chloride and d) barium chloride.
Chemical properties of chlorine compounds
Various allotropic modifications are stable under different conditions element sulfur. Under normal conditions simple matter sulfur is a yellow brittle crystalline substance, consisting of eight-atomic molecules:
This is the so-called rhombic sulfur (or -sulfur) S 8. (The name comes from a crystallographic term characterizing the symmetry of the crystals of this substance). When heated, it melts (113 ° C), turning into a mobile yellow liquid, consisting of the same molecules. With further heating, the cycles are broken and very long polymer molecules are formed - the melt darkens and becomes very viscous. This is the so-called -sulfur S n. Sulfur boils (445 ° C) in the form of diatomic molecules S 2, similar in structure to oxygen molecules. The structure of these molecules, as well as oxygen molecules, cannot be described in terms of the covalent bond model. In addition, there are other allotropic modifications of sulfur.
In nature, there are deposits of native sulfur, from which it is mined. Most of the sulfur that is mined is used to produce sulfuric acid. Part of sulfur is used in agriculture for plant protection. Purified sulfur is used in medicine for the treatment of skin diseases.
From hydrogen compounds sulfur, hydrogen sulfide (monosulfan) H 2 S is of the greatest importance. It is a colorless poisonous gas with the smell of rotten eggs. It is slightly soluble in water. Physical dissolution. To a small extent, protolysis of hydrogen sulfide molecules occurs in an aqueous solution, and to an even lesser extent, hydrosulfide ions formed in this case (see Appendix 13). However, a solution of hydrogen sulfide in water is called hydrosulfide acid (or hydrogen sulfide water).
Hydrogen sulfide burns in air:
2H 2 S + 3O 2 \u003d 2H 2 O + SO 2 (with excess oxygen).
A qualitative reaction to the presence of hydrogen sulfide in the air is the formation of black lead sulfide (blackening of filter paper moistened with a solution of lead nitrate:
H 2 S + Pb 2 + 2H 2 O \u003d PbS + 2H 3O
The reaction proceeds in this direction due to the very low solubility of lead sulfide.
In addition to hydrogen sulfide, sulfur forms other sulfanes H 2 S n, for example, disulfan H 2 S 2 , similar in structure to hydrogen peroxide. It is also a very weak acid; its salt is pyrite FeS 2 .
In accordance with the valence capabilities of its atoms, sulfur forms two oxide: SO 2 and SO 3 . Sulfur dioxide (the trivial name is sulfur dioxide) is a colorless gas with a pungent, cough-inducing odor. Sulfur trioxide (the old name is sulfuric anhydride) is a solid, extremely hygroscopic non-molecular substance that turns into a molecular substance when heated. Both oxides are acidic. When reacted with water, they form sulfur and sulfuric acids, respectively. acids.
In dilute solutions, sulfuric acid is a typical strong acid with all their characteristic properties.
Pure sulfuric acid, as well as its concentrated solutions, are very strong oxidizing agents, and the oxidizing atoms here are not hydrogen atoms, but sulfur atoms, passing from the oxidation state + VI to the oxidation state + IV. As a result, OVR with concentrated sulfuric acid usually produces sulfur dioxide, for example:
Cu + 2H 2 SO 4 \u003d CuSO 4 + SO 2 + 2H 2 O;
2KBr + 3H 2 SO 4 \u003d 2KHSO 4 + Br 2 + SO 2 + 2H 2 O.
Thus, even metals that are in the voltage series to the right of hydrogen (Cu, Ag, Hg) react with concentrated sulfuric acid. At the same time, some fairly active metals (Fe, Cr, Al, etc.) do not react with concentrated sulfuric acid, this is due to the fact that a dense protective film is formed on the surface of such metals under the action of sulfuric acid, preventing further oxidation. This phenomenon is called passivation.
Being a dibasic acid, sulfuric acid forms two rows salts: medium and sour. Acid salts are isolated only for alkaline elements and ammonium, the existence of other acid salts is doubtful.
Most medium sulfates are soluble in water and, since the sulfate ion is practically not an anionic base, they do not undergo anionic hydrolysis.
Modern industrial methods for the production of sulfuric acid are based on the production of sulfur dioxide (1st stage), its oxidation to trioxide (2nd stage) and the interaction of sulfur trioxide with water (3rd stage).
Sulfur dioxide is obtained by burning sulfur or various sulfides in oxygen:
S + O 2 \u003d SO 2;
4FeS 2 + 11O 2 \u003d 2Fe 2 O 3 + 8SO 2.
The process of roasting sulfide ores in non-ferrous metallurgy is always accompanied by the formation of sulfur dioxide, which is used to produce sulfuric acid.
Under normal conditions, sulfur dioxide cannot be oxidized with oxygen. Oxidation is carried out by heating in the presence of a catalyst, vanadium(V) oxide or platinum. Although the reaction
2SO 2 + O 2 2SO 3 + Q
reversible, the yield reaches 99%.
If the resulting gas mixture of sulfur trioxide with air is passed through pure water, most of the sulfur trioxide is not absorbed. To prevent losses, the gas mixture is passed through sulfuric acid or its concentrated solutions. In this case, disulfuric acid is formed:
SO 3 + H 2 SO 4 \u003d H 2 S 2 O 7.
A solution of disulfuric acid in sulfuric acid is called oleum and is often represented as a solution of sulfur trioxide in sulfuric acid.
By diluting oleum with water, both pure sulfuric acid and its solutions can be obtained.
1.Compose structural formulas
a) sulfur dioxide, b) sulfur trioxide,
c) sulfuric acid, d) disulfuric acid.
Chlorine forms four oxygen-containing acids: chlorous, chlorine, chloric and perchloric.
Hypochlorous acid HClO It is formed by the interaction of chlorine with water, as well as its salts with strong mineral acids. It is a weak acid and is very unstable. The composition of the reaction products of its decomposition depends on the conditions. With strong illumination of hypochlorous acid, the presence of a reducing agent in the solution, as well as prolonged standing, it decomposes with the release of atomic oxygen: HclO \u003d HCl + O
In the presence of water-removing substances, chlorine oxide (I) is formed: 2 HClO \u003d 2 H2O + Cl2O
Therefore, when chlorine interacts with a hot alkali solution, salts are formed not of hydrochloric and hypochlorous, but of hydrochloric and hypochlorous acids: 6 NaOH + 3 Cl2 = 5 NaCl + NaClO3 + 3 H2O
Salts of hypochlorous acid - g and p about chlorites are very strong oxidizing agents. They are formed by the interaction of chlorine with alkalis in the cold. At the same time hydrochloric acid salts are formed. Of these mixtures, bleach and shale water are most widely used.
Chloric acid HClO2 is formed by the action of concentrated sulfuric acid on alkali metal chlorites, which are obtained as intermediate products during the electrolysis of solutions of alkali metal chlorides in the absence of a diaphragm between the cathode and anode spaces. It is a weak, unstable acid, a very strong oxidizing agent in an acidic environment. When it interacts with hydrochloric acid, chlorine is released: HClO2 + 3 HC1 = Cl2 + 2 H2O
Perchloric acid HClO3 is formed by the action of its salts - chlorates- sulfuric acid. It is a very unstable acid, a very strong oxidizing agent. Can only exist in dilute solutions. By evaporating a solution of HClO3 at low temperature in a vacuum, a viscous solution containing about 40% perchloric acid can be obtained. At a higher acid content, the solution decomposes with an explosion. Explosive decomposition also occurs at lower concentrations in the presence of reducing agents. In dilute solutions, perchloric acid exhibits oxidizing properties, and the reactions proceed quite calmly:
HClO3 + 6 HBr = HCl + 3 Br2 + 3 H2O
Salts of chloric acid - chlorates - are formed during the electrolysis of chloride solutions in the absence of a diaphragm between the cathode and anode spaces, as well as when chlorine is dissolved in a hot alkali solution, as shown above. Potassium chlorate (Berthollet salt) formed during electrolysis is slightly soluble in water and easily separated from other salts in the form of a white precipitate. Like an acid, chlorates are fairly strong oxidizing agents:
KClO3 + 6 HCl = KCl + 3 Cl2 + 3 H2O
Chlorates are used for the production of explosives, as well as for the production of oxygen in the laboratory and salts of perchloric acid - perchlorates. When Bertolet salt is heated in the presence of manganese dioxide MnO2, which plays the role of a catalyst, oxygen is released. If potassium chlorate is heated without a catalyst, then it decomposes with the formation of potassium salts of hydrochloric and perchloric acids:
2 KClO3 = 2 KCl + 3 O2
4 KClO3 = KCl + 3 KClO4
When perchlorates are treated with concentrated sulfuric acid, perchloric acid can be obtained:
KClO4 + H2SO4 = KHSO4 + HclO4
This is the strongest acid. It is the most stable of all oxygen-containing chlorine acids, but anhydrous acid can explode explosively when heated, shaken, or in contact with reducing agents. Dilute solutions of perchloric acid are quite stable and safe to use. Chlorates of potassium, rubidium, cesium, ammonium and most organic bases are poorly soluble in water.
In industry, potassium perchlorate is obtained by electrolytic oxidation of berthollet salt:
2 H+ + 2 e- \u003d H2 (at the cathode)
СlО3- - 2 e- + Н2О = СlO4- + 2 Н+ (on the anode)
biological role.
it belongs to the essential essential elements. In the human body 100 g.
Chlorine ions play a very important biological role. Entering together with K+, Mg2+, Ca2+, HCO~, H3P04 ions and proteins, they play a leading role in creating a certain level of osmotic pressure (osmotic homeostasis) of blood plasma, lymph, cerebrospinal fluid, etc.
Chlorine ion is involved in the regulation of water-salt metabolism and the volume of fluid retained by tissues, maintaining the pH of the intracellular fluid and the membrane potential created by the operation of the sodium-potassium pump, which is explained (as in the case of its participation in osmosis) by the ability to diffuse through cell membranes like the way Na +, K + ions do it. Chlorine ion is a necessary component (together with H2PO4, HSO4 ions, enzymes, etc.) of gastric juice, which is part of hydrochloric acid.
Promoting digestion, hydrochloric acid destroys a variety of pathogenic bacteria.
At raising the.OK . chlorine acid resistance too growing .
The increase in stability is explained by:
a) hardening bonds in anions due to a decrease in the number of NEPs in chlorine,
b) increasing attitude the number of π-overlaps to the number of σ-bonds from 0/1 in ClO − to 3/4 in ClO − 4 . Compare the graphic formulas of acids:
H - O - Cl, H - O - Cl \u003d O, H - O - Cl \u003d O H - O - Cl \u003d O
c) from HClO to HClO 4 grows symmetry anion (both by increasing
the number of oxygen atoms, and as a result of a decrease polarizing actions
hydrogen due to the weakening of its bond with the anion).
d) is decreasing attack angle chlorine atom (i.e. its spatial availability for interaction).
Acidic properties of halogen hydroxides. Acid-base properties
of any hydroxide depend on the ratio of the bond strengths H - O and O - E in
fragment H - O - E. Obviously, the greater the electronegativity of the element, the more the electron density from the H - O bond is shifted to the O - E bond
(H - O - E) and the more acidic properties the hydroxide exhibits.
Therefore, an important factor is nature halogen. So, in the transition from chlorine to iodine, in accordance with a decrease in the value of E.O. acidic properties of hydroxides are reduced. And so much so that iodous acid dissociates along acidic type in lesser degrees HIO → H + + IO - (K d \u003d 4 ∙ 10 - 13),
than in the main: IOH → I + + OH - (K d = 3 ∙10 - 10).
Even a neutralization reaction is possible (but reversible): IOH + HNO 3 → INO 3 + H 2 O.
Salts of chlorine acids, as more stable (than acids) compounds, all
isolated in the free state, but also their activity increases with a decrease in Art. Cl. So, KClO 3 (Bertolet's salt) oxidizes iodide ions only in an acidic environment, and KClO - in a neutral one.
2.8.1. Hypochlorous acid HCl +1 O H–O–Cl (hypochlorites)
physical properties. It exists only in the form of dilute aqueous solutions.
Receipt.
Cl 2 + H 2 O ↔ HCl + HClO
Chemical properties.
HClO is a weak acid and a strong oxidizing agent:
1) Decomposes, releasing atomic oxygen
HClO – exposed to light → HCl + O HClO – vol. conv. → H 2 O + Cl 2 O НClO --- t → НCl + НClO 3
2) With alkalis gives salts - hypochlorites
HClO + KOH → KClO + H 2 O CaOCl 2 - bleach (bleach)
CaOCl 2 + CO 2 + H 2 O → CaCO 3 + CaCl 2 + HClO (HCl + O)
3) with a strong reducing agent HI
2HI + HClO → I 2 ↓ + HCl + H 2 O
2.8.2. Chloric acid HCl +3 O 2 H–O–Cl=O (chlorites)
physical properties. Exists only in aqueous solutions.
Receipt
It is formed by the interaction of hydrogen peroxide with chlorine (IV) oxide, which is obtained from Berthollet salt and oxalic acid in H 2 SO 4 medium:
2KClO 3 + H 2 C 2 O 4 + H 2 SO 4 → K 2 SO 4 + 2CO 2 + 2СlO 2 + 2H 2 O
2ClO 2 + H 2 O 2 → 2HClO 2 + O 2
Chemical properties
HClO 2 is a weak acid and a strong oxidizing agent.
1) HClO 2 + KOH → KClO 2 + H 2 O
KClO 2 + KI + H 2 SO 4 → I 2 + KCl + K 2 SO 4 + H 2 O
2) Unstable, decomposes during storage
4HClO 2 → HCl + HClO 3 + 2ClO 2 + H 2 O
5HClO 2 ---t→ 3HClO 3 + Cl 2 + H 2 O
2.8.3. Perchloric acid HCl +5 O 3 (chlorates)
Physical properties: Stable only in aqueous solutions.
Receipt: Ba (ClO 3) 2 + H 2 SO 4 → 2HClO 3 + BaSO 4 ↓
Chemical properties
HClO 3 - Strong acid and strong oxidizing agent; salts of chloric acid -
chlorates:
6P + 5HClO 3 → 3P 2 O 5 + 5HCl HClO 3 + KOH → KClO 3+ H2O
- KClO 3 - Berthollet's salt; it is obtained by passing chlorine through a heated (40 ° C) KOH solution: 3Cl 2 + 6KOH → 5KCl + KClO 3 + 3H 2 O
Berthollet's salt is used as an oxidizing agent; when heated, it decomposes:
4KClO 3 - no cat → KCl + 3KClO 4 2KClO 3 - MnO2 cat → 2KCl + 3O 2
2.8.4. Perchloric acid HCl +7 O 4 (perchlorates)
Physical properties: Colorless liquid, bp = 25°C, t°pl.= -101°C.
Receipt: KClO 4 + H 2 SO 4 → KHSO 4 + HClO 4
Chemical properties:
HClO 4 is a very strong acid and a very strong oxidizing agent;
salts of perchloric acid - perchlorates .
1) HClO 4 + KOH → KClO 4 + H 2 O
2) When heated, perchloric acid and its salts decompose:
4HClO 4 - t ° → 4ClO 2 + 3O 2 + 2H 2 O KClO 4 - t ° → KCl + 2O 2
Hydrogen bromide HBr (BROMIDES)
Physical Properties
Colorless gas, highly soluble in water; t°boiling = -67°C; t°pl. = -87°С.
Receipt
1) 2NaBr + H 3 PO 4 - t ° → Na 2 HPO 4 + 2HBr 2) PBr 3 + 3H 2 O → H 3 PO 3 + 3HBr
Chemical properties
An aqueous solution of hydrogen bromide - hydrobromic acid is even stronger than hydrochloric acid. It enters into the same reactions as HCl
1) Dissociation: HBr ↔ H+ + Br -
2) With metals in the voltage series up to hydrogen:
Mg + 2HBr → MgBr 2 + H 2
3) with metal oxides:
CaO + 2HBr → CaBr 2 + H 2 O
4) with bases and ammonia:
NaOH + HBr → NaBr + H 2 O Fe(OH) 3 + 3HBr → FeBr 3 + 3H 2 O NH 3 + HBr → NH 4 Br
5) with salts
MgCO 3 + 2HBr → MgBr 2 + H 2 O + CO 2
Qualitative response: AgNO 3 + HBr → AgBr↓ + HNO 3
The formation of a yellow precipitate of silver bromide insoluble in acids serves to detect the anion Br - in solution.
6) restorative properties:
2HBr + H 2 SO 4 (conc.) → Br 2 + SO 2 + 2H 2 O 2HBr + Cl 2 → 2HCl + Br 2
Of the oxygen acids of bromine are known
Weak bromous HBr +1 O and
Strong brominate HBr +5 O 3 .
Hydrogen iodide (iodides)
Physical properties: Colorless gas with a pungent odor, soluble in water
t°boiling = -35°С; t°pl. = -51°С.
Receipt:
1) I 2 + H 2 S → S + 2HI 2) 2P + 3I 2 + 6H 2 O → 2H 3 PO 3 + 6HI
Chemical properties
1) Solution of HI in water - strong hydroiodic acid:
HI ↔ H + + I - 2HI + Ba(OH) 2 → BaI 2 + 2H 2 O
Salts of hydroiodic acid - iodides (for other HI reactions, see St. HCl and HBr)
2) HI is a very strong reducing agent:
2HI + Cl 2 → 2HCl + I 2
8HI + H 2 SO 4 (conc.) → 4I 2 + H 2 S + 4H 2 O
5HI + 6KMnO 4 + 9H 2 SO 4 → 5HIO 3 + 6MnSO 4 + 3K 2 SO4 + 9H 2 O
3)Qualitative response: The formation of a dark yellow precipitate of silver iodide, insoluble in acids, serves to detect the iodine anion in solution.
NaI + AgNO 3 → AgI↓ + NaNO 3 HI + AgNO 3 → AgI↓ + HNO 3
3.0.1. Oxygen acids of iodine ( iodates )
a) Iodonic acid HI +5 O 3
Colorless crystalline substance, t°pl.= 110°С, highly soluble in water.
Get: 3I 2 + 10HNO 3 → 6HIO 3 + 10NO + 2H 2 O
HIO 3 is a strong acid (salts - iodates) and a strong oxidizing agent.
b) Iodine acid H 5 I +7 O 6
Crystalline hygroscopic substance, highly soluble in water,
t°pl.= 130°С. Weak acid (salts - periodates); strong oxidizing agent.
All chlorine oxides have a pungent odor, are thermally and photochemically unstable, and prone to explosive decomposition. +1 Cl 2 O T. Pl o. C bp °C -120, 6 +3 +4 +4 +5 +6 +7 Cl 2 O 3 Cl. O 2 Cl 2 O 4 Cl 2 O 5 Cl 2 O 6 Cl 2 O 7 not received Hcl. O 2 -117 9.7 2.0 -59 not obtained 44.5 -93.4203 HCl. O 3 3 87 HCl. O 4 chloric chloride silt chlorine strong very strong hypochlorites Na chlorates. Cl. O 2 KCl. О 3 perchlorates weak medium strength chlorine KCl. About 4
§ All compounds with chlorine in positive degrees are very strong oxidizing agents. § The most strongly oxidizing properties are expressed in hypochlorous acid, although it is weak and unstable. § Free oxygen-containing acids of chlorine are unstable and, in addition to perchloric acid, exist only in solution. All of them are strong oxidizing agents. § The strength of acids and their oxidizing properties are different concepts. § In the HCl series. O - HCl. O 2 - HCl. O 3 - HCl. O 4 stability and strength of acids increases, and the reactivity decreases.
The ratio of halogens to water ü 2 F 20 + 2 H 2 O− 2 → 4 HF + O 2 interaction, F 2 - oxidizing agent, ü Cl 20 + H 2 O ↔ HCl + 1 O + HCl - 1 interaction, Cl 20 - oxidizing agent , reducing agent; reaction - disproportionation, ü Br 20 + H 2 O ↔ HBr + 1 O + HBr - 1 is highly soluble, there is practically no interaction; Br 20 - oxidizing agent, reducing agent; reaction - disproportionation, ü I 2 + H 2 O ≠ poorly soluble, interaction practically does not occur; ü At 2 + H 2 O ≠ poorly soluble, practically no interaction occurs
Chlorine oxides Comparison parameter Chlorine (I) oxide Chlorine (IV) State of aggregation at n. y. , color Brownish-yellow gas; at t°
Comparison parameter Chlorine oxide (I) Chlorine oxide (IV) Chlorine oxide (VII) Thermal stability Thermally unstable, decomposes in the light Thermally very unstable Decomposes slowly at room temperature The most stable chlorine oxide, decomposes when heated to 120°C pathways Toxic Highly toxic Toxic Relation to water Dissolves well, interacts with water
Methods for the production of chlorine oxides Chlorine oxide Name of the method, CCR Chlorine oxide (I) Interaction of mercury oxide (II) with chlorine at 0°C: Hg. O (solid) + 2 Cl 2 (gas) → Hg. Cl 2 + Cl 2 O Chlorine (IV) oxide 1) Interaction of potassium chlorate with oxalic acid: KCl. O 3 + H 2 C 2 O 4 → K 2 CO 3 + 2 Cl. O 2 + CO 2 + H 2 O (laboratory method); 2) Passing sulfur dioxide SO 2 into an acidified sodium chlorate solution: 2 Na. Cl. O 3 + SO 2 + H 2 SO 4 \u003d 2 Na. HSO 4 + 2 Cl. O 2 (industrial method) Chlorine (VI) oxide Oxidation of chlorine (IV) oxide with ozone: 2 Cl. O 2 + 2 O 3 \u003d 2 O 2 + Cl 2 O 6 Chlorine oxide (VII) Interaction of perchloric acid with phosphoric anhydride - phosphorus oxide (V): 8 HCl. O 4 + P 4 O 10 → 4 Cl 2 O 7 + 2 H 4 P 2 O 7
Chemical properties of chlorine oxides Cl 2 O - chlorine oxide (I) Cl 2 + 1 O + H 2 O \u003d 2 HCl + 1 O not OVR, Cl 2 + 1 O + 2 KOH \u003d 2 KCl + 1 O + H 2 O not OVR, Cl. O 2 - chlorine oxide (IV) 2 Cl + 4 O 2 + H 2 O \u003d HCl + 3 O 2 + HCl + 5 O 3 OVR, Cl + 4 - both reducing agent and oxidizing agent 2 Cl + 4 O 2 + 2 KOH \u003d KCl + 3 O 2 + KCl + 5 O 3 + H 2 O OVR, Cl + 4 - both reducing agent and oxidizing agent Cl 2 O 6 - chlorine oxide (VI) Cl 2 + 6 O 6 + H 2 O \u003d HCl + 5 O 3 + HCl + 7 O 4 OVR, Cl + 6 - both a reducing agent and an oxidizing agent Cl 2 + 6 O 6 + 2 KOH \u003d KCl + 5 O 3 + KCl + 7 O 4 + H 2 O OVR, Cl + 6 - both reducing agent and oxidizing agent Cl 2 O 7 - chlorine oxide (VII) Cl 2+7 O 7 + H 2 O \u003d 2 HCl + 7 O 4 not OVR, but Cl 2 + 7 O 7 + 2 KOH \u003d 2 KCl + 7 O 4 + H 2 O not OVR,
Oxygen-containing chlorine acids Physical properties, preparation methods Chemical properties - relation to heating, alkali solutions and basic oxides
Oxygen-containing acids of chlorine Acid formula Oxidation state of Cl in acid HCl + 1 O HCl + 3 O 2 HCl + 5 O 3 HCl + 7 O 4 +1 +3 +5 +7 Increases Thermal stability Increases Acid strength Increases Very weak acid Weak acid partially dissociates in water Form of existence Acid of medium strength, closer to strong One of the strongest acids dissociates almost irreversibly exists only in solution isolated in free form
Hypochlorous acid is obtained by dissolving chlorine (I) oxide in water (1): (1) Cl 2 O + H 2 O → 2 HCl. O Hypochlorous acid is chlorine water, a solution of chlorine in water. Obtained in a chlorinator by passing chlorine into water until saturation (1 volume of water dissolves about 2.2 volumes of gaseous chlorine at 20 ° C) (2): (2) Cl 2 + H 2 O ⇌ HCl. O + HCl Formed HCl. O decomposes in the light into O 2 and HCl. Chlorine water is a strong oxidizing agent used to disinfect water and bleach fabrics.
Chlorous acid An acid solution is obtained from its salts - chlorites Ba (Cl. O 2) 2 + H 2 SO 4 → 2 HCl. O2 +Ba. SO 4↓ And also according to the reaction: 2 Cl. O 2 + H 2 O → HCl. O 2 + HCl. O 3 Perchloric acid is an acid of medium strength, closer to weak. Chlorites are used for bleaching.
Perchloric acid in aqueous solutions at a concentration below 30% in the cold is quite stable; decomposes in more concentrated solutions: In 8 HCl. O 3 \u003d 4 HCl. O 4 + 3 O 2 + 2 Cl 2 + 2 H 2 O. Perchloric acid is a strong oxidizing agent; oxidizing power increases with increasing concentration and temperature, for example, 40% acid ignites filter paper. Chloric acid in the laboratory is obtained by reacting barium chlorate with dilute sulfuric acid: Ba (Cl. O 3) 2 + H 2 SO 4 = Ba. SO 4↓+ 2 HCl. About 3.
Perchloric acid Anhydrous perchloric acid is obtained by the interaction of sodium or potassium perchlorates with concentrated sulfuric acid or aqueous solutions of perchloric acid with oleum, as well as by the interaction of chlorine oxide (VII) with water: KCl. O 4 + H 2 SO 4 → KHSO 4 + HCl. O 4 Cl 2 O 7 + H 2 O → 2 HCl. O 4
Thermal stability of acids - relation to heating Perchloric acid (HCl. O 4) Ø Can be isolated in free form; Ø With moderate heating with phosphoric anhydride, § 2 HCl decomposes. O 4 + P 2 O 5 \u003d Cl 2 O 7 + 2 HPO 3 Perchloric acid (HCl. O 3) Ø With weak heating, § 8 HCl decomposes. O 3 \u003d 4 HCl. O 4 + 3 O 2 + 2 Cl 2 + 2 H 2 O Chloric acid (HCl. O 2) Ø Very unstable, decomposes at room temperature in the light § 4 HCl. O 2 \u003d HCl + HCl. O 3 + 2 Cl. O 2 + H 2 O Hypochlorous acid (HCl. O) Ø 2 HCl. O \u003d 2 HCl + O 2 (under the influence of light)
Relation to alkali solutions When oxygen-containing acids of chlorine interact with alkali solutions, a salt of this acid and water are formed by an exchange reaction. A neutralization reaction occurs. HCl. O 2 + Na. OH=Na. Cl. O 2 + H 2 O; HCl. O 3 + KOH \u003d KCl. O 3 + H 2 O; Relation to basic oxides When oxygen-containing acids of chlorine interact with basic oxides, a salt of this acid and water are formed by an exchange reaction. 2 HCl. O + Na 2 O = 2 Na. Cl. O + H 2 O; 2 HCl. O 4 + Cu. O \u003d Cu (Cl. O 4) 2 + H 2 O
