S-elements. Chemistry of s-elements General characteristics of s-elements Chemistry of s-elements general characteristics

General consideration.

s-elements are the elements of the main subgroups of groups I and II of the Periodic Table, as well as helium. All of them, except hydrogen and helium, are metals. Group I metals are called alkali metals because they all react with water to form alkalis. Group II metals, with the exception of beryllium, are usually called alkaline earth metals. The emergence of this term is associated with the ancient name of the oxides of these metals - alkaline earths. Francium, which completes group I, and radium, which completes group II, are radioactive elements. The only natural isotope has a short half-life, so its chemical properties ah, not much is known.

All metals have one or two electrons in the outer shell of their atoms. These metals can easily donate their -electrons, forming ions with stable electron configurations of noble gases.

All s-metals are in the solid state under ordinary conditions; none of them forms allotropic modifications. Group I metals are very soft and have low density compared to other metals. Lithium, sodium and potassium are lighter than water and float on its surface, reacting with it. Metals of group II are harder than metals of group I. They have a comparatively higher density, although it is much less than that of transition metals.

Chemical properties of metals.

All metals have a shiny surface when freshly cut, but when they come into contact with oxygen in the air, they vigorously oxidize and quickly become dull. Therefore, all s-metals, with the exception of beryllium and magnesium, must be stored under a layer of kerosene or liquid paraffin to prevent their contact with air. Beryllium and magnesium form a protective oxide layer on the surface and therefore corrode relatively slowly.

All s-metals burn in an air atmosphere, forming oxides of one or more types - normal oxides of composition (Group I) and (II Group), peroxides of composition (Group I) and (II Group), superoxides of composition (Group I) and (II group).

For example, only lithium burns in air to form an oxide

and sodium forms a mixture of peroxide and superoxide

Sodium and potassium oxides can only be obtained if special conditions, for example, when heating a mixture of peroxide with excess metal in the absence of oxygen:

All -metals of groups I and II combine with hydrogen when heated, forming hydrides, for example:

Also, all s-metals, when heated, react with halogens, sulfur, nitrogen, phosphorus and carbon, forming halides

All s-metals of groups I and II reduce cold water to hydroxides and hydrogen:

Their reactivity increases from top to bottom of the group. Thus, lithium reacts with water relatively slowly, while potassium reacts explosively with water, spontaneously igniting and burning with a violet flame on the surface of the water.

The activity of metals of groups I and II towards acids also increases from top to bottom in the group

All alkali metals react explosively with acids, so such reactions are usually not carried out in laboratories.

Compounds of s-metals.

It was stated above that -metals form three types of oxides, which have typical basic properties. With the exception of beryllium and magnesium oxides, oxides, peroxides and superoxides of other elements readily react with water, forming strongly alkaline solutions, for example:

Hydroxides KOH and NaOH are the most important chemical compounds alkali metals. In industry they are obtained by electrolysis of chloride solutions.

S-elements

1. Characteristics of s-elements

The block of s-elements includes 13 elements, common to which is the building of an external energy level in their s-sublevel atoms.

Although hydrogen and helium are classified as s-elements, due to the specific nature of their properties, they should be considered separately. Hydrogen, sodium, potassium, magnesium, calcium are vital elements.

Compounds of s-elements exhibit general patterns in properties, which is explained by the similarity electronic structure their atoms. All outer electrons are valence electrons and take part in the formation of chemical bonds. Therefore, the maximum oxidation state of these elements in compounds is equal to number electrons in the outer layer and is accordingly equal to the number of the group in which the element is located. The oxidation state of s-element metals is always positive. Another feature is that after the electrons of the outer layer are separated, an ion with a noble gas shell remains. When increasing serial number element, atomic radius, the ionization energy decreases (from 5.39 eV y Li to 3.83 eV y Fr), and the reduction activity of the elements increases.

The vast majority of compounds of s-elements are colorless (unlike compounds of d-elements), since the transition of d-electrons from low energy levels to higher energy levels, which causes color, is excluded.

Compounds of elements of groups IA - IIA are typical salts; in an aqueous solution they almost completely dissociate into ions and are not subject to cation hydrolysis (except for Be 2+ and Mg 2+ salts).

hydrogen hydride ionic covalent

Complexation is not typical for s-element ions. Crystalline complexes of s - elements with ligands H 2 O-crystalline hydrates have been known since ancient times, for example: Na 2 B 4 O 7 10H 2 O-borax, KAl (SO 4) 2 12H 2 O-alum. Water molecules in crystalline hydrates are grouped around the cation, but sometimes completely surround the anion. Due to the small ion charge and large ion radius, alkali metals are least prone to forming complexes, including aqua complexes. Lithium, beryllium, and magnesium ions act as complexing agents in complex compounds of low stability.

2. Hydrogen. Chemical properties of hydrogen

Hydrogen is the lightest s-element. Its electronic configuration in the ground state is 1S 1. A hydrogen atom consists of one proton and one electron. The peculiarity of hydrogen is that its valence electron is located directly in the sphere of action atomic nucleus. Hydrogen does not have an intermediate electron layer, so hydrogen cannot be considered an electronic analogue of alkali metals.

Like alkali metals, hydrogen is a reducing agent and exhibits an oxidation state of +1. The spectra of hydrogen are similar to the spectra of alkali metals. What makes hydrogen similar to alkali metals is its ability to produce a hydrated, positively charged H + ion in solutions.

Like a halogen, the hydrogen atom is missing one electron. This determines the existence of the H - hydride ion.

In addition, like halogen atoms, hydrogen atoms are characterized by a high ionization energy (1312 kJ/mol). Thus, hydrogen occupies a special position in the Periodic Table of Elements.

Hydrogen is the most abundant element in the universe, accounting for up to half the mass of the sun and most stars.

On the sun and other planets, hydrogen is in the atomic state, in the interstellar medium in the form of partially ionized diatomic molecules.

Hydrogen has three isotopes; protium 1 H, deuterium 2 D and tritium 3 T, and tritium is a radioactive isotope.

Hydrogen molecules are distinguished by high strength and low polarizability, small size and low mass, and have high mobility. Therefore, hydrogen has a very low temperatures melting (-259.2 o C) and boiling (-252.8 o C). Due to the high dissociation energy (436 kJ/mol), the disintegration of molecules into atoms occurs at temperatures above 2000 o C. Hydrogen is a colorless gas, odorless and tasteless. It has a low density - 8.99·10 -5 g/cm At very high pressures, hydrogen transforms into a metallic state. It is believed that on distant planets solar system- On Jupiter and Saturn, hydrogen is in a metallic state. There is an assumption that the composition of the earth's core also includes metallic hydrogen, where it is found at ultra-high pressure created by the earth's mantle.

Chemical properties. At room temperature molecular hydrogen reacts only with fluorine, when irradiated with light - with chlorine and bromine, when heated with O 2, S, Se, N 2, C, I 2.

Reactions of hydrogen with oxygen and halogens proceed by a radical mechanism.

Interaction with chlorine is an example of an unbranched reaction when irradiated with light (photochemical activation) or when heated (thermal activation).

Сl+ H2 = HCl + H (chain development)

H+ Cl 2 = HCl + Cl

The explosion of a detonating gas - a hydrogen-oxygen mixture - is an example of a branched chain process, when the initiation of the chain includes not one, but several stages:

H 2 + O 2 = 2OH

H+ O 2 = OH+O

O+ H 2 = OH+ H

OH + H 2 = H 2 O + H

An explosion process can be avoided if you work with pure hydrogen.

Since hydrogen is characterized by a positive (+1) and negative (-1) oxidation state, hydrogen can exhibit both reducing and oxidizing properties.

The reducing properties of hydrogen manifest themselves when interacting with non-metals:

H 2 (g) + Cl 2 (g) = 2HCl (g),

2H 2 (g) + O 2 (g) = 2H 2 O (g),

These reactions proceed with the release large quantity heat, which indicates the high energy (strength) of the H-Cl, H-O bonds. Therefore, hydrogen exhibits reducing properties towards many oxides and halides, for example:

This is the basis for the use of hydrogen as a reducing agent for the production of simple substances from halide oxides.

An even stronger reducing agent is atomic hydrogen. It is formed from a molecular electron discharge under low pressure conditions.

Hydrogen has a high reducing activity at the moment of release during the interaction of a metal with an acid. This hydrogen reduces CrCl 3 to CrCl 2:

2CrCl 3 + 2HCl + 2Zn = 2CrCl 2 + 2ZnCl 2 +H 2 ^

The interaction of hydrogen with nitrogen oxide (II) is important:

2NO + 2H2 = N2 + H2O

Used in purification systems for the production of nitric acid.

As an oxidizing agent, hydrogen interacts with active metals:

In this case, hydrogen behaves like a halogen, forming similar to halides hydrides.

Hydrides of s-elements of group I have an ionic structure of the NaCl type. Chemically, ionic hydrides behave like basic compounds.

Covalent hydrides include hydrides of non-metallic elements that are less electronegative than hydrogen itself, for example, hydrides of the composition SiH 4, BH 3, CH 4. By chemical nature, non-metal hydrides are acidic compounds.

A characteristic feature of the hydrolysis of hydrides is the release of hydrogen; the reaction proceeds via a redox mechanism.

Basic hydride

Acid hydride

Due to the release of hydrogen, hydrolysis proceeds completely and irreversibly (?H<0, ?S>0). In this case, basic hydrides form an alkali, and acidic hydrides form an acid.

The standard potential of the system is B. Therefore, the H ion is a strong reducing agent.

In the laboratory, hydrogen is produced by reacting zinc with 20% sulfuric acid in a Kipp apparatus.

Technical zinc often contains small impurities of arsenic and antimony, which are reduced by hydrogen at the time of release to poisonous gases: arsine SbH 3 and stabine SbH This hydrogen can poison you. With chemically pure zinc, the reaction proceeds slowly due to overvoltage and a good hydrogen current cannot be obtained. The rate of this reaction is increased by adding crystals of copper sulfate; the reaction is accelerated due to the formation of a Cu-Zn galvanic couple.

More pure hydrogen is formed by the action of alkali on silicon or aluminum when heated:

In industry, pure hydrogen is produced by electrolysis of water containing electrolytes (Na 2 SO 4, Ba (OH) 2).

A large amount of hydrogen is produced as a by-product during the electrolysis of an aqueous sodium chloride solution with a diaphragm separating the cathode and anode spaces,

The largest amount of hydrogen is obtained by gasification of solid fuel (anthracite) with superheated water steam:

Or by conversion of natural gas (methane) with superheated steam:

The resulting mixture (synthesis gas) is used in the production of many organic compounds. The yield of hydrogen can be increased by passing synthesis gas over the catalyst, which converts CO into CO 2 .

Application. A large amount of hydrogen is consumed in the synthesis of ammonia. For the production of hydrogen chloride and hydrochloric acid, for the hydrogenation of vegetable fats, for the reduction of metals (Mo, W, Fe) from oxides. Hydrogen-oxygen flame is used for welding, cutting and melting metals.

Liquid hydrogen is used as rocket fuel. Hydrogen fuel is environmentally friendly and more energy-intensive than gasoline, so in the future it can replace petroleum products. Already, several hundred cars in the world are powered by hydrogen. The problems of hydrogen energy are related to the storage and transportation of hydrogen. Hydrogen is stored in underground tankers in a liquid state under a pressure of 100 atm. Transporting large quantities of liquid hydrogen poses serious risks.

3. Hydrides. Hydrogen peroxide

Hydrides are compounds of elements with hydrogen. According to the nature of the bond, ionic, covalent and metal hydrides are distinguished.

Ionic (or salt-like) hydrides are formed by alkali or alkaline earth metals and are obtained by heating the metal in a hydrogen atmosphere.

These are white crystalline substances, whose structure is built from H ions? and metal cations.

Ionic hydrides are strong reducing agents. When dissolved in air, the following ignite:

CaH 2 + O 2 = CaO + H 2 O.

They are easily decomposed by water and can be used to produce small quantities of hydrogen:

CaH 2 + 2H 2 O = Ca (OH) 2 + H 2 ^.

Covalent hydrides consist of molecules. Nonmetal hydrides (HCk, H 2 S, NH 3, CH 4, H 2 Se) have a molar structure.

Beryllium, magnesium, and aluminum hydrides have a polymer structure. Here, the metal atoms are united in chains and layered hydride ions, which form three-center two-electron bonds with the metal atoms, for example, AlHAl.

Transition d- and f-elements form metal hydrides.

When moving from left to right in a period, the properties of hydrides change from neutral (SiH 4) to basic (PH 3) and acidic (HCl).

In complex hydrides, H ions? play the role of ligands. An example is aluminum hydrides? and borohydrides [ВH4]? .

Borohydrides are fairly stable compounds, while aluminum hydrides are easily decomposed by water, releasing hydrogen:

4H 2 O = Al (OH) 3) + OH? + 4H 2.

This reaction is used to produce hydrogen. Aluminum hydrides are also used to prepare hydrides of other elements:

GeCl 4 + Li > GeH4 + LiCl + AlCl.

Greatest practical significance has hydrogen peroxide (peroxide) H 2 O 2. Energy O-O communications(210 kJ/mol) is significantly lower than the O-H bond energy (468 kJ/mol). Due to the asymmetrical distribution of bonds H-O molecule H 2 O 2 is highly polar (m = 0.7 10 -29 C m). A strong hydrogen bond occurs between hydrogen peroxide molecules, leading to their association. Therefore, under normal conditions, hydrogen peroxide is a colorless, viscous, transparent liquid with a high boiling point (150.2 o C). Hydrogen peroxide mixes with water in any way, due to the formation of new hydrogen bonds. In the laboratory, 3% and 30% solutions of H 2 O 2 are usually used (the latter is called perhydrol).

In aqueous solutions, hydrogen peroxide is a weak acid:

hydroperoxide ion

IN chemical reactions The peroxide radical can, without changing, transform into other compounds:

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

BaO 2 + H 2 SO 4 = BaSO 4 + H 2 O 2

More often reactions occur that are accompanied by the destruction of the O-O bond or a change in the charge of the O 2 2 - ion. The oxidation state of oxygen in H 2 O 2 is - 1, so hydrogen peroxide can exhibit both the properties of a reducing agent and the properties of an oxidizing agent.

An example of a reaction in which hydrogen peroxide acts as an oxidizing agent is:

When interacting with a very strong oxidizing agent, for example with PbO 2, peroxide acts as a reducing agent:

reducing agent

The oxidizing properties of peroxide are most pronounced in acidic and neutral environments. And reducing ones - in alkaline:

Cl 2 + H 2 O 2 + 2naCl = 2NaCl + 2H 2 O + O 2 ^.

Hydrogen peroxide is characterized by decomposition according to the type of disproportionation:

This decomposition is accelerated in the presence of impurities, lighting, and heating. 30-60% solutions are stable. Hydrogen peroxide is stored in a dark container and in the cold.

The decomposition of hydrogen peroxide is accelerated in the presence of salts heavy metals. The metal ion-catalyzed decomposition of H 2 O 2 can lead to the formation of radicals, the most important of which are hydroxide HO and hydroperoxide HO 2. For example, under the influence of Fe 2+, bonds - O-O- are broken:

Fe 2+ + H 2 O 2 > Fe 3+ + OH - + HO

The resulting radicals are very toxic for the cell. Hydrogen peroxide is used in medical practice as an external bactericidal agent, and H 2 O 2 solutions are used as a disinfectant. Hydrogen peroxide is used to bleach paper, leather, and textile materials.

4. Water chemistry

Water is the main hydrogen compound with unique properties and of vital importance.

The structure of water. Water is one of the most common substances in nature. Its total amount is 1.4 10 18 tons, it covers approximately four-fifths of the earth's surface area. Water is a component of many minerals, rocks and soil. She plays exclusively important role in nature, in the life of plants, animals and humans. Water accounts for approximately 1/3 of the human body weight. Many food products(vegetables, fruits, milk, eggs, meat) are 95-65% water.

There are nine established isotopes of water, of which H 16 2 O is 99.73% (mol fraction), and H 18 2 O is 0.2%. A small amount is due to heavy water D 2 O. Water contains a small amount of radioactive isotope (T 2 O).

It is difficult to overestimate the role of water in technology, agriculture, medicine, as well as in technological processes various industries national economy. At fuel and nuclear power plants, water, for example, is the main working substance - the coolant, and at hydroelectric power plants it is the carrier of mechanical energy. The exclusive role of water in nature and technology is due to its properties. Water is a thermodynamically stable compound. The standard Gibbs energy of formation of liquid water at a temperature of 298 K is 237.57 kJ/mol, water vapor is 228.94 kJ/mol. Accordingly, the dissociation constant of water vapor during decomposition into hydrogen and oxygen is very small:

The dissociation constant approaches unity only at temperatures above 4000K.

Physical properties water. The melting point of water is 0 o C, boiling point is 100 o C. Density at 20 o C is 0.998 g/cm The properties of water differ significantly from the properties of hydrogen compounds of group IV elements (H 2 S, H 2 Se, H 2 Te). Water under normal conditions is in a liquid state, while these compounds are gases. The temperature of crystallization and evaporation of water is significantly higher than the temperature of crystallization and evaporation of hydrogen compounds of group IV elements. Water has its maximum density at a temperature of 4 o C. This is also unusual. Unlike other compounds, the density of water does not increase during crystallization, but decreases. Water has a very high dielectric constant. Thus, at 298 K its dielectric constant is 78.5, while for H 2 S it is less than 10. Water is a good solvent for polar liquids and compounds with ionic bonds.

Water forms crystalline hydrates with many compounds. For example, CH 4 nH 2 O, C 2 H 5 Cl mH 2 O (clathrates or inclusion compounds).

The unusual properties of water are due to three reasons: the polar nature of the molecules, the presence of lone electron pairs on oxygen atoms and the formation of hydrogen bonds. The water molecule has an angular shape with a HOH angle of 104.5°, close to tetrahedral; at the vertex there is an oxygen atom connected to two hydrogen atoms (protons) of the polar covalent bond. Two pairs of electrons are shared between the protons and the oxygen atom, two pairs of lone electrons are oriented on the other side of the oxygen. The water molecule is polar. Due to its polarity, water dissolves polar liquids and compounds with ionic bonds well. The presence of lone pairs of electrons in oxygen and the displacement of shared electron pairs from the hydrogen atom to the oxygen atom determine the formation of hydrogen bonds between oxygen and hydrogen.

Although hydrogen bonds are weaker than covalent and ionic bonds, they are much stronger than van der Waals bonds and determine the association of water molecules in the liquid state and some anomalous properties of water, in particular high melting and vaporization temperatures, high dielectric constant, maximum density at 4 o C, as well as a special structure ice. In ice crystals, a water molecule forms four hydrogen bonds with neighboring water molecules (due to oxygen's two lone electron pairs and two protons), which gives rise to the tetrahedral crystal structure of ice.

In liquid water, molecules are associated, i.e. combined into larger particles. Moreover, an equilibrium is established between water molecules bound into associates and free water molecules. The presence of associates increases the temperature of crystallization and evaporation of water and the dielectric constant. As the temperature increases, the proportion of free molecules increases.

When water evaporates, the associates are destroyed, and water vapor at low pressures consists of free H2O molecules. However, as the pressure increases, the water molecules come closer and form hydrogen bonds. An association of molecules occurs. As the pressure increases, the vapor approaches its structure to a liquid state. This causes an increase in the solubility of compounds with ionic bonds in the vapor.

Chemical properties of water. Water partially dissociates into hydrogen and hydroxide ions (K d.298 = 2·10 -16).

A proton interacts with H 2 O, forming H 3 O +. Water - amphoteric compound, i.e. maybe like acid

and the basis

Water can be both an oxidizing agent and a reducing agent. Redox duality is associated with the possibility of two processes occurring:

(1) oxidation of hydrogen H 2 O + e?SN 2 + OH -, E 0 (pH = 7) = - 0.410 V

(2) reduction of oxygen O 2 + 4H + + 4e = 4H 2 O, E 0 (pH = 7) = 0.815 V.

Strong oxidizing agents oxidize it, releasing oxygen:

H 2 O + F 2 = 2HF + SO 2

Strong reducing agents reduce it with the release of hydrogen, for example:

2H 2 O + Ca = Ca (OH) 2 + H 2

At elevated temperatures, water vapor interacts with CO (on a Fe-catalyst), methane (on a Na- or Co-catalyst):

CO + H 2 O = CO 2 + H 2

CH 4 + 2H 2 O = CO 2 + 4H 2

Water is a ligand and is coordinated by both cations [M (H 2 O) m ] n + and anions [A (H 2 O) m ] n - .

Water catalyzes many reactions. For example, alkali metals react at room temperature even in the presence of traces of water. Since water molecules are polar, they dissolve well many polar compounds that dissociate into ions. Substances that form hydrogen bonds with water (SO 2 , NH 3 , C 2 H 5 OH , etc.) are highly soluble in water. The solubility in water of low-polar substances is low.

4.1 Composition natural waters

Humanity widely uses natural water for its needs. The total water reserves on Earth are enormous. However, the bulk of the water comes from the World Ocean. According to UNESCO (1970), water reserves are distributed as follows: oceans - 97.2%, glaciers and ice caps - 2.15%, groundwater- 0.625%, fresh lakes and rivers - 9·10 - 3%, salt lakes and inland seas - 8·10 - 3%, atmosphere - 10 - 3%, rivers - 10 - 4%, fresh water reserves available for use , constitute only 0.15% of the volume of the hydrosphere (about 0.2 million km 3).

In nature there is a continuous water cycle. Water, evaporating, enters the atmosphere and then falls into precipitation over the ocean (65-75%) and land (35-25%). Natural water is in continuous interaction with environment. It reacts with the atmosphere, soil, vegetation, minerals and various rocks. In this case, water dissolves organic and inorganic compounds. The composition of natural waters is determined by the nature of this interaction.

All impurities in natural waters can be divided into three groups depending on the particle size: truly dissolved, colloidal and suspended. True solutes are in the form of ions and molecules and have sizes less than 1 nm. Colloidal particles have sizes from 1 to 200 nm. Suspended or coarse particles have sizes greater than 0.1 microns. By chemical composition impurities are divided into organic and inorganic. The former usually have very complex composition and are in a colloidal or truly dissolved state. Inorganic impurities are found mainly in the form of ions: Na +, Ca 2+, Mg 2+, K +, Cl -, SO 4 2 -, HCO 3 -. Nitrogen, oxygen, carbon dioxide and other gases are dissolved in water. Between carbonic acid and its anions establish an equilibrium called carbon dioxide:

With increasing pH, the equilibrium shifts towards the formation of carbonate ions, which dominate at pH>10. When the pH decreases, the equilibrium shifts towards the formation of H 2 CO 3, which prevails at pH<6. Вода, у которой угольная кислота, гидрокарбонат - и карбонат-ионы находятся в равновесии, называется стабильной. При сдвиге равновесия в сторону образования угольной кислоты вода становится агрессивной, при этом повышается её коррозионная активность. При сдвиге равновесия в сторону образования карбонат-ионов из воды выпадает малорастворимый карбонат кальция.

To obtain water suitable for drinking, natural waters are purified. The main stages of water treatment include:

1. Separation of large mechanical impurities by passing through a layer of river sand, a filter, and drum screens.

2. Clarification (Water treatment with aluminum sulfate for the purpose of adsorption of mineral and organic impurities that cause color by the resulting aluminum hydroxide).

Disinfection (chlorination or ozonation).

4. Softening.

Clarifying water allows you to get rid of colloidal impurities and heavy metal ions. When aluminum sulfate gets into water, it reacts with the hydrocarbonates it contains:

Al 2 (SO 4) 3 + 3Ca (HCO 3) 2 = 3CaSO 4 v + Al (OH) 3 + 6CO 2

Flaky amorphous hydroxide Al (OH) 3 with a highly developed surface is formed.

Positively charged aluminum ions neutralize the negative charges of colloidal particles, they stick together and are enveloped in Al (OH) flakes. Hydroxo groups located on the surface of the sediment bind the heavy metal ions present in the solution.

The composition of natural waters is characterized by certain technological indicators, including hardness, environmental reaction, alkalinity, salinity, and oxidability. The hardness of water reflects the content of calcium and magnesium ions in it. It is expressed in mmol/l: F = ( + ). There are carbonate and non-carbonate hardness. Carbonate called hardness caused by calcium and magnesium bicarbonates. Non-carbonate hardness is the difference between total and carbonate hardness.

Alkalinity water is expressed by the sum of the concentrations of hydroxide ions and weak acid anions HCO - ; CO 3 2- .

Water is characterized salt content, which is equal to the total salt concentration. The composition of natural waters depends on their type and location of the reservoir or water source. River waters usually have low salt content: 0.5-0.6 g/l. Groundwater has a higher salinity content. The salt content in the waters of the oceans and open seas is approximately the same and amounts to 35 g/l, with the main ions being Na + and Cl -. The salt content of inland seas is lower than that of the oceans. For example, the salt content of the Caspian Sea is 3-23 g/l, and the Black Sea is 17-18 g/l.

Oxidability reflects the content of impurities that can interact with oxidizing agents.

Biochemical oxygen demand (BOD)) determines the oxygen consumption for the decomposition of organic substances through oxidation by bacteria. It is determined by the change in oxygen concentration in water before and after keeping it in the dark for five days at 20 0 C (BOD 5). The BOD is used to judge the degree of water pollution. Water with a BOD of up to 30 mg/l is considered practically clean, with a BOD of 30-80 mg/l - slightly polluted, and with a BOD>80 - highly polluted.

Water use. Fresh natural waters are used in agriculture (about 82%), mainly for irrigation, in everyday life (about 10%), in industry (about 8%) for cooling, and also as an energy carrier, vehicle, solvent.

Table 4

Maximum permissible concentration ions in drinking water(with max, mg/l)

Hardness salts and other poorly soluble impurities of industrial waters are deposited on the walls of boilers and other devices, reducing the efficiency of these devices. Sodium chlorides and some other impurities in boilers turn into steam and then, deposited on turbine blades, change their profile and accordingly reduce the efficiency of power plants. Oxygen, carbon dioxide, iron ions and nitrite ions dissolved in water cause corrosion of metals.

Therefore, natural waters are purified from a significant part of impurities before use.

4.2 Basic chemical and physical-chemical methods of water treatment

The choice of method for removing impurities from water is determined by the nature and properties of the impurities. Thus, suspended impurities are most easily removed from water by filtration, colloidal impurities by coagulation. If ionic impurities can form a poorly soluble compound, then they can be converted into this compound, oxidizing impurities can be eliminated by reduction, and reducing impurities can be eliminated by oxidation. Adsorption is widely used to remove impurities, with uncharged impurities adsorbed on activated carbon or other adsorbents, and ions on ion exchangers. Charged impurities can also be removed by electrochemical methods. Thus, knowledge of the composition and properties of impurities allows you to choose a method of water purification.

For poorly soluble salts at a constant temperature, the constancy of the ion activity products (PR) is observed.

The concentration of an ion in a poorly soluble compound can be reduced by increasing the concentration of an ion of the opposite sign in the same compound. For example, the concentration of Ca 2+ and Mg 2+ ions can be reduced by increasing the concentration of CO 3 2 - and OH - ions, respectively.

The method of precipitation of poorly soluble compounds is used to purify water, for example, to soften it (reduce hardness). To reduce carbonate hardness, the liming method is used, in which lime Ca (OH) 2 is introduced into the treated water. As a result electrolytic dissociation lime:

Ca (OH) 2 >Ca 2+ + 2OH -

The pH of the water increases, which leads to a shift in the carbon dioxide balance towards the formation of carbonate ions:

As a result, the product of the solubility of calcium carbonate with subsequent precipitation is achieved:

Ca 2+ + CO 3 2 - > CaCO 3 v

In addition, with an increase in the concentration of hydroxide ions, the product of the solubility of magnesium hydroxide is achieved, followed by precipitation:

Mg 2+ + 2ОH - > Mg (ОH) 2 v

The reactions that occur when lime is added can be written in molecular form equations:

Ca (HCO 3) 2 + Ca (OH) 2 = 2CaCO 3 + 2H 2 O

Mg (HCO 3) 2 + 2Ca (OH) 2 =Mg (OH) 2 + 2CaCO 3 + 2H 2 O

H 2 CO 3 + Ca (OH) 2 = CaCO 3 + 2H 2 O

As you can see, with the introduction of lime, the concentration of Ca 2+ and Mg 2+ ions decreases (softening), HCO 3 - (decreased alkalinity) and H 2 CO

The liming method is not suitable for reducing non-carbonate hardness. For these purposes, it is necessary to introduce a highly soluble salt containing carbate ions. Usually, Na 2 CO 3 soda is used for this, which, when dissociated, gives CO 3 2 - ions:

Na 2 CO 3 > 2Na + + CO 3 2 - ; CO 3 2 - +Ca 2+ >CaCO 3 v

The carbon dioxide equation can also be shifted to the right when heated:

As a result, the concentration of carbonate ions increases and the solubility product of calcium carbonate, which precipitates, is achieved.

To purify natural waters from impurities, methods of cationization, anionization, and chemical desalination are widely used.

The removal of cations (Mg 2+, Ca 2+, Na +, etc.) is carried out using cation exchangers, and anions (Cl -, SO 4 2, HCO 3 -, etc.) - using anion exchangers.

For example, hardness ions are removed by Na-cationization.

Anions can be removed by OH anionization.

where the subscript(s) indicates the ion exchange resin.

If you carry out OH anionization and remove anions from the solution and H cationization to remove cations from the solution

then H + and OH - ions will pass into the solution, which are neutralized, forming water:

Thus, as a result of ion exchange reactions, cations and anions are removed from the solution, i.e. salt, or in other words, chemical desalting occurs. To remove salts from seawater, the electrolysis method is also used, which is produced in a multi-chamber electrolyzer. Each chamber has a membrane on one side that is permeable only to anions. As a result of electrolysis sea water in some chambers it is enriched with salts (brine is obtained), in other chambers it is depleted in salts (water purification occurs).

Disinfection. To destroy pathogenic bacteria, viruses and microorganisms. Water that causes biological fouling of pipelines and equipment is treated with oxidizing agents. The most common chlorination of water is liquid or gaseous chlorine, hypochlorites NaClO or Ca (ClO) 2. The bactericidal effect of chlorine is mainly caused by hypochlorous acid, which is formed when chlorine reacts with water:

When chlorine interacts with organic substances the appearance of small amounts of toxic substances, for example CHCl 3, is possible, therefore the treatment of water with ozone O 3 (ozonation) is of increasing interest.

5. Group IA elements

S - elements of the first group (lithium, sodium, potassium, rubidium, cesium, francium) - alkali metals. Some information about these elements is given in table.

The atoms of the elements in question have a single valence electron. Compared to elements of other subgroups, they have the lowest ionization energies, the sizes of atoms and ions are the largest, and they have strongly pronounced metallic characteristics. In the atomic and condensed state, these are unconditional reducing agents. The standard electrode potentials of these metals are very low, which indicates their high reducing activity.

Natural resources . Sodium and potassium compounds are very common, and Li, Rb and Cs are rare elements. Rb and Cs are classified as trace elements, their compounds are satellites of potassium minerals. France is negligibly small in nature (one of the Fr isotopes is a product of the decay of actinium).

In the free state, alkali metals are not found, but are found in the form of compounds: Na 2 OAi 2 O 3 6SiO 2 - sodium feldspar, K 2 OAi 2 O 3 6SiO 2 - potassium feldspar, NaCI - halite or rock salt, KS1-sylvite , KS1MgCl 2 6H 2 O - carnallite. The thickness of rock salt layers can be more than a kilometer. In the ash land plants contains K 2 CO 3, in algae ash Na 2 CO Lithium is found in the form of aluminosilicates and aluminophosphates, from which its other compounds are obtained.

Table 5

Properties of group IA elements

Properties
Atomic mass
Valence electrons
Atomic radius, nm
Ion radius, nm
Ionization energy, eV

V earth's crust, %
Standard electrode potential, IN

Receipt . Lithium metal is produced by electrolysis of LiCl and KC1 melts.

Lithium is also obtained by reduction of its oxides:

Si + 2Li 2 O 4Li + SiO 2.

Sodium is obtained by electrolysis of melts containing sodium chloride, as well as by electrolysis of NaOH melt:

Anode: Cathode:

4OH - 4eO 2 +2H 2 O Na + +leNa

Due to high reactivity potassium, several methods for its production have been developed:

1) reduction of potassium from molten KOH or KC1 with sodium;

2) electrolysis of a melt mixture of KS1 and K 2 CO 3 (cathode - liquid lead) followed by distillation from the alloy with lead. A convenient method for obtaining rubidium and cesium is thermal reduction from chlorides With using calcium in a vacuum:

2CsC?+ Ca CaC? 2+2Cs,

2RbC? + Ca CaC? 2 + 2Rb.

Highly volatile rubidium and cesium are distilled off. Na, K, Rb, Cs are purified by vacuum distillation.

Metallic Li, Na, K are stored in sealed iron containers, Rb and Cs in sealed glass ampoules. Small amounts of Li, Na, K are stored in kerosene in laboratories due to their high chemical activity.

Properties . In the solid state in the absence of moisture and air, Li, Na, K, Rb have a metallic luster and a silvery-white color, and Cs-golden-yellow. In air, the metallic luster quickly disappears, and the metal surface becomes covered with an oxide film. Alkali metals have high compressibility and high electrical and thermal conductivity. These are light metals, lithium being the lightest of the solids. Working with alkali metals requires great care, because they ignite easily and react violently with water and other substances. After work, the remaining alkali metals are destroyed by throwing them in small portions into ethanol, which produces sodium alkoxide

2Na + 2C 2 H 5 OН2C 2 H 5 ONa + H 2 .

Connections . Alkali metals react with dry hydrogen, forming EN hydrides:

2Na + H 2 = 2NaH,

2K + N 2 = 2KN.

Alkali metal hydrides are solid crystalline substances having an ionic lattice. The thermal stability of hydrides decreases in the order from LiH to CsH. Alkali metal hydrides are strong reducing agents. They react vigorously with water, releasing hydrogen:

EN + H 2 OEON + H 2,

NaH + H 2 ONaOH + H 2 .

Interact with carbon dioxide:

NaH + CO 2 NaCOOH.

sodium formate

The reactivity of hydrides increases when going from LiH to CsH.

All alkali metals react vigorously with oxygen, forming oxides, peroxides, superoxides:

4Li + O 2 2Li 2 O (lithium oxide),

2Na + O 2 Na 2 O 2 (sodium peroxide).

Potassium, rubidium, cesium with oxygen form superoxides:

Rb + O 2 = RbO 2 (rubidium superoxide),

Cs + O 2 = CsO 2 (cesium superoxide).

Alkali metal oxides E 2 O can be obtained with a lack of oxygen. Oxides Li 2 O, Na 2 O - colorless; K 2 O, Rb 2 O - yellow; Cs 2 O - orange (as the size of the ion increases, and therefore its polarizability, the compounds become colored). Superoxide KO 2 has crystal lattice type KS?, in which the superoxide ion O 2 - is located in the position of chlorine ions. Peroxides are salts of hydrogen peroxide H 2 O 2 . Acid properties H 2 O 2 are weakly expressed and peroxides, when dissolved in water, undergo almost complete hydrolysis:

Na 2 O 2 + 2HOpNaOH + H 2 O 2 .

The hydrolysis of superoxides produces H 2 O 2 and O 2, 2KO 2 + 2HOpKOH + 2H 2 O 2 + O 2.

Peroxides and superoxides of alkali metals are strong oxidizing agents.

Alkali metal oxides react vigorously with water, forming hydroxides:

E 2 O + H 2 O 2EON,

Na 2 O + H 2 O2NaOH.

Alkali metals react even more actively with water:

2Cs + 2H 2 O2CsOH + H 2 (the reaction proceeds explosively).

Chemical properties . Alkali metal hydroxides are colorless crystalline substances. They are fusible and very soluble in water (with the exception of NaOH). These are alkalis (alkali are bases that are highly soluble in water). In practice, NaOH and KOH are used (caustic soda and caustic potash - technical names). Alkalis greedily absorb moisture and CO 2 from the air:

NaOH + CO 2 = NaHCO 3

NaOH + H 2 O = NaOH? H 2 O (crystalline hydrate NaOH)

When melting, alkalis destroy glass and porcelain:

2NaOH (k) + SiO 2 (k) = Na 2 SiO 3 (k) + H 2 O (g).

When exposed to oxygen, alkalis destroy platinum; they are melted in vessels made of silver, nickel or iron, and stored in polyethylene vessels. Solid alkalis and their concentrated solutions destroy living tissue, so working with them requires precautions (rubber gloves, safety glasses). Of the alkalis, NaOH is of greatest practical importance; it is obtained:

1) electrolysis of an aqueous solution of NaCI:

2NaCl + 2H 2 OCl 2 + H 2 + 2NaOH

2) heating a solution of soda with milk of lime:

Na 2 CO 3 + Ca (OH) 2 CaCO 3 + 2NaOH.

All alkali metals react with acids to form salts:

2E + 2NS1N 2 + 2ES1.

Interact with halogens:

2Na + Cl 2 2NaCl,

and also with chalcogens:

2NaOH + H 2 SNa 2 S + 2H 2 O (neutralization reaction),

NaOH + H 2 S NaHS + H 2 O.

Alkali metals with polybasic acids form medium salts (Na 2 CO 3, KNO 3, K 2 SO 4, K 3 PO 4, etc.) and acid salts (NaHCO 3, KHSO 3, K 2 HPO 4, NaH 2 PO 4, NaHSO 4, etc.). Salts of alkali metals and weak acids (CH 3 COOH, HCN, H 2 CO 3, etc.) are hydrolyzed, their aqueous solutions have an alkaline reaction:

Alkali metal salts (with the exception of Li salts) are highly soluble in water. Of the alkali metal salts, sodium carbonate Na 2 CO 3 (soda ash) is of practical importance. It is obtained using the ammonia method:

NH 3 + H 2 O + CO 2 NH 4 HCO 3, ammonium bicarbonate

NH 4 HCO 3 + NaC? NaHCO 3 + NH 4 C?,

2NaHCO 3 Na 2 CO 3 + CO 2 + H 2 O.

The released CO 2 is returned to the process. When heated, alkali metal nitrates decompose:

4LiNO 3 2Li 2 O + 4NO 2 + O 2, 2KNO 3 2KNO 2 + O 2.

Application . Of the alkali metals, sodium is most used; it is used to produce sodium peroxide, in organic syntheses, in metallothermy, as a coolant in nuclear reactors, and for drying organic solvents. Potassium is used in metallothermy; superoxide KO 2 is obtained from potassium, used in submarines and spaceships for CO 2 absorption and oxygen regeneration:

4KO 2 + 2CO 2 2K 2 CO 3 + 3O 2.

Sodium peroxide is used for the same purpose:

2Na 2 O 2 + 2CO 2 2Na 2 CO 3 + O 2 .

Lithium is an additive to some alloys; it is used in chemical power sources to produce lithium aluminum hydride. In aviation, the construction material A1-Li is used. Cesium is used in photovoltaic cells. Alkali metal salts are widely used. NaCl is a food seasoning and preservative in food industry, and is also used in the production of soap and organic dyes. KS1 is used as fertilizer. NaOH is used for the production of artificial fiber and the purification of petroleum products. Sodium peroxide - for bleaching, disinfection. Salts Na 2 SO 4, K 2 CO 3 are used for the production of glass, KNO 3 - for the production of fertilizer, Na 2 CO 3 are used for the production of aluminum, glass, and in soap making; NaHCO 3 is used in the food industry. Li 2 O - is part of special types of glass with a low melting point.

6. Biological role of group IA elements

The biological role of lithium as a trace element has not yet been fully elucidated. It has been proven that at the level cell membranes Lithium ions (at sufficient concentration) compete with sodium ions when entering cells. The replacement of sodium ions by lithium ions in cells is associated with greater covalency of lithium compounds, as a result of which they are better soluble in phospholipids.

Sodium is the main extracellular ion. The human body contains sodium in the form of its soluble salts - chlorides, phosphates, bicarbonates. Sodium enters the human body in the form of table salt. The daily requirement for sodium is 1 g. Although the average consumption of this element is 4-7 g. Excess sodium consumption contributes to the development of hypertension. Sodium chloride is used to prepare hypertonic solutions. In case of silver nitrate poisoning, the stomach is washed with a 2-5% NaCl solution.

Sodium bicarbonate NaHCO 3 (soda) is used for diseases associated with high acidity. Sodium sulfate (Glauber's salt) NaSO 4 · 10H 2 O is used as a laxative.

Potassium is the main intracellular anion, accounting for 2/3 of the total number of active cellular anions.

Potassium ions play an important role in physiological processes - the normal functioning of the heart, muscle contraction, and the behavior of nerve impulses. Potassium is an antagonist of sodium. Potassium and sodium ions take part in biocatalysis. For potassium depletion, take potassium chloride KCl 4-5 times a day, 1 g.

Rubidium and cesium belong to microelements. A potassium synergist, rubidium activates many of the same enzymes as potassium.

Radioactive isotopes 127 Cs and 87 Rb used in radiotherapy of malignant tumors.

France - It is a radioactive chemical element produced artificially. Francium is able to selectively accumulate in tumors in the early stages of their development, which is useful in diagnosing cancer.

7. Group IIA elements

The main subgroup of group II includes the elements: beryllium, magnesium, calcium. strontium, barium and radium. All these elements, except beryllium, have pronounced metallic properties. in the free state they are silvery-white substances. More solid. Than alkali metals, with a fairly high melting point. In terms of density, all of them, except radium, belong to light metals. Their most important properties are given in Table 6. The second period element beryllium beryllium differs in its properties from other elements of this subgroup. Thus, the Be 2+ ion, due to its very small ionic radius (0.027 nm), high charge density, and high ionization energies, is stable only in the gas phase at high temperatures. Therefore, the chemical bond in binary beryllium compounds even with the most electronegative elements (BeO, BeF 2) has a high degree of covalence.

For alkaline earth metals(Ca, Sr, Ba, Ra) are characterized by the formation of ionic bonds and high coordination numbers. Magnesium occupies an intermediate position, since on the one hand it is similar to alkaline earth, predominantly ionic compounds, the formation of the 2+ ion, and in a number of properties (dissolution from oles, hydroxide basicity) - to beryllium. The degree of ionicity of the bond in salts and hydroxides is less than in alkali metal compounds. In many cases, the bonds in the crystal structure are so precise that alkali salts (sulfates, carbonates, orthophosphates) turn out to be poorly soluble.

Mg and Ca are widespread in nature, Sr and Ba are rare, Be is a rare element, Ra in insignificant quantities accompanies uranium, during the decay of which it is formed.

Elements of subgroup II A are not found in the free state (native magnesium is found in very small quantities). Mg and Ca are part of natural silicates, aluminosilicates and carbonates:

2МgОSiО 2 (olivine); MgOAI 2 O 3 (spinel); MgС1 2 6Н 2 O (bischofite); MgCO 3 (magnesite); CaCO3 (limestone, marble, chalk). CaCO 3 MgCO3 (dolomite), CaF 2 (fluorite).

Chapter 14. Chemistry of s-elements. Sodium and potassium. Magnesium and calcium

14.1. General characteristics of elements of IA and IIA groups

Group IA includes lithium, sodium, potassium, rubidium and cesium. These elements are called alkaline elements. The same group includes the artificially obtained little-studied radioactive (unstable) element francium. Sometimes hydrogen is also included in the IA group (see Chapter 10). Thus, this group includes elements from each of the 7 periods.
Group IIA includes beryllium, magnesium, calcium, strontium, barium and radium. The last four elements have a group name - alkaline earth elements.
When talking about how often atoms of a particular element are found in nature, they usually indicate its prevalence in the earth’s crust. The earth's crust refers to the atmosphere, hydrosphere and lithosphere of our planet. Thus, four of these thirteen elements are most abundant in the earth's crust: Na ( w=2.63%), K ( w= 2.41%), Mg ( w= 1.95%) and Ca ( w= 3.38%). The rest are much less common, and francium is not found at all.
The orbital radii of the atoms of these elements (except hydrogen) vary from 1.04 A (for beryllium) to 2.52 A (for cesium), that is, for all atoms they exceed 1 angstrom. This leads to the fact that all of these elements are true metal forming elements, and beryllium is an amphoteric metal forming element.
The general valence electronic formula of group IA elements is ns 1, and group IIA elements – ns 2 .
The large sizes of atoms and the small number of valence electrons lead to the fact that the atoms of these elements (except beryllium) tend to give up their valence electrons. The atoms of group IA elements give up their valence electrons most easily (see Appendix 6), while singly charged cations are formed from atoms of alkali elements, and doubly charged cations are formed from atoms of alkaline earth elements and magnesium. The oxidation state in compounds of alkaline elements is +I, and that of group IIA elements is +II.
Simple substances formed by the atoms of these elements are metals. Lithium, sodium, potassium, rubidium, cesium and francium are called alkali metals because their hydroxides are alkalis. Calcium, strontium and barium are called alkaline earth metals. The chemical activity of these substances increases as the atomic radius increases.
Of the chemical properties of these metals, the most important are their reducing properties. Alkali metals are the strongest reducing agents. Metals of Group IIA elements are also quite strong reducing agents.
All of them (except beryllium) react with water (magnesium when boiled):
2M + 2H 2 O = 2M aq+2OH aq+H2,
M + 2H 2 O = M 2 + 2OH + H 2.

In the case of magnesium, calcium and strontium, due to the low solubility of the resulting hydroxides, the reaction is accompanied by the formation of a precipitate:

M 2 + 2OH = Mg(OH) 2

Alkali metals react with most nonmetals:
2M + H 2 = 2MH (when heated),
4M + O 2 = 2M 2 O (M – Li),
2M + Cl 2 = 2MCl (under normal conditions),
2M + S = M 2 S (when heated).

Of the alkali metals, when burned in oxygen, the usual oxide forms only lithium. The remaining alkali metals form peroxides (M 2 O 2) or superoxides(MO 2 – compounds containing superoxide ion with a formal charge of –1 e).
Like alkali metals, metals of Group IIA elements react with many nonmetals, but under more severe conditions:
M + H 2 = MH 2 (when heated; except beryllium),
2M + O 2 = 2MO (under normal conditions; Be and Mg - when heated),
M + Cl 2 = MCl 2 (under normal conditions),
M + S = MS (when heated).
Unlike alkali metals, they form ordinary oxides with oxygen.
Only magnesium and beryllium react calmly with acids; other simple substances react very violently, often with explosion.
Beryllium reacts with concentrated alkali solutions:
Be + 2OH + 2H 2 O = 2 + H 2

In accordance with their position in the voltage series, only beryllium and magnesium react with salt solutions; the remaining metals in this case react with water.
Being strong reducing agents, alkali and alkaline earth metals reduce many less active metals from their compounds, for example, when heated, the following reactions occur:
4Na + MnO 2 = 2Na 2 O + Mn;
2Ca + SnO 2 = 2CaO + Sn.
Common to all alkali metals and group IIA metals industrial method production – electrolysis of molten salts.
Except beryllium oxides of all the elements under consideration are basic oxides, and hydroxides– strong bases (in beryllium these compounds are amphoteric, magnesium hydroxide is a weak base).
The strengthening of the basic properties of hydroxides with an increase in the atomic number of an element in a group is easily observed in the series of hydroxides of group IIA elements. Be(OH) 2 is an amphoteric hydroxide, Mg(OH) 2 is a weak base, Ca(OH) 2, Sr(OH) 2 and Ba(OH) 2 are strong bases, but with increasing atomic number their solubility increases, and Ba( OH) 2 can already be classified as alkalis.

SUPEROXIDES
1.Make abbreviated electronic formulas and energy diagrams of atoms of elements of groups IA and IIA. Indicate the outer and valence electrons.
2. For what reasons is hydrogen placed in group IA, and for what reasons is it placed in group VIIA?
3. Make up equations for the reactions of the following substances with excess oxygen: Li, Na, K, LiH, NaH, Li 3 N, Na 2 C 2.
4.Crystals of a certain substance consist of singly charged ions. Each ion contains 18 electrons. Make up a) the simplest formula of the substance; b) abbreviated electronic formulas of ions; c) the equation of one of the reactions for producing this substance; d) two reaction equations involving this substance.

14.2. Sodium and potassium

Sodium and potassium are the most important alkaline elements.
Simple substances, formed by these elements, are soft, fusible silvery metals, easily cut with a knife, and quickly oxidize in air. They are stored under a layer of kerosene. The melting point of sodium is 98 °C, and potassium is 64 °C.
Oxides These elements are typical basic oxides. They are very hygroscopic: absorbing water, they turn into hydroxides.
Hydroxides sodium and potassium are alkalis. These are solid, colorless crystalline substances that melt without decomposition. Like oxides, they are very hygroscopic: absorbing water, they turn into concentrated solutions. Both solid hydroxides and their concentrated solutions are very dangerous substances: if they come into contact with the skin, they cause difficult-to-heal ulcers, and inhalation of their dust leads to damage to the respiratory tract. Sodium hydroxide (trivial names - caustic soda, caustic soda) is one of the most important products of the chemical industry - it is used to create an alkaline environment in many chemical industries. Potassium hydroxide (the common name is “caustic potassium”) is used to produce other potassium compounds.
Majority medium salts sodium and potassium are thermally stable substances and decompose only at very high temperatures. With moderate heating, only salts of halogenated oxoacids, nitrates and some other compounds decompose:

NaClO 4 = NaCl + 2O 2;
8NaClO 3 = 6NaClO 4 + 2NaCl;
2NaNO 3 = 2NaNO 2 + O 2;
Na 2 = Na 2 ZnO 2 + 2H 2 O.

Acid salts less stable; when heated, they all decompose:

2NaHS = Na 2 S + H 2 S;
2NaHSO 4 = Na 2 S 2 O 7 + H 2 O;
2NaHCO 3 = Na 2 CO 3 + H 2 O + CO 2;
NaH 2 PO 4 = NaPO 3 + H 2 O;
Na 2 HPO 4 = Na 4 P 2 O 7 + H 2 O.

These elements do not form basic salts.

Of the salts, the most important is sodium chloride - table salt. This is not only necessary component food, but also raw materials for the chemical industry. From it sodium hydroxide, baking soda (NaHCO 3), soda (Na 2 CO 3) and many other sodium compounds are obtained. Potassium salts are necessary mineral fertilizers.
Almost all sodium and potassium salts are soluble, so they are available qualitative reactions on the ions of these elements not. (Qualitative reactions are chemical reactions that make it possible to detect atoms or ions of any chemical element in a compound, while proving that it is these atoms or ions that have been detected, and not some others similar to them in chemical properties. Also called reactions, allowing to detect any substance in a mixture) The presence of sodium or potassium ions in a compound can be determined by the coloring of the colorless flame when the test sample is added to it: in the case of sodium, the flame is colored yellow, and in the case of potassium - purple.

QUALITATIVE REACTIONS
Write down reaction equations characterizing the chemical properties of a) sodium, b) potassium hydroxide, c) sodium carbonate, d) sodium hydrosulfide.
Flame coloring with sodium and potassium salts

14.3. Magnesium and calcium

The simple substances magnesium and calcium are metals. Calcium quickly oxidizes in air, but magnesium is much more stable under these conditions - it oxidizes only from the surface. Calcium is stored under a layer of kerosene. The melting points of magnesium and calcium are 650 and 851 °C, respectively. Magnesium and calcium are significantly more solids than alkali metals. The low density of magnesium (1.74 g/cm3) with significant strength makes it possible to use its alloys in the aviation industry.
Both magnesium and calcium are strong reducing agents (especially when heated). They are often used to reduce other, less active metals from their oxides (magnesium in the laboratory, and calcium in industry).
Magnesium and calcium are among the few metals that react with nitrogen. When heated, they form with it the nitrides Mg 3 N 2 and Ca 3 N 2. Therefore, when burned in air, magnesium and calcium are converted into a mixture of oxides and nitrides.
Calcium reacts easily with water, but magnesium only reacts when boiled. In both cases, hydrogen is released and poorly soluble hydroxides are formed.
Oxides magnesium and calcium are ionic substances; in chemical behavior they are basic oxides. Magnesium oxide does not react with water, but calcium oxide (the trivial name is “quicklime”) reacts violently, releasing heat. The resulting calcium hydroxide is called "slaked lime" in the industry.
Hydroxide Magnesium is insoluble in water, nevertheless it is a base. Calcium hydroxide is noticeably soluble in water; its saturated solution is called “lime water”, it is an alkaline solution (changes the color of the indicators). Calcium hydroxide in a dry, and especially in a wet state, absorbs carbon dioxide from the surrounding air and turns into calcium carbonate. This property of slaked lime has been used in construction for many centuries: slaked lime as the main component was part of building lime mortars, which are now almost completely replaced by cement ones. Both hydroxides decompose when heated moderately without melting.
Salts Magnesium and especially calcium are found in many rock-forming minerals. Of these rocks, the most famous are chalk, marble and limestone, the main substance of which is calcium carbonate. When heated, calcium and magnesium carbonates decompose into the corresponding oxides and carbon dioxide. With water containing dissolved carbon dioxide, these carbonates react to form solutions of bicarbonates, for example:

MCO 3 + CO 2 + H 2 O = M 2 + 2HCO 3.

When heated, and even when trying to separate bicarbonates from solution by removing water at room temperature, they decompose by the reverse reaction:

M 2 + 2HCO 3 = MCO 3 + CO 2 + H 2 O.

Hydrated calcium sulfate CaSO 4 ·2H 2 O is a colorless crystalline substance, slightly soluble in water. When heated, it is partially dehydrated, turning into a crystalline hydrate of the composition 2CaSO 4 ·H 2 O. The trivial name for a dihydrate hydrate is gypsum, and a hemihydrate is alabaster. When alabaster is mixed with water, it hydrates, forming a dense, solid mass of gypsum. This property of alabaster is used in medicine (plaster casts) and construction (reinforced plaster partitions, sealing defects). Sculptors use alabaster to make plaster models and molds.
Calcium carbide (acetylenide) CaC 2. Structural formula (Ca2)(CC). Obtained by sintering quicklime with coal:

CaO + 3C = CaC 2 + CO

This ionic substance is not a salt and is completely hydrolyzed by water to form acetylene, which for a long time was obtained in this way:

CaC 2 + 2H 2 O = C 2 H 2 + Ca(OH) 2.

Hydrated magnesium ion 2 is a cationic acid (see Appendix 13), therefore soluble salts magnesium undergo hydrolysis. For the same reason, magnesium can form basic salts, for example Mg(OH)Cl. The hydrated calcium ion is not a cationic acid.
Calcium in the compound can be detected by the coloration of the flame. The color of the flame is orange-red. Qualitative reaction into the ions Ca 2 , Sr 2 and Ba 2 , which does not, however, allow one to distinguish these ions from each other - precipitation of the corresponding sulfates with a dilute solution of sulfuric acid (or any solution of sulfate in an acidic medium):

M 2 + SO 4 2 = MSO 4.

1.Why don’t magnesium and calcium form singly charged ions?
2. Write descriptive equations for all reactions given in the paragraph.
3. Create equations for reactions characterizing the chemical properties of a) calcium, b) calcium oxide, c) magnesium hydroxide, d) calcium carbonate, e) magnesium chloride.
Study of the properties of magnesium and calcium compounds

S-elements. Chemistry of s-elements General characteristics of s-elements Chemistry of s-elements general characteristics

General consideration.

Chemical properties of metals.

Compounds of s-metals.

Similar documents

Chapter 14. Chemistry of s-elements. Sodium and potassium. Magnesium and calcium