Of the entire periodic system, most of the elements represent a group of metals. amphoteric, transitional, radioactive - there are a lot of them. All metals play a huge role not only in nature and human biological life, but also in various industries. No wonder the 20th century was called "Iron".

Metals: general characteristics

All metals share common chemical and physical properties that make them easy to distinguish from non-metals. So, for example, the structure of the crystal lattice allows them to be:

• conductors of electric current;
• good heat conductors;
• malleable and plastic;
• strong and shiny.

Of course, there are differences among them. Some metals shine with a silvery color, others with a more matte white, and still others with red and yellow in general. There are also differences in terms of thermal and electrical conductivity. However, all the same, these parameters are common to all metals, while non-metals have more differences than similarities.

By chemical nature, all metals are reducing agents. Depending on the reaction conditions and specific substances, they can also act as oxidizing agents, but rarely. Capable of forming numerous substances. Chemical compounds of metals are found in nature in large quantities in the composition of ore or minerals, minerals and other rocks. The degree is always positive, it can be constant (aluminum, sodium, calcium) or variable (chromium, iron, copper, manganese).

Many of them are widely used as building materials and are used in various branches of science and technology.

Chemical compounds of metals

Among these, several main classes of substances should be mentioned, which are products of the interaction of metals with other elements and substances.

1. Oxides, hydrides, nitrides, silicides, phosphides, ozonides, carbides, sulfides and others - binary compounds with non-metals, most often belong to the class of salts (except oxides).
2. Hydroxides - the general formula is Me + x (OH) x.
3. Salt. Compounds of metals with acidic residues. May be different:

• medium;
• sour;
• double;
• basic;
• complex.

4. Compounds of metals with organic substances - organometallic structures.

5. Compounds of metals with each other - alloys, which are obtained in different ways.

Metal connection options

Substances that can contain two or more different metals at the same time are divided into:

• alloys;
• double salts;
• complex compounds;
• intermetallics.

Methods for connecting metals to each other also vary. For example, to obtain alloys, the method of melting, mixing and solidifying the resulting product is used.

Intermetallic compounds are formed as a result of direct chemical reactions between metals, often occurring with an explosion (for example, zinc and nickel). Such processes require special conditions: very high temperature, pressure, vacuum, lack of oxygen, and others.

Soda, salt, caustic are all alkali metal compounds found in nature. They exist in their pure form, forming deposits, or are part of the combustion products of certain substances. Sometimes they are obtained in the laboratory. But these substances are always important and valuable, as they surround a person and form his life.

Alkali metal compounds and their uses are not limited to sodium. Also common and popular in the sectors of the economy are salts such as:

• potassium chloride;
• (potassium nitrate);
• potassium carbonate;
• sulfate.

All of them are valuable mineral fertilizers used in agriculture.

Alkaline earth metals - compounds and their applications

This category includes elements of the second group of the main subgroup of the system of chemical elements. Their permanent oxidation state is +2. These are active reducing agents that easily enter into chemical reactions with most compounds and simple substances. Show all the typical properties of metals: brilliance, ductility, heat and electrical conductivity.

The most important and common of these are magnesium and calcium. Beryllium is amphoteric, while barium and radium are rare elements. All of them are capable of forming the following types of connections:

• intermetallic;
• oxides;
• hydrides;
• binary salts (compounds with non-metals);
• hydroxides;
• salts (double, complex, acidic, basic, medium).

Consider the most important compounds from a practical point of view and their applications.

Magnesium and calcium salts

Such compounds of alkaline earth metals as salts are important for living organisms. After all, calcium salts are the source of this element in the body. And without it, the normal formation of the skeleton, teeth, horns in animals, hooves, hair and coat, and so on, is impossible.

So, the most common salt of the alkaline earth metal calcium is carbonate. Its other names are:

• marble;
• limestone;
• dolomite.

It is used not only as a supplier of calcium ions to a living organism, but also as a building material, raw material for chemical industries, in the cosmetic industry, glass, and so on.

Alkaline earth metal compounds such as sulfates are also important. For example, barium sulfate (medical name "barite porridge") is used in X-ray diagnostics. Calcium sulfate in the form of crystalline hydrate is a gypsum found in nature. It is used in medicine, construction, stamping casts.

Phosphorus from alkaline earth metals

These substances have been known since the Middle Ages. Previously, they were called phosphors. This name still occurs today. By their nature, these compounds are sulfides of magnesium, strontium, barium, calcium.

With a certain processing, they are able to exhibit phosphorescent properties, and the glow is very beautiful, from red to bright purple. This is used in the manufacture of road signs, workwear and other things.

Complex compounds

Substances that include two or more different elements of a metallic nature are complex compounds of metals. Most often they are liquids with beautiful and multi-colored colors. Used in analytical chemistry for the qualitative determination of ions.

Such substances are capable of forming not only alkali and alkaline earth metals, but also all the others. There are hydroxocomplexes, aquacomplexes and others.

alkali metals- these are elements of the 1st group of the periodic table of chemical elements (according to the outdated classification - elements of the main subgroup of group I): lithium Li, sodium Na, potassium K, rubidium rb, cesium cs, francium Fr, and ununenniy Uue. When alkali metals are dissolved in water, soluble hydroxides are formed, called alkalis.

Chemical properties of alkali metals

Due to the high chemical activity of alkali metals with respect to water, oxygen, and sometimes even nitrogen (Li, Cs), they are stored under a layer of kerosene. To carry out the reaction with an alkali metal, a piece of the required size is carefully cut off with a scalpel under a layer of kerosene, the metal surface is thoroughly cleaned from the products of its interaction with air in an argon atmosphere, and only then the sample is placed in the reaction vessel.

1. Interaction with water. An important property of alkali metals is their high activity with respect to water. Lithium reacts most calmly (without explosion) with water:

When carrying out a similar reaction, sodium burns with a yellow flame and a small explosion occurs. Potassium is even more active: in this case, the explosion is much stronger, and the flame is colored purple.

2. Interaction with oxygen. The combustion products of alkali metals in air have a different composition depending on the activity of the metal.

· Only lithium burns in air to form an oxide of stoichiometric composition:

・When burning sodium Na 2 O 2 peroxide is mainly formed with a small admixture of NaO 2 superoxide:

In combustion products potassium, rubidium and cesium contains mainly superoxides:

To obtain oxides of sodium and potassium, mixtures of hydroxide, peroxide or superoxide are heated with an excess of metal in the absence of oxygen:

For oxygen compounds of alkali metals, the following regularity is characteristic: as the radius of the alkali metal cation increases, the stability of oxygen compounds containing peroxide ion O 2 2− and superoxide ion O 2 − increases.

Heavy alkali metals are characterized by the formation of fairly stable ozonides composition of EO 3 . All oxygen compounds have different colors, the intensity of which deepens in the series from Li to Cs:

Alkali metal oxides have all the properties of basic oxides: they react with water, acidic oxides and acids:

Peroxides and superoxides exhibit the properties of strong oxidizers:

Peroxides and superoxides interact intensively with water, forming hydroxides:

3. Interaction with other substances. Alkali metals react with many non-metals. When heated, they combine with hydrogen to form hydrides, with halogens, sulfur, nitrogen, phosphorus, carbon and silicon to form, respectively, halides, sulfides, nitrides, phosphides, carbides and silicides:

When heated, alkali metals are able to react with other metals, forming intermetallics. Alkali metals react actively (with an explosion) with acids.

Alkali metals dissolve in liquid ammonia and its derivatives - amines and amides:

When dissolved in liquid ammonia, an alkali metal loses an electron, which is solvated by ammonia molecules and gives the solution a blue color. The resulting amides are easily decomposed by water with the formation of alkali and ammonia:

Alkali metals interact with organic substances, alcohols (with the formation of alcoholates) and carboxylic acids (with the formation of salts):

4. Qualitative determination of alkali metals. Since the ionization potentials of alkali metals are small, when a metal or its compounds is heated in a flame, an atom is ionized, coloring the flame in a certain color:

Coloring the flame with alkali metals
and their compounds

alkaline earth metals.

alkaline earth metals- chemical elements of the II group of the periodic table of elements: beryllium, magnesium, calcium, strontium, barium and radium.

Physical properties

All alkaline earth metals are gray, solid substances at room temperature. Unlike alkali metals, they are much harder, and they are mostly not cut with a knife (the exception is strontium). The density of alkaline earth metals with a serial number increases, although an increase is clearly observed only starting from calcium, which has the lowest density among them (ρ = 1.55 g / cm³), the heaviest is radium, whose density is approximately equal to the density of iron.

Chemical properties

Alkaline earth metals have an electronic configuration of the outer energy level ns², and are s-elements, along with alkali metals. Having two valence electrons, alkaline earth metals easily donate them, and in all compounds they have an oxidation state of +2 (very rarely +1).

The chemical activity of alkaline earth metals increases with increasing serial number. Beryllium in a compact form does not react with either oxygen or halogens, even at a red heat temperature (up to 600 ° C, an even higher temperature is needed to react with oxygen and other chalcogens, fluorine is an exception). Magnesium is protected by an oxide film at room temperature and higher (up to 650 °C) temperatures and does not oxidize further. Calcium oxidizes slowly and at room temperature in depth (in the presence of water vapor), and burns out with slight heating in oxygen, but is stable in dry air at room temperature. Strontium, barium, and radium rapidly oxidize in air to give a mixture of oxides and nitrides, so they, like the alkali metals (and calcium), are stored under a layer of kerosene.

Oxides and hydroxides of alkaline earth metals tend to increase in basic properties with increasing serial number: Be (OH) 2 - amphoteric, water-insoluble hydroxide, but soluble in acids (and also exhibits acidic properties in the presence of strong alkalis), Mg (OH) 2 - weak base, insoluble in water, Ca (OH) 2 - strong, but slightly soluble in water base, Sr (OH) 2 - more soluble in water than calcium hydroxide, strong base (alkali) at high temperatures close to the boiling point water (100 ° C), Ba (OH) 2 - a strong base (alkali), not inferior in strength to KOH or NaOH, and Ra (OH) 2 - one of the strongest alkalis, a very corrosive substance

Being in nature

All alkaline earth metals are found (in varying amounts) in nature. Due to their high chemical activity, all of them are not found in the free state. The most common alkaline earth metal is calcium, the amount of which is 3.38% (of the mass of the earth's crust). Magnesium is slightly inferior to it, the amount of which is 2.35% (of the mass of the earth's crust). Barium and strontium are also common in nature, which, respectively, are 0.05 and 0.034% of the mass of the earth's crust. Beryllium is a rare element, the amount of which is 6·10 −4% of the mass of the earth's crust. As for radium, which is radioactive, it is the rarest of all alkaline earth metals, but it is always found in small quantities in uranium ores. In particular, it can be separated from there by chemical means. Its content is 1 10 −10% (of the mass of the earth's crust)

Aluminum.

Aluminum- an element of the main subgroup of the third group of the third period of the periodic system of chemical elements of D. I. Mendeleev, with atomic number 13. It is indicated by the symbol Al(lat. Aluminum). Belongs to the group of light metals. The most common metal and the third most common chemical element in the earth's crust (after oxygen and silicon).

simple substance aluminum- light, paramagnetic silver-white metal, easily molded, cast, machined. Aluminum has a high thermal and electrical conductivity, resistance to corrosion due to the rapid formation of strong oxide films that protect the surface from further interaction.

Aluminum was first obtained by the Danish physicist Hans Oersted in 1825 by the action of potassium amalgam on aluminum chloride, followed by distillation of mercury. The modern method of obtaining was developed independently by the American Charles Hall and the Frenchman Paul Héroux in 1886. It consists in dissolving aluminum oxide Al 2 O 3 in a melt of Na 3 AlF 6 cryolite followed by electrolysis using consumable coke or graphite electrodes. This method of obtaining requires large amounts of electricity, and therefore was in demand only in the 20th century.

The production of 1000 kg of crude aluminum requires 1920 kg of alumina, 65 kg of cryolite, 35 kg of aluminum fluoride, 600 kg of anode mass and 17 thousand kWh of DC electricity

Alkali metals easily react with non-metals:

2K + I 2 = 2KI

2Na + H2 = 2NaH

6Li + N 2 = 2Li 3 N (the reaction is already at room temperature)

2Na + S = Na 2 S

2Na + 2C = Na 2 C 2

In reactions with oxygen, each alkali metal exhibits its own individuality: when burned in air, lithium forms an oxide, sodium a peroxide, and potassium a superoxide.

4Li + O 2 = 2Li 2 O

2Na + O 2 \u003d Na 2 O 2

K + O 2 = KO 2

Obtaining sodium oxide:

10Na + 2NaNO 3 \u003d 6Na 2 O + N 2

2Na + Na 2 O 2 \u003d 2Na 2 O

2Na + 2NaOH \u003d 2Na 2 O + H 2

Interaction with water leads to the formation of alkali and hydrogen.

2Na + 2H 2 O \u003d 2NaOH + H 2

Interaction with acids:

2Na + 2HCl \u003d 2NaCl + H 2

8Na + 5H 2 SO 4 (conc.) = 4Na 2 SO 4 + H 2 S + 4H 2 O

2Li + 3H 2 SO 4 (conc.) = 2LiHSO 4 + SO 2 + 2H 2 O

8Na + 10HNO 3 \u003d 8NaNO 3 + NH 4 NO 3 + 3H 2 O

When interacting with ammonia, amides and hydrogen are formed:

2Li + 2NH 3 = 2LiNH 2 + H 2

Interaction with organic compounds:

H ─ C ≡ C ─ H + 2Na → Na ─ C≡C ─ Na + H 2

2CH 3 Cl + 2Na → C 2 H 6 + 2NaCl

2C 6 H 5 OH + 2Na → 2C 6 H 5 ONa + H 2

2CH 3 OH + 2Na → 2CH 3 ONa + H 2

2CH 3 COOH + 2Na → 2CH 3 COOONa + H 2

A qualitative reaction to alkali metals is the coloring of the flame by their cations. Li + ion colors the flame carmine red, Na + ion yellow, K + violet

Alkali metal compounds

Oxides.

Alkali metal oxides are typical basic oxides. They react with acidic and amphoteric oxides, acids, water.

3Na 2 O + P 2 O 5 \u003d 2Na 3 PO 4

Na 2 O + Al 2 O 3 \u003d 2NaAlO 2

Na 2 O + 2HCl \u003d 2NaCl + H 2 O

Na 2 O + 2H + = 2Na + + H 2 O

Na 2 O + H 2 O \u003d 2NaOH

Peroxides.

2Na 2 O 2 + CO 2 \u003d 2Na 2 CO 3 + O 2

Na 2 O 2 + CO \u003d Na 2 CO 3

Na 2 O 2 + SO 2 \u003d Na 2 SO 4

2Na 2 O + O 2 \u003d 2Na 2 O 2

Na 2 O + NO + NO 2 \u003d 2NaNO 2

2Na 2 O 2 \u003d 2Na 2 O + O 2

Na 2 O 2 + 2H 2 O (cold) = 2NaOH + H 2 O 2

2Na 2 O 2 + 2H 2 O (gor.) \u003d 4NaOH + O 2

Na 2 O 2 + 2HCl \u003d 2NaCl + H 2 O 2

2Na 2 O 2 + 2H 2 SO 4 (razor. Hor.) \u003d 2Na 2 SO 4 + 2H 2 O + O 2

2Na 2 O 2 + S = Na 2 SO 3 + Na 2 O

5Na 2 O 2 + 8H 2 SO 4 + 2KMnO 4 \u003d 5O 2 + 2MnSO 4 + 8H 2 O + 5Na 2 SO 4 + K 2 SO 4

Na 2 O 2 + 2H 2 SO 4 + 2NaI \u003d I 2 + 2Na 2 SO 4 + 2H 2 O

Na 2 O 2 + 2H 2 SO 4 + 2FeSO 4 = Fe 2 (SO 4) 3 + Na 2 SO 4 + 2H 2 O

3Na 2 O 2 + 2Na 3 \u003d 2Na 2 CrO 4 + 8NaOH + 2H 2 O

Bases (alkalis).

2NaOH (excess) + CO 2 = Na 2 CO 3 + H 2 O

NaOH + CO 2 (excess) = NaHCO 3

SO 2 + 2NaOH (excess) = Na 2 SO 3 + H 2 O

SiO 2 + 2NaOH Na 2 SiO 3 + H 2 O

2NaOH + Al 2 O 3 2NaAlO 2 + H 2 O

2NaOH + Al 2 O 3 + 3H 2 O \u003d 2Na

NaOH + Al(OH) 3 = Na

2NaOH + 2Al + 6H 2 O \u003d 2Na + 3H 2

2KOH + 2NO 2 + O 2 = 2KNO 3 + H 2 O

KOH + KHCO 3 \u003d K 2 CO 3 + H 2 O

2NaOH + Si + H 2 O \u003d Na 2 SiO 3 + H 2

3KOH + P 4 + 3H 2 O \u003d 3KH 2 PO 2 + PH 3

2KOH (cold) + Cl 2 = KClO + KCl + H 2 O

6KOH (hot) + 3Cl 2 = KClO 3 + 5KCl + 3H 2 O

6NaOH + 3S \u003d 2Na 2 S + Na 2 SO 3 + 3H 2 O

2NaNO 3 2NaNO 2 + O 2

NaHCO 3 + HNO 3 \u003d NaNO 3 + CO 2 + H 2 O

NaI → Na + + I –

at the cathode: 2H 2 O + 2e → H 2 + 2OH - 1

at the anode: 2I – – 2e → I 2 1

2H 2 O + 2I - H 2 + 2OH - + I 2

2H2O + 2NaI H 2 + 2NaOH + I 2

2NaCl 2Na + Cl2

at the cathode at the anode

2Na 2 HPO 4 Na 4 P 2 O 7 + H 2 O

KNO 3 + 4Mg + 6H 2 O \u003d NH 3 + 4Mg (OH) 2 + KOH

4KClO 3 KCl + 3KClO 4

2KClO 3 2KCl + 3O 2

KClO 3 + 6HCl \u003d KCl + 3Cl 2 + 3H 2 O

Na 2 SO 3 + S \u003d Na 2 S 2 O 3

Na 2 S 2 O 3 + H 2 SO 4 = Na 2 SO 4 + S↓ + SO 2 + H 2 O

2NaI + Br 2 = 2NaBr + I 2

2NaBr + Cl 2 = 2NaCl + Br 2

I A group.

1. Electric discharges were passed over the surface of the sodium hydroxide solution poured into the flask, while the air in the flask turned brown, which disappears after a while. The resulting solution was carefully evaporated and found that the solid residue is a mixture of two salts. When this mixture is heated, gas is released and only one substance remains. Write the equations of the described reactions.

2. The substance released at the cathode during the electrolysis of a melt of sodium chloride was burned in oxygen. The resulting product was placed in a gasometer filled with carbon dioxide. The resulting substance was added to a solution of ammonium chloride and the solution was heated. Write the equations of the described reactions.

3) The nitric acid was neutralized with baking soda, the neutral solution was carefully evaporated and the residue was calcined. The resulting substance was introduced into a solution of potassium permanganate acidified with sulfuric acid, and the solution became colorless. The nitrogen-containing reaction product was placed in a sodium hydroxide solution and zinc dust was added, and a gas with a pungent odor was released. Write the equations of the described reactions.

4) The substance obtained at the anode during the electrolysis of a sodium iodide solution with inert electrodes was introduced into a reaction with potassium. The reaction product was heated with concentrated sulfuric acid, and the evolved gas was passed through a hot solution of potassium chromate. Write the equations of the described reactions

5) The substance obtained at the cathode during the electrolysis of a melt of sodium chloride was burned in oxygen. The obtained product was sequentially treated with sulfur dioxide and barium hydroxide solution. Write the equations of the described reactions

6) White phosphorus dissolves in a solution of caustic potash with the release of a gas with a garlic odor, which ignites spontaneously in air. The solid product of the combustion reaction reacted with caustic soda in such a ratio that the resulting white substance contains one hydrogen atom; when the latter substance is calcined, sodium pyrophosphate is formed. Write the equations of the described reactions

7) An unknown metal was burned in oxygen. The product of the reaction interacts with carbon dioxide, forms two substances: a solid, which interacts with a solution of hydrochloric acid with the release of carbon dioxide, and a gaseous simple substance that supports combustion. Write the equations of the described reactions.

8) A brown gas was passed through an excess of caustic potash solution in the presence of a large excess of air. Magnesium shavings were added to the resulting solution and heated, nitric acid was neutralized by the evolved gas. The resulting solution was carefully evaporated, the solid reaction product was calcined. Write the equations of the described reactions.

9) During the thermal decomposition of salt A in the presence of manganese dioxide, a binary salt B and a gas that supports combustion and is part of the air were formed; when this salt is heated without a catalyst, salt B and a salt of a higher oxygen-containing acid are formed. When salt A reacts with hydrochloric acid, a yellow-green gas (a simple substance) is released and salt B is formed. Salt B colors the flame purple, and when it interacts with a solution of silver nitrate, a white precipitate forms. Write the equations of the described reactions.

10) Copper shavings were added to heated concentrated sulfuric acid and the released gas was passed through a solution of caustic soda (excess). The reaction product was isolated, dissolved in water, and heated with sulfur, which dissolved as a result of the reaction. Dilute sulfuric acid was added to the resulting solution. Write the equations of the described reactions.

11) Table salt was treated with concentrated sulfuric acid. The resulting salt was treated with sodium hydroxide. The resulting product was calcined with an excess of coal. The resulting gas reacted in the presence of a catalyst with chlorine. Write the equations of the described reactions.

12) Sodium reacted with hydrogen. The reaction product was dissolved in water, and a gas was formed that reacted with chlorine, and the resulting solution, when heated, reacted with chlorine to form a mixture of two salts. Write the equations of the described reactions.

13) Sodium was burned in an excess of oxygen, the resulting crystalline substance was placed in a glass tube and carbon dioxide was passed through it. The gas coming out of the tube was collected and burned in its atmosphere of phosphorus. The resulting substance was neutralized with an excess of sodium hydroxide solution. Write the equations of the described reactions.

14) To the solution obtained as a result of the interaction of sodium peroxide with water during heating, a solution of hydrochloric acid was added until the reaction was completed. The resulting salt solution was subjected to electrolysis with inert electrodes. The gas formed as a result of electrolysis at the anode was passed through a suspension of calcium hydroxide. Write the equations of the described reactions.

15) Sulfur dioxide was passed through a solution of sodium hydroxide until an average salt was formed. An aqueous solution of potassium permanganate was added to the resulting solution. The formed precipitate was separated and treated with hydrochloric acid. The evolved gas was passed through a cold solution of potassium hydroxide. Write the equations of the described reactions.

16) A mixture of silicon (IV) oxide and magnesium metal was calcined. The simple substance obtained as a result of the reaction was treated with a concentrated solution of sodium hydroxide. The evolved gas was passed over heated sodium. The resulting substance was placed in water. Write the equations of the described reactions.

17) The reaction product of lithium with nitrogen was treated with water. The resulting gas was passed through a solution of sulfuric acid until the chemical reactions ceased. The resulting solution was treated with barium chloride solution. The solution was filtered and the filtrate was mixed with sodium nitrate solution and heated. Write the equations of the described reactions.

18) Sodium was heated in a hydrogen atmosphere. When water was added to the resulting substance, gas evolution and the formation of a clear solution were observed. A brown gas was passed through this solution, which was obtained as a result of the interaction of copper with a concentrated solution of nitric acid. Write the equations of the described reactions.

19) Sodium bicarbonate was calcined. The resulting salt was dissolved in water and mixed with a solution of aluminum, as a result, a precipitate formed and a colorless gas was released. The precipitate was treated with an excess of nitric acid solution, and the gas was passed through a solution of potassium silicate. Write the equations of the described reactions.

20) Sodium was fused with sulfur. The resulting compound was treated with hydrochloric acid, the evolved gas completely reacted with sulfur oxide (IV). The resulting substance was treated with concentrated nitric acid. Write the equations of the described reactions.

21) Sodium was burned in excess oxygen. The resulting substance was treated with water. The resulting mixture was boiled, after which chlorine was added to the hot solution. Write the equations of the described reactions.

22) Potassium was heated in a nitrogen atmosphere. The resulting substance was treated with an excess of hydrochloric acid, after which a suspension of calcium hydroxide was added to the resulting mixture of salts and heated. The resulting gas was passed through hot copper (II) oxide. Write the equations for the described reactions.

23) Potassium was burned in an atmosphere of chlorine, the resulting salt was treated with an excess of an aqueous solution of silver nitrate. The precipitate formed was filtered off, the filtrate was evaporated and heated carefully. The resulting salt was treated with an aqueous solution of bromine. Write the equations of the described reactions.

24) Lithium reacted with hydrogen. The reaction product was dissolved in water, and a gas was formed that reacted with bromine, and the resulting solution, when heated, reacted with chlorine to form a mixture of two salts. Write the equations of the described reactions.

25) Sodium was burned in the air. The resulting solid absorbs carbon dioxide, releasing oxygen and salt. The last salt was dissolved in hydrochloric acid, and a solution of silver nitrate was added to the resulting solution. As a result, a white precipitate formed. Write the equations of the described reactions.

26) Oxygen was subjected to an electric discharge in an ozonator. The resulting gas was passed through an aqueous solution of potassium iodide, and a new colorless and odorless gas was released, supporting combustion and respiration. Sodium was burned in the atmosphere of the latter gas, and the resulting solid reacted with carbon dioxide. Write the equations of the described reactions.

I A group.

1. N 2 + O 2 2NO

2NO + O 2 \u003d 2NO 2

2NO 2 + 2NaOH \u003d NaNO 3 + NaNO 2 + H 2 O

2NaNO 3 2NaNO 2 + O 2

2. 2NaCl 2Na + Cl2

at the cathode at the anode

2Na + O 2 \u003d Na 2 O 2

2Na 2 O 2 + 2CO 2 \u003d 2Na 2 CO 3 + O 2

Na 2 CO 3 + 2NH 4 Cl \u003d 2NaCl + CO 2 + 2NH 3 + H 2 O

3. NaHCO 3 + HNO 3 \u003d NaNO 3 + CO 2 + H 2 O

2NaNO 3 2NaNO 2 + O 2

5NaNO 2 + 2KMnO 4 + 3H 2 SO 4 = 5NaNO 3 + 2MnSO 4 + K 2 SO 4 + 3H 2 O

NaNO 3 + 4Zn + 7NaOH + 6H 2 O = 4Na 2 + NH 3

4. 2H2O + 2NaI H 2 + 2NaOH + I 2

2K + I 2 = 2KI

8KI + 5H 2 SO 4 (conc.) = 4K 2 SO 4 + H 2 S + 4I 2 + 4H 2 O

3H 2 S + 2K 2 CrO 4 + 2H 2 O = 2Cr(OH) 3 ↓ + 3S↓ + 4KOH

5. 2NaCl 2Na + Cl2

at the cathode at the anode

2Na + O 2 \u003d Na 2 O 2

Na 2 O 2 + SO 2 \u003d Na 2 SO 4

Na 2 SO 4 + Ba(OH) 2 = BaSO 4 ↓ + 2NaOH

6. P 4 + 3KOH + 3H 2 O \u003d 3KH 2 PO 2 + PH 3

2PH 3 + 4O 2 = P 2 O 5 + 3H 2 O

P 2 O 5 + 4NaOH \u003d 2Na 2 HPO 4 + H 2 O

“Lithium is the lightest metal; it has a specific gravity of 0.59, as a result of which it floats even on oil; melts at about 185°, but does not volatilize in red-hot heat. It resembles sodium in color and, like sodium, has a yellow tint.

D. I. Mendeleev. Fundamentals of chemistry.

When in 1817 the 25-year-old Swedish chemist Johan August Arfvedson (1792-1841) isolated a new “flammable alkali of a hitherto unknown nature” from the mineral petalite (it was lithium hydroxide), his teacher, the famous Swedish chemist Jens Jakob Berzelius (1779-1848), proposed to call it lithion, from the Greek. lithos - stone.

This alkali, in contrast to the already known sodium and potassium, was first discovered in the "kingdom" of stones. In 1818, the English chemist Humphrey Davy (1778-1829) obtained a new metal from lithion, which he called lithium. The same Greek root is in the words "lithosphere", "lithography" (an impression from a stone mold), etc.

Lithium is the lightest of the solids: its density is only 0.53 g/cm3 (half that of water). Lithium is obtained by electrolysis of a melt of lithium chloride. A rare property of metallic lithium is the reaction with nitrogen under normal conditions to form lithium nitride.

Lithium is increasingly used in the production of lithium-ion batteries. As a result, the world production of lithium in 2012 amounted to 37 thousand tons - five times more than in 2005.

Lithium compounds are used in the glass and ceramic industries. Lithium hydroxide is an absorber of excess carbon dioxide in the cabins of spacecraft and submarines. Lithium carbonate is used in psychiatry to treat certain disorders. The average human contains less than 1 mg of lithium.

Sodium

“The production of metallic sodium is one of the most important discoveries in chemistry, not only because the concept of simple bodies has thereby expanded and become more correct, but especially because chemical properties are visible in sodium, only weakly expressed in other well-known metals.”

D. I. Mendeleev. Fundamentals of chemistry.

The Russian name "sodium" (it is also in Swedish and German) comes from the word "natron": this is how the ancient Egyptians called dry soda, which was used in the process of mummification. In the XVIII century, the name "natron" was assigned to the "mineral alkali" - caustic soda. Now soda lime is called a mixture of caustic soda and calcium oxide (in English soda lime), and sodium in English (and in many other languages ​​- sodium). The word "soda" comes from the Latin name of the plant hodgepodge (sodanum). This is a coastal marine plant whose ashes were used in the manufacture of glass in ancient times. This ash contains sodium carbonate, which is called soda. And now soda is the most important component of the charge for the production of most glass, including window glass.

Halite is the main mineral of sodium

The first person to see what metallic sodium looks like was G. Davy, who isolated the new metal by electrolysis. He also proposed the name of the new element - sodium.

Sodium is a very active metal; it quickly oxidizes in air, becoming covered with a thick crust of reaction products with oxygen and water vapor. A lecture experience is known: if a small piece of sodium is thrown into water, it will begin to react with it, releasing hydrogen. A lot of heat is released in the reaction, which melts the sodium, and its ball runs along the surface. Water cools the sodium and prevents the hydrogen from flaring up, but if the piece of sodium is large, a fire and even an explosion is possible.

Sodium metal is widely used in various syntheses as a reducing agent and also as a desiccant for non-aqueous liquids. It is present in high-capacity sodium-sulphur batteries. A low-melting alloy of sodium and potassium, liquid at room temperature, works as a coolant that removes excess thermal energy from nuclear reactors. Everyone knows the yellow color of the flame in the presence of sodium: this is how the flame of a gas burner is colored if the smallest drop of salty soup gets into it. Sodium vapor glows yellow in economical gas-discharge lamps that illuminate the streets.

For centuries, salt has been the only way to preserve food. Without table salt, long-distance sea voyages, round-the-world expeditions and great geographical discoveries would be impossible. The history of Russia knows a grand uprising, called the Salt Riot, which began in 1648 and swept across the country. One of the reasons for the uprising was the increase in the tax on salt.

Once upon a time, hundreds of thousands of tons of sodium were produced per year: it was used to produce tetraethyl lead, which increases the octane number of gasoline. The ban on leaded gasoline in many countries has led to a decline in sodium production. Now the world production of sodium is about 100 thousand tons per year.

The mineral halite (sodium chloride) forms huge deposits of rock salt. Only in Russia, its reserves amount to tens of billions of tons. Halite usually contains up to 8% other salts, mainly magnesium and calcium. More than 280 million tons of sodium chloride are mined annually, this is one of the largest productions. Once upon a time, sodium nitrate was mined in large quantities in Chile, hence its name - Chilean nitrate.

Other sodium salts, of which many are currently known, are also used. One of the most famous is sodium sulfate. If this salt contains water, it is called Glauber's. Huge amounts of it are formed during the evaporation of water in the Kara-Bogaz-Gol Bay of the Caspian Sea (Turkmenistan), as well as in some salt lakes. Currently, sodium sulfate solutions are used as a heat accumulator in devices that store solar energy, in the production of glass, paper, and fabrics.

Salt

Sodium is a vital element. Sodium ions are found mainly in the extracellular fluid and are involved in the mechanism of muscle contractions (a lack of sodium causes convulsions), in maintaining water-salt balance (sodium ions retain water in the body) and acid-base balance (maintaining a constant blood pH value). Hydrochloric acid is produced from sodium chloride in the stomach, without which it is impossible to digest food. The content of sodium in the body of an average person is about 100 g. Sodium enters the body mainly in the form of table salt, its daily dose is 3-6 g. A single dose of more than 30 g is life-threatening.

Potassium

In Arabic, al-qili is ash, and also something calcined. They also began to call the product obtained from the ashes of plants, i.e. potassium carbonate. In sunflower ash, potassium is more than 30%. Without the Arabic article, this word in Russian turned into "potassium". In addition to Russian and Latin (kalium), this term has been preserved in many European languages: German, Dutch, Danish, Norwegian, Swedish (with the Latin ending -um), Greek (κάλιο), as well as in a number of Slavic languages: Serbian (kalyum ), Macedonian (kalium), Slovenian (kalij).

Potassium is one of the most abundant elements in the earth's crust. Its main minerals are sylvin (potassium chloride), sylvinite (mixed potassium and sodium chloride) and carnallite (mixed potassium and magnesium chloride). Silvin, as well as potassium nitrate (potash, it is also Indian nitrate) are used in large quantities as potash fertilizers. Together with nitrogen and phosphorus, potassium is one of the three most important elements for plant nutrition.

Sylvin is one of the main potassium minerals (along with sylvinite and carnallite).

The English name for the element (potassium), like the Russian name for potassium carbonate (potash), is borrowed from the languages ​​of the Germanic group; in English, German and Dutch ash is ash, pot is a pot, i.e. potash is “ash from a pot”. Previously, potassium carbonate was obtained by evaporating the extract from the ash in vats; it was used to make soap. Potassium soap, unlike sodium soap, is liquid. From the Arabic name for ash came the name of alkali in many European languages: English. and goll. alkali, German Alkali, French and ital. alcali etc. The same root is present in the word "alkaloids" i.e. "like alkalis").

Potassium was the first element discovered by G. Davy (he also received lithium, barium, calcium, strontium, magnesium and boron for the first time). Davy electrolyzed a wet chunk of potassium hydroxide. At the same time, according to Davy, “small balls with a strong metallic sheen appeared on its surface, outwardly no different from mercury. Some of them, immediately after their formation, burned out with an explosion and with the appearance of a bright flame, while others did not burn out, but only dimmed, and their surface was covered with a white film. Potassium is a very active metal. His small piece, brought into the water, explodes.

Potassium is an important bioelement, the human body contains from 160 to 250 g of potassium, more than sodium. Potassium ions are involved in the passage of nerve impulses. Fruits and vegetables contain a lot of potassium.

Potassium hydroxide is used to make soap. It serves as an electrolyte in alkaline batteries - iron-nickel, nickel-metal hydride. Previously, potassium nitrate (potassium nitrate) was consumed in large quantities for the production of black powder; now it is used as a fertilizer.

Natural potassium contains 0.0117% of the long-lived radionuclide 40K with a half-life of 1.26 billion years. This explains the fact that potassium-40 "survived" to our time from the moment of its synthesis in nuclear reactions in stars. However, since the formation of the Earth 4.5 billion years ago, the content of 40K on the planet has decreased by 12.5 times due to its decay! A human body weighing 70 kg contains approximately 20 mg 40K, or 3 x 1020 atoms, of which more than 5000 atoms decay every second! It is possible that such "internal" irradiation (enhanced by the decay of carbon-14) was one of the causes of mutations in the course of the evolution of wildlife. World production of potassium metal is small: about 200 tons per year.

rubidium and cesium

Rubidium and cesium are the first chemical elements discovered using spectral analysis. This method was developed by German scientists and friends - the physicist Gustav Robert Kirchhoff (1824-1887) and the chemist Robert Wilhelm Bunsen (1811-1899), who worked at the University of Heidelberg. With this extremely sensitive method, they analyzed all the substances they came across in the hope of finding something new. And in the early 1860s. discovered two new elements. This happened when they analyzed the dry residue obtained by evaporating water from the mineral springs of the Bad Dürkheim resort, 30 km from Heidelberg. In the spectrum of this substance, in addition to the lines of sodium, potassium and lithium already known to them, Kirchhoff and Bunsen noticed two weak blue lines. They realized that these lines belong to an unknown chemical element that is present in water in very small quantities. According to the light of the spectral lines, a new element

Continuing their research, Kirchhoff and Bunsen discovered in the aluminosilicate mineral lepido (lithium mica) sent to them from Saxony, one more element, in the spectrum of which dark red lines stood out. It was called rubidium: from lat. rubidus - red. The same element was found in mineral water, from where the chemist Bunsen managed to isolate it. It is worth mentioning that in order to obtain several grams of rubidium salt, 44 tons of mineral water and over 180 kg of lepidolite had to be processed.

Cesium crystals can be stored in a sealed ampoule.

And just as at the end of the 19th century, in the no less titanic work on the isolation of radium salt, the “compass” for Marie Curie was radioactivity, the similar “compass” for Kirchhoff and Bunsen was the spectroscope.

Rubidium and cesium are typical alkali metals. This was confirmed when the chemist Bunsen, by reducing the salt of rubidium, obtained this element in the form of a metal. More active cesium was obtained in pure form only in 1881 by the Swedish chemist Carl Theodor Setterberg (1853-1941) by electrolysis of molten cesium cyanide. Cesium is one of the most fusible metals. In its pure form, it has a golden color. But it is not easy to obtain pure cesium: in air it instantly ignites spontaneously. Pure rubidium melts at only 39.3 °C, cesium - 10 degrees lower, and on a very hot summer day, samples of these metals in ampoules become liquid.

The world production of metallic rubidium is small - about 3 tons per year. Rubidium-87 is used in medicine: its atoms are absorbed by blood cells, and by the emission of fast electrons by them, with the help of special equipment, “bottlenecks” in the blood vessels can be seen. Rubidium is used in solar cells.

Gustav Kirchhoff (left) and Robert Bunsen discovered rubidium using a spectroscope. In the spectrum of lepidolite, they found dark red lines and gave the name to the new element - rubidium.

The body of a middle-aged person contains approximately 0.7 g of rubidium, and cesium - only 0.04 mg.

Electronic transitions in cesium atoms are used in extremely accurate "atomic clocks". All over the world there are now more than 70 such most accurate clocks - time standards: the error is less than a second in 100 million years. A cesium clock has a unit of time - a second.

It was proposed to use cesium ions to accelerate the rocket using an electric jet engine. In it, ions are accelerated in a strong electrostatic field and ejected through a nozzle.

Electric rocket engines with low thrust are capable of operating for a long time and flying over long distances.

France

This element was discovered (by its radioactivity) in 1939 by Marguerite Perey (1909-1975), an employee of the Radium Institute in Paris, and she named it in honor of her homeland in 1946.

Francium is a neighbor of cesium in the Periodic Table of the Elements. D. I. Mendeleev called the then undiscovered element - ekacesium. This last and heaviest alkali metal is strikingly different from all others in its group. Firstly, no one has ever seen and will not see even the smallest piece of France. Secondly, francium does not have such physical properties as density, melting point and boiling point. So the term "heaviest metal" can only be attributed to its atoms, but not to a simple substance. And all because francium is an artificially obtained highly radioactive element, its longest-lived isotope 223 Fr has a half-life of only 22 minutes. And in order to study the physical properties of a substance, you need to have it in the form of at least the smallest piece. But for France it is impossible.

Marguerite Perey is the first woman elected (in 1962) to the French Academy of Sciences.

Francium is obtained artificially. And as it fuses, its atoms rapidly decay. Moreover, the more accumulated atoms, the more of them decays per unit time. So, in order to simply keep the number of francium atoms constant, they must be synthesized at a rate no less than the rate of their decay. During the synthesis of francium in Dubna by irradiating uranium with a powerful beam of protons, about a million atoms of this element were produced every second. At this rate of synthesis, the rate of decay of the sample becomes equal to the rate of its formation when the number of its atoms is equal to two billion. This is a completely negligible amount of substance, it is not even visible under a microscope.

In addition, these atoms are not assembled into a piece of metal, but are distributed over the surface of the uranium target. So it is not surprising that in the entire globe at any moment there will be no more than two or three tens of grams of francium scattered singly in radioactive rocks.

ALKALI METALS
SUB-GROUP IA. ALKALI METALS
LITHIUM, SODIUM, POTASSIUM, RUBIDIUM, CESIUM, FRANCE
The electronic structure of alkali metals is characterized by the presence of one electron on the outer electron shell, which is relatively weakly bound to the nucleus. Each alkali metal starts a new period in the periodic table. The alkali metal is able to donate its outer electron more easily than any other element of this period. The cut of an alkali metal in an inert medium has a bright silvery sheen. Alkali metals are characterized by low density, good electrical conductivity, and melt at relatively low temperatures (Table 2).
Due to their high activity, alkali metals do not exist in pure form, but occur in nature only in the form of compounds (excluding francium), for example, with oxygen (clays and silicates) or with halogens (sodium chloride). Chlorides are raw materials for obtaining alkali metals in a free state. Sea water contains ALKALINE METALS 3% NaCl and trace amounts of other salts. Obviously, lakes and inland seas, as well as underground salt deposits and brines, contain alkali metal halides in greater concentrations than sea water. For example, the salt content in the waters of the Great Salt Lake (Utah, USA) is 13.827.7%, and in the Dead Sea (Israel) up to 31%, depending on the area of ​​the water surface, which varies with the season. It can be assumed that the insignificant content of KCl in sea water compared to NaCl is explained by the assimilation of the K+ ion by marine plants.
In the free form, alkali metals are obtained by electrolysis of melts of salts such as NaCl, CaCl2, CaF2 or hydroxides (NaOH), since there is no more active metal capable of displacing the alkali metal from the halide. During the electrolysis of halides, it is necessary to isolate the metal released at the cathode, since at the same time gaseous halogen is released at the anode, which actively reacts with the released metal.
See also ALKALI PRODUCTION
Since alkali metals have only one electron on the outer layer, each of them is the most active in its period, so Li is the most active metal in the first period of eight elements, Na, respectively, in the second, and K is the most active metal of the third period, containing 18 elements (first transition period). In the alkali metal subgroup (IA), the ability to donate an electron increases from top to bottom.
Chemical properties. All alkali metals actively react with oxygen, forming oxides or peroxides, differing from each other in this: Li turns into Li2O, and other alkali metals into a mixture of M2O2 and MO2, while Rb and Cs ignite. All alkali metals form with hydrogen salt-like, thermally stable at high temperatures, hydrides of composition M + H, which are active reducing agents; hydrides are decomposed by water with the formation of alkalis and hydrogen and the release of heat, causing ignition of the gas, and the rate of this reaction for lithium is higher than for Na and K.
See also HYDROGEN; OXYGEN.
In liquid ammonia, alkali metals dissolve to form blue solutions, and (unlike with water) can be isolated again by evaporating ammonia or adding an appropriate salt (for example, NaCl from its ammonia solution). When reacting with gaseous ammonia, the reaction proceeds similarly to the reaction with water:

Alkali metal amides exhibit basic properties similar to hydroxides. Most alkali metal compounds, except for some lithium compounds, are highly soluble in water. In terms of atomic size and charge density, lithium is close to magnesium, so the properties of the compounds of these elements are similar. In terms of solubility and thermal stability, lithium carbonate is similar to magnesium and beryllium carbonates of subgroup IIA elements; these carbonates decompose at relatively low temperatures due to the stronger binding of MO. Lithium salts are better soluble in organic solvents (alcohols, ethers, petroleum solvents) than other alkali metal salts. Lithium (like magnesium) reacts directly with nitrogen to form Li3N (magnesium forms Mg3N2), while sodium and other alkali metals can only form nitrides under severe conditions. The metals of subgroup IA react with carbon, but the most easy reaction is with lithium (apparently due to its small radius) and the least easy with cesium. Conversely, active alkali metals directly react with CO, forming carbonyls (for example, K(CO)x), while less active Li and Na only under certain conditions.
Application. Alkali metals are used both in industry and in chemical laboratories, for example, for syntheses. Lithium is used to produce hard light alloys, which differ, however, in brittleness. Large amounts of sodium are consumed to obtain the Na4Pb alloy, from which tetraethyl lead Pb(C2H5)4 is obtained as an antiknock gasoline fuel. Lithium, sodium and calcium are used as components of soft bearing alloys. The only and therefore mobile electron on the outer layer makes alkali metals excellent conductors of heat and electricity. Potassium and sodium alloys, which remain liquid over a wide temperature range, are used as a heat exchange fluid in some types of nuclear reactors and, due to the high temperatures in a nuclear reactor, are used to produce steam. Sodium metal in the form of supply busbars is used in electrochemical technology to transmit high power currents. Lithium hydride LiH is a convenient source of hydrogen released as a result of the reaction of the hydride with water. Lithium aluminum hydride LiAlH4 and lithium hydride are used as reducing agents in organic and inorganic synthesis. Due to the small ionic radius and correspondingly high charge density, lithium is active in reactions with water, therefore lithium compounds are highly hygroscopic, and lithium chloride LiCl is used to dry the air during the operation of devices. Alkali metal hydroxides are strong bases, highly soluble in water; they are used to create an alkaline environment. Sodium hydroxide, as the cheapest alkali, is widely used (in the USA alone, more than 2.26 million tons of it are consumed per year).
Lithium. The lightest metal, has two stable isotopes with atomic masses 6 and 7; the heavy isotope is more common, its content is 92.6% of all lithium atoms. Lithium was discovered by A. Arfvedson in 1817 and isolated by R. Bunsen and A. Mathisen in 1855. It is used in the production of thermonuclear weapons (hydrogen bomb), to increase the hardness of alloys and in pharmaceuticals. Lithium salts are used to increase the hardness and chemical resistance of glass, in the technology of alkaline batteries, and to bind oxygen during welding.
Sodium. Known since antiquity, it was isolated by H. Davy in 1807. It is a soft metal, its compounds such as alkali (sodium hydroxide NaOH), baking soda (sodium bicarbonate NaHCO3) and soda ash (sodium carbonate Na2CO3) are widely used. Metal is also used in the form of vapors in dim gas-discharge lamps for street lighting.
Potassium. Known since antiquity, it was also identified by H. Davy in 1807. Potassium salts are well known: potassium nitrate (potassium nitrate KNO3), potash (potassium carbonate K2CO3), caustic potash (potassium hydroxide KOH), etc. Potassium metal also finds various applications in technologies of heat exchange alloys.
Rubidium was discovered by spectroscopy by R. Bunsen in 1861; contains 27.85% radioactive rubidium Rb-87. Rubidium, like other metals of subgroup IA, is highly reactive and must be stored under a layer of oil or kerosene to avoid oxidation by atmospheric oxygen. Rubidium finds a variety of applications, including photovoltaic technology, radio vacuum devices and pharmaceuticals.
Cesium. Cesium compounds are widely distributed in nature, usually in small quantities together with compounds of other alkali metals. The mineral pollucite silicate contains 34% cesium oxide Cs2O. The element was discovered by R. Bunsen by spectroscopy in 1860. The main application of cesium is the production of photocells and electronic lamps, one of the radioactive isotopes of cesium Cs-137 is used in radiation therapy and scientific research.
France. The last member of the alkali metal family, francium, is so radioactive that it does not exist in the earth's crust in more than trace amounts. Information about francium and its compounds is based on the study of its insignificant amount, artificially obtained (at a high-energy accelerator) during the a-decay of actinium-227. The longest-lived isotope 22387Fr decays in 21 min into 22388Ra and b-particles. According to a rough estimate, the metallic radius of francium is 2.7 . Francium has most of the properties of other alkali metals and is highly electron-donating. It forms soluble salts and hydroxide. Francium exhibits oxidation state I in all compounds.

Collier Encyclopedia. - Open society. 2000 .