Class: 9

Lesson type: learning new material.

Type of lesson: combined lesson

Lesson objectives:

Tutorials: the formation of students' knowledge about alkaline earth elements as typical metals, the concept of the relationship between the structure of atoms and properties (physical and chemical).

Developing: skill development research activities, the ability to extract information from various sources, compare, generalize, draw conclusions.

Educators: nurturing a sustainable interest in the subject, nurturing such moral qualities as accuracy, discipline, independence, responsible attitude to the task assigned.

Methods: problematic, search, laboratory work, independent work students.

Equipment: computer, safety table, disk “Virtual laboratory in chemistry”, presentation .

During the classes

1. Organizational moment.

2. Introductory word of the teacher.

We study the section, metals, and you know what metals have great importance in life modern man. In previous lessons, we got acquainted with the elements of group I of the main subgroup - alkali metals. Today we are starting to study the metals of group II of the main subgroup - alkaline earth metals. In order to assimilate the material of the lesson, we need to remember the most important questions that were considered in the previous lessons.

3. Actualization of knowledge.

Conversation.

Where are the alkali metals found in periodic system DI. Mendeleev?

Student:

In the periodic system, alkali metals are located in group I of the main subgroup, on the outer level 1 electron, which alkali metals easily give away, therefore, in all compounds they exhibit an oxidation state of +1. With an increase in the size of atoms from lithium to francium, the ionization energy of atoms decreases and, as a rule, their chemical activity increases.

Teacher:

Physical Properties alkali metals?

Student:

All alkali metals are silvery-white in color with slight tints, light, soft and fusible. Their hardness and melting point naturally decrease from lithium to cesium.

Teacher:

We will check the knowledge of the chemical properties of alkali metals in the form of a small test work on the options:

• Ioption: Write the reaction equations for the interaction of sodium with oxygen, chlorine, hydrogen, water. Specify the oxidizing agent and reducing agent.
• I option: Write the reaction equations for the interaction of lithium with oxygen, chlorine, hydrogen, water. Specify the oxidizing agent and reducing agent.
• I I I option: Write the reaction equations for the interaction of potassium with oxygen, chlorine, hydrogen, water. Specify the oxidizing agent and reducing agent.

Teacher: The theme of our lesson alkaline earth metals”

Lesson objectives: To give general characteristics alkaline earth metals.

Consider their electronic structure, compare physical and chemical properties.

Learn about the most important compounds of these metals.

Determine the scope of these compounds.

Our lesson plan is written on the board, we will work according to the plan, look at the presentation.

1. The position of metals in the periodic system D.I. Mendeleev.
2. The structure of the alkali metal atom.
3. physical properties.
4. Chemical properties.
5. The use of alkaline earth metals.

Conversation.

Teacher:

Based on the knowledge gained earlier, we will answer the following questions: To answer, we will use the periodic system chemical elements DI. Mendeleev.

1. List the alkaline earth metals

Student:

These are magnesium, calcium, strontium, barium, radium.

Teacher:

2. Why are these metals called alkaline earth?

Student:

The origin of this name is due to the fact that their hydroxides are alkalis, and their oxides are similar in refractoriness to oxides of aluminum and iron, which previously bore the common name "earth"

Teacher:

3. Location of alkaline earth metals in PSCE D.I. Mendeleev.

Student:

Group II is the main subgroup. Metals of group II of the main subgroup have 2 electrons at the external energy level, located at a smaller distance from the nucleus than alkali metals. Therefore, their reducing properties, although great, are still less than those of the elements of group I. Gain reducing properties also observed during the transition from Mg to Ba, which is associated with an increase in the radii of their atoms, in all compounds they exhibit an oxidation state of +2.

Teacher: Physical properties of alkaline earth metals?

Student:

Metals of group II of the main subgroup are silvery-white substances that conduct heat well and electricity. Their density increases from Be to Ba, while the melting point, on the contrary, decreases. They are much harder than alkali metals. All, except beryllium, have the ability to color the flame in different colors.

Problem: How are alkaline earth metals found in nature?

Why do alkaline earth metals mostly exist in nature in the form of compounds?

Answer: In nature, alkaline earth metals are in the form of compounds, because they have high chemical activity, which in turn depends on the characteristics electronic structure atoms (the presence of two unpaired electrons in the outer energy level)

Fizkultminutka - rest for the eyes.

Teacher:

Knowing the general physical properties, the activity of metals, assume the chemical properties of alkaline earth metals. What substances do alkali metals interact with?

Student:

Alkaline earth metals interact with both simple and complex substances. They actively interact with almost all non-metals (with halogens, hydrogen, forming hydrides). From complex substances with water - forming water-soluble bases - alkalis and with acids.

Teacher:

And now, in experiments, we will verify the correctness of our assumptions about the chemical properties of alkaline earth metals.

4. Laboratory work on the virtual laboratory.

Target: carry out reactions confirming the chemical properties of alkaline earth metals.

We repeat the safety rules for working with alkaline earth metals.

• work in a fume hood
• on a tray
• with dry hands
• take in small quantities

We work with the text that we read in the virtual laboratory.

Experience No. 1. Interaction of calcium with water.

Experience number 2. Combustion of magnesium, calcium, strontium, barium

Write down the reaction and observation equations in a notebook.

5. Summing up the lesson, grading.

5. Reflection.

What do you remember about the lesson, what did you like?

6. Homework.

§ 12 exercise 1(b) exercise 4

Literature.

1. Rudzitis G.E., Feldman F.G. Chemistry 9.- Moscow.: Education, 2001
2. Gabrielyan O.S. Chemistry 9.-Moscow.: Bustard, 2008
3. Gabrielyan O.S., Ostroumov I.G. Handbook of the teacher. Chemistry 9.-Moscow.: Bustard 2002
4. Gabrielyan O.S. Control and verification work. Chemistry 9.-Moscow.: Bustard, 2005.
5. Collection Virtual Lab. Educational electronic edition

All elements of the main subgroups of groups I and II of the Periodic system, as well as hydrogen and helium, belong to the s-elements. In addition to hydrogen and helium, all these elements are metals. Group I metals of the Periodic Table are called alkaline, because they react with water to form alkalis. Group II metals of the Periodic Table, with the exception of beryllium and magnesium, are called alkaline earth. Francium, which completes group I, and radium, which completes group II, - radioactive elements.

Some properties of s-metals 3

Table 15.1

	Metal radius, nm	Ionic radius, nm		EO according to Pauling
I group
11 group

and PI is the potential (energy) of ionization; EO - electronegativity.

All s-metals have one or two electrons on the outer shell and can easily donate them, forming ions with a stable electronic configuration of noble gases. The high reducing activity of these metals is manifested in very low ionization potentials (PI) and low electronegativity (EO) (Table 15.1). Compare the ionization potentials of alkali metals and noble gases (of all elements, noble gases have the lowest EO and highest PI; see Table 18.1).

physical properties. Under normal conditions, s-metals are in a solid state, forming crystals with metallic bond. All Group I metals have body-centered cubic lattice(BCC, see § 4.4). Beryllium and magnesium are characterized hexagonal close packing(hcp), in calcium and strontium face-centered cubic lattice (fcc), for barium body-centered cubic(BCC).

Group I metals are soft and have a low density compared to others. Lithium, sodium and potassium lighter than water and float on its surface, reacting with it. Group II metals are harder and denser than alkali metals. The low melting and boiling points of s-metals (see Table 15.1) are explained by the relatively weak metallic bond in the crystal lattices; binding energy (in eV): lithium 1.65, sodium 1.11, potassium 0.92, rubidium 0.84, cesium 0.79, beryllium 3.36, magnesium 1.53, calcium 1.85, strontium 1, 70, barium 1.87.

For comparison of binding energies (in eV): aluminum 3.38, zinc 1.35, iron 4.31, copper 3.51, silver 2.94, titanium 4.87, molybdenum 6.82, tungsten 8.80.

A metallic bond is formed by delocalized valence electrons holding the positive ions of metal atoms together (see § 3.6). The larger the metallic radius, the more delocalized electrons that are distributed in a "thin layer" between positive ions, and the less strength of the crystal lattice. This explains the low melting and boiling points of metals of groups I and II. The melting and boiling points of group II elements, in contrast to alkali metals, change non-systematically, which is explained by differences in crystal structures (see above).

distribution in nature. All s-metals are found in nature only in the form of compounds: fossil mineral salts and their deposits (KS1, NaCl, CaCO 3 and others) and ions in sea ​​water. Calcium, sodium, potassium and magnesium are fifth, sixth, seventh and eighth in abundance on Earth, respectively. Strontium is common in moderation. The content of other s-metals in earth's crust and oceanic waters is negligible. For example, the content of sodium in the earth's crust is 2.3% and sea water is 1.1%, cesium in the earth's crust is 3 10 ~ 4% and in sea water 3 10 -8%.

Sodium, cesium and beryllium have only one stable isotope each, lithium, potassium and rubidium each have two: |Li 7.5% and |Li 92.5%; 93.26% and Central Committee 6.74%; f^Rb 72.17% and fpRb 27.83%. Magnesium has three stable isotopes (| 2 Mg 79.0%, j|Mg 10.0% and j|Mg 11.0%). Other alkaline earth metals have more stable isotopes; the main ones: 4 °С 96.94% and CA 2.09%; ||Sr 82.58%, 8 |Sr 9.86% and ||Sr 7.0%; 1 ||Ba 71.7%, 18 |Ba 11.23%, 18 ®Ba 7.85% and 18 |Ba 6.59%.

The main subgroup of the second group of the periodic system covers the elements: beryllium, magnesium, calcium, strontium, barium and radium. According to the main representatives of this subgroup - calcium, strontium and barium - known under the general name of alkaline earth metals, the entire main subgroup of the second group is also called the subgroup alkaline earth metals.

The name "alkaline earth" these metals (sometimes magnesium is added to them) was received because their oxides, in their chemical properties are intermediate, on the one hand, between alkalis, i.e. oxides or hydroxides of alkali metals and, on the other hand, "earths", i.e. oxides of such elements, a typical representative of which is aluminum - the main component clay Owing to this intermediate position, the oxides of calcium, strontium and barium were given the name "alkaline earths".

The first element of this subgroup, beryllium (if its valency is not taken into account), is much closer in its properties to aluminum than to the higher analogues of the top group to which it belongs. The second element of this group, magnesium, is also in some respects significantly different from the alkaline earth metals in the narrow sense of the term. Some reactions bring it closer to the elements of the secondary subgroup of the second group, especially zinc; thus, sulfates of magnesium and zinc, in contrast to sulfates of alkaline earth metals, are easily soluble, isomorphic to each other and form double salts similar in composition. Previously, the rule was indicated, according to which the first element reveals properties that are transitional to the next main subgroup, the second - to a secondary subgroup of the same group; and usually only the third element possesses the properties characteristic of the group; this rule is especially evident in the group of alkaline earth metals.

The heaviest of the elements of the second group - radium - in terms of its chemical properties, of course, corresponds to typical representatives of the alkaline earth metals. However, it is usually not customary to include it in the group of alkaline earth metals in a narrower sense. In connection with the peculiarities of its distribution in nature, and also due to its most characteristic property - radioactivity, it is more expedient to give it a special place. In discussion common properties elements of this subgroup, radium will not be considered, since the corresponding physicochemical properties have not yet been studied enough.

With the exception of radium, all elements of the alkaline earth subgroup are light metals. Light metals are called, the specific gravity of which does not exceed 5. In terms of hardness, the metals of the main subgroup of group II are significantly superior to alkali metals. The softest of these, barium (whose properties are closest to the alkali metals) has approximately the hardness of lead. The melting points of this group of metals are much higher than those of the alkali metals.

Common to all elements of the main subgroup of group II is their property to show a positive valence 2 in their compounds, and only in completely exceptional cases are they positively monovalent. The valency 2+ typical for them, as well as the serial numbers of the elements, undoubtedly force us to attribute these metals to the main subgroup of the second group. In addition, they all show a strongly electropositive character, which is determined by their position on the left side of the electrochemical voltage series, as well as by their strong affinity for electronegative elements.

In accordance with the value of the normal potentials of the elements of the main subgroup of the second group, all of the listed metals decompose water; however, the action of beryllium and magnesium on water proceeds very slowly due to the low solubility of the hydroxides resulting from this reaction, for example, for magnesium:

Mg + 2HOH \u003d Mg (OH) 2 + H 2

Formed on the surface of the metal, Be and Mg hydroxides hinder the further course of the reaction. Therefore, even small errors of magnesium have to be kept at ordinary temperature in contact with water for several days before they completely turn into magnesium hydroxide. The remaining alkaline earth metals react with water much more vigorously, which is explained by the better solubility of their hydroxides. Barium hydroxide is the easiest to dissolve; the normal potential Ba has the lowest value in comparison with other elements of the group, therefore it reacts with water, as well as with alcohol, very vigorously. The resistance of alkaline earth metals to air decreases in the direction from magnesium to barium. In accordance with their position in the series of stresses, the named metals displace all heavy metals from solutions of their salts.

Normal M II O oxides are always obtained as products of combustion of alkaline earth metals. Peroxides of alkaline earth metals are much less stable than those of the alkali metal series.

Alkaline earth metal oxides combine with water to form hydroxides. moreover, the energy of this reaction very noticeably increases in the direction from BeO to BaO. The solubility of hydroxides also greatly increases from beryllium hydroxide and barium hydroxide; but even the solubility of the latter at normal temperature is very low. In the same order, the basic nature of these compounds also increases - from amphoteric beryllium hydroxide to strongly basic caustic barium.

It is interesting to note the strong affinity of the elements of the main subgroup of the second group for nitrogen. The propensity to form compounds with nitrogen increases in these elements with an increase in atomic weight (despite the fact that the heats of formation of nitrides decrease in this direction); in the alkaline earth metals proper, the tendency to form nitrides is so great that the latter are slowly combined with nitrogen even at ordinary nitride.

alkaline earth metals like alkali metals, they combine with hydrogen to form hydrides, for example:

Ca + H 2 \u003d CaH 2.

Ethn hydrides also have a salt-like character, and therefore it should be assumed that in them, as in alkali metal hydrides, hydrogen is an electronegative component.

It is more difficult to obtain MgH 2 directly from the elements, but it was generally not possible to synthesize BeH 2 in this way. MgH 2 and BeH 2 are solid and non-volatile compounds, like alkaline earth metal hydrides, but unlike the latter, they do not have a pronounced salt-like character.

All elements of the main subgroup of the second group form colorless ions with a positive charge of 2: Be 2+, Mg 2+, Ca 2+, Sr 2+, Ba 2+, Ra 2+. Beryllium also forms colorless anions [BeO 2 ] 2+ and [Be(OH) 4 ] 2+ . Colorless and all salts M II X 2 of these elements, if they are not derivatives of colored anions.

Radium salts themselves are also colorless. However, some of them, such as radium chloride and bromide, are gradually colored by the radiation of the radium contained in them, and finally acquire a color from brown to black. Upon recrystallization, they turn white again.

Many alkaline earth metal salts are sparingly soluble in water. A certain pattern is often found in the change in the solubility of these salts: for example, in sulfates, the solubility decreases rapidly with an increase in the atomic weight of the alkaline earth metal. The solubility of chromites also changes approximately in the same way. Most of the salts formed by alkaline earth metals with weak acids and with acids of medium strength are difficult to dissolve, such as phosphates, oxalates and carbonates; some of them, however, are easily soluble; the latter include sulfides, cyanides, thiocyanates and acetates. Due to the weakening of the basic character of hydroxides upon passing from Ba to Be, the degree of hydrolysis of their carbonates increases in the same sequence. Their thermal stability also changes in the same direction: while barium carbonate decomposes far from completely even at a temperature of white heat, calcium carbonate can be completely decomposed into CaO and CO 2 even with relatively weak calcination, and magnesium carbonate decomposes even more easily.

From the point of view of Kossel's theory, the reason for the divalence of the elements of the alkaline earth group is the fact that in the periodic table they are all removed from the corresponding inert gases with: 2 elements, therefore each of them has 2 electrons more than the previous inert gas. Due to the tendency of atoms to take on the configuration of inert gases in the elements of the alkaline earth group, a slight splitting off of two electrons occurs, but no more, since further splitting off would have already caused the destruction of the configuration of inert gases.

The most active among the metal group are the alkali and alkaline earth metals. These are soft light metals that react with simple and complex substances.

general description

Active metals occupy the first and second groups periodic table Mendeleev. Full list alkali and alkaline earth metals:

• lithium (Li);
• sodium (Na);
• potassium (K);
• rubidium (Rb);
• cesium (Cs);
• francium (Fr);
• beryllium (Be);
• magnesium (Mg);
• calcium (Ca);
• strontium (Sr);
• barium (Ba);
• radium (Ra).

Rice. 1. Alkali and alkaline earth metals in the periodic table.

Electronic configuration of alkali metals - ns 1 , alkaline earth metals - ns 2 .

Accordingly, the constant valence of alkali metals is I, alkaline earth - II. Due to the small number of valence electrons in the outer energy level active metals exhibit powerful reducing agent properties by donating outer electrons in reactions. The more energy levels, the less the bond between the outer electrons and the atomic nucleus. Therefore, the metallic properties increase in groups from top to bottom.

Due to the activity, metals of groups I and II are found in nature only in the composition rocks. Pure metals are isolated by electrolysis, calcination, substitution reactions.

Physical Properties

Alkali metals have a silvery-white color with a metallic sheen. Cesium is a silvery yellow metal. These are the most active and soft metals. Sodium, potassium, rubidium, cesium are cut with a knife. The softness is like wax.

Rice. 2. Cutting sodium with a knife.

Alkaline earth metals are gray in color. Compared to alkali metals, they are harder, denser substances. Only strontium can be cut with a knife. The densest metal is radium (5.5 g/cm3).

The lightest metals are lithium, sodium and potassium. They float on the surface of the water.

Chemical properties

Alkali and alkaline earth metals react with simple substances and complex compounds, forming salts, oxides, alkalis. The main properties of active metals are described in the table.

Interaction	alkali metals	alkaline earth metals
With oxygen	Self-igniting in air. They form superoxides (RO 2), except for lithium and sodium. Lithium forms an oxide when heated above 200°C. Sodium forms a mixture of peroxide and oxide. 4Li + O 2 → 2Li 2 O; 2Na + O 2 → Na 2 O 2; Rb + O 2 → RbO 2	Protective oxide films quickly form in air. When heated to 500 ° C, they spontaneously ignite. 2Mg + O 2 → 2MgO; 2Ca + O 2 → 2CaO
with non-metals	React when heated with sulfur, hydrogen, phosphorus: 2K + S → K 2 S; 2Na + H 2 → 2NaH; 2Cs + 5P → Cs 2 P 5 . Only lithium reacts with nitrogen, lithium and sodium react with carbon: 6Li + N 2 → 2Li 3 N; 2Na + 2C → Li 2 C 2	React when heated: Ca + Br 2 → CaBr 2; Be + Cl 2 → BeCl 2; Mg + S → MgS; 3Ca + 2P → Ca 3 P 2 ; Sr + H2 → SrH2
With halogens	React violently to form halides: 2Na + Cl 2 → 2NaCl
	Alkalis are formed. The lower the metal is located in the group, the more actively the reaction proceeds. Lithium interacts calmly, sodium burns with a yellow flame, potassium with a flash, cesium and rubidium explode. 2Na + 2H 2 O → 2NaOH + H 2 -; 2Li + 2H 2 O → 2LiOH + H 2	Less active than alkali metals, react at room temperature: Mg + 2H 2 O → Mg (OH) 2 + H 2; Ca + 2H 2 O → Ca (OH) 2 + H 2
With acids	With weak and dilute acids they react explosively. They form salts with organic acids. 8K + 10HNO 3 (conc) → 8KNO 3 + N 2 O + 5H 2 O; 8Na + 5H 2 SO 4 (conc) → 4Na 2 SO 4 + H 2 S + 4H 2 O; 10Na + 12HNO 3 (diff) → N 2 + 10NaNO 3 + 6H 2 O; 2Na + 2CH 3 COOH → 2CH 3 COONa + H 2	Form salts: 4Sr + 5HNO 3 (conc) → 4Sr(NO 3) 2 + N 2 O + 4H 2 O; 4Ca + 10H 2 SO 4 (conc) → 4CaSO 4 + H 2 S + 5H 2 O
With alkalis		Of all the metals, only beryllium reacts: Be + 2NaOH + 2H 2 O → Na 2 + H 2
With oxides		All metals react with the exception of beryllium. Replace less active metals: 2Mg + ZrO 2 → Zr + 2MgO

Rice. 3. Reaction of potassium with water.

Alkali and alkaline earth metals can be detected using a qualitative reaction. When burning, metals are painted in a certain color. For example, sodium burns with a yellow flame, potassium with violet, barium with light green, and calcium with dark orange.

What have we learned?

Alkaline and alkaline earth are the most active metals. It's soft simple substances gray or silvery color with a small density. Lithium, sodium, potassium float on the surface of the water. Alkaline earth metals are harder and denser than alkali metals. They oxidize quickly in air. Alkali metals form superoxides and peroxides, oxide forms only lithium. React violently with water at room temperature. React with non-metals when heated. Alkaline earth metals react with oxides, displacing less active metals. Only beryllium reacts with alkalis.

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To the family alkaline earth elements include calcium, strontium, barium and radium. D. I. Mendeleev included magnesium in this family. Alkaline earth elements are named for the reason that their hydroxides, like alkali metal hydroxides, are soluble in water, that is, they are alkalis. “... They are called earthy because in nature they are found in the state of compounds that form an insoluble mass of the earth, and themselves, in the form of RO oxides, have an earthy appearance,” Mendeleev explained in Fundamentals of Chemistry.

General characteristics of the elements of group IIa

Metals of the main subgroup of group II have an electronic configuration of the external energy level ns², and are s-elements.

Easily donate two valence electrons, and in all compounds they have an oxidation state of +2

Strong reducing agents

The activity of metals and their reducing ability increases in the series: Be–Mg–Ca–Sr–Ba

Alkaline earth metals include only calcium, strontium, barium and radium, less often magnesium

Beryllium is closer to aluminum in most properties.

Physical properties of simple substances

Alkaline earth metals (compared to alkali metals) have higher t°pl. and t ° boiling., ionization potentials, densities and hardness.

Chemical properties of alkaline earth metals + Be

1. Reaction with water.

Under normal conditions, the surface of Be and Mg is covered with an inert oxide film, so they are resistant to water. In contrast, Ca, Sr and Ba dissolve in water with the formation of alkalis:

Mg + 2H 2 O - t ° → Mg (OH) 2 + H 2

Ca + 2H 2 O → Ca (OH) 2 + H 2

2. Reaction with oxygen.

All metals form oxides RO, barium peroxide - BaO 2:

2Mg + O 2 → 2MgO

Ba + O 2 → BaO 2

3. Form binary compounds with other non-metals:

Be + Cl 2 → BeCl 2 (halides)

Ba + S → BaS (sulfides)

3Mg + N 2 → Mg 3 N 2 (nitrides)

Ca + H 2 → CaH 2 (hydrides)

Ca + 2C → CaC 2 (carbides)

3Ba + 2P → Ba 3 P 2 (phosphides)

Beryllium and magnesium react relatively slowly with non-metals.

4. All alkaline earth metals dissolve in acids:

Ca + 2HCl → CaCl 2 + H 2

Mg + H 2 SO 4 (dec.) → MgSO 4 + H 2

5. Beryllium dissolves in aqueous solutions alkalis:

Be + 2NaOH + 2H 2 O → Na 2 + H 2

6. Volatile compounds of alkaline earth metals give the flame a characteristic color:

calcium compounds - brick red, strontium - carmine red, and barium - yellowish green.

Beryllium, like lithium, is an s-element. The fourth electron that appears in the Be atom is placed in the 2s orbital. The ionization energy of beryllium is higher than that of lithium due to the larger nuclear charge. In strong bases, it forms the BeO 2-2 beryllate ion. Therefore, beryllium is a metal, but its compounds are amphoteric. Beryllium, although a metal, is much less electropositive than lithium.

The high ionization energy of the beryllium atom differs markedly from other elements of the PA subgroup (magnesium and alkaline earth metals). Its chemistry is largely similar to that of aluminum (diagonal similarity). Thus, this is an element with the presence of amphoteric qualities in its compounds, among which the basic ones still prevail.

The electronic configuration of Mg: 1s 2 2s 2 2p 6 3s 2 has one significant feature compared to sodium: the twelfth electron is placed in the 2s orbital, where there is already 1e - .

Magnesium and calcium ions are indispensable elements of the vital activity of any cell. Their ratio in the body must be strictly defined. Magnesium ions are involved in the activity of enzymes (for example, carboxylase), calcium - in the construction of the skeleton and metabolism. Increasing the calcium content improves the absorption of food. Calcium excites and regulates the work of the heart. Its excess sharply increases the activity of the heart. Magnesium plays part of the role of a calcium antagonist. The introduction of Mg 2+ ions under the skin causes anesthesia without a period of excitement, paralysis of muscles, nerves and heart. Getting into the wound in the form of metal, it causes long-term non-healing purulent processes. Magnesium oxide in the lungs causes the so-called foundry fever. Frequent contact of the skin surface with its compounds leads to dermatitis. The most widely used calcium salts in medicine are CaSO 4 sulfate and CaCL 2 chloride. The first is used for plaster casts, and the second is used for intravenous infusions and as an internal remedy. It helps fight swelling, inflammation, allergies, relieves spasms of the cardiovascular system, improves blood clotting.

All barium compounds except BaSO 4 are poisonous. Cause mengoencephalitis with damage to the cerebellum, damage to smooth cardiac muscles, paralysis, and in large doses - degenerative changes in the liver. In small doses, barium compounds stimulate the activity of the bone marrow.

When strontium compounds are introduced into the stomach, its disorder, paralysis, and vomiting occur; lesions are similar in signs to lesions from barium salts, but strontium salts are less toxic. Of particular concern is the appearance in the body of the radioactive isotope of strontium 90 Sr. It is extremely slowly excreted from the body, and its long half-life and, therefore, the duration of action can cause radiation sickness.

Radium is dangerous for the body with its radiation and a huge half-life (T 1/2 = 1617 years). Initially, after the discovery and production of radium salts in a more or less pure form, it began to be used quite widely for fluoroscopy, the treatment of tumors, and some serious diseases. Now, with the advent of other more accessible and cheaper materials, the use of radium in medicine has practically ceased. In some cases, it is used to produce radon and as an additive to mineral fertilizers.

The filling of the 4s orbital is completed in the calcium atom. Together with potassium, it forms a pair of s-elements of the fourth period. Calcium hydroxide is quite strong base. In calcium - the least active of all alkaline earth metals - the nature of the bond in the compounds is ionic.

According to its characteristics, strontium occupies an intermediate position between calcium and barium.

The properties of barium are closest to those of alkali metals.

Beryllium and magnesium are widely used in alloys. Beryllium bronzes are elastic copper alloys with 0.5-3% beryllium; aviation alloys (density 1.8) contain 85-90% magnesium ("electron"). Beryllium differs from other metals of group IIA - it does not react with hydrogen and water, but it dissolves in alkalis, since it forms amphoteric hydroxide:

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

Magnesium actively reacts with nitrogen:

3 Mg + N 2 \u003d Mg 3 N 2.

The table shows the solubility of hydroxides of elements of group II.

Traditional technical problem - hardness of water associated with the presence of Mg 2+ and Ca 2+ ions in it. From bicarbonates and sulfates on the walls of heating boilers and pipes with hot water magnesium and calcium carbonates and calcium sulfate precipitate. They especially interfere with the work of laboratory distillers.

S-elements in a living organism perform an important biological function. The table shows their content.

The extracellular fluid contains 5 times more sodium ions than inside the cells. An isotonic solution (“physiological fluid”) contains 0.9% sodium chloride, it is used for injections, washing wounds and eyes, etc. Hypertonic solutions (3-10% sodium chloride) are used as lotions in the treatment of purulent wounds (“stretching » pus). 98% of potassium ions in the body are inside the cells and only 2% in the extracellular fluid. A person needs 2.5-5 g of potassium per day. 100 g of dried apricots contain up to 2 g of potassium. In 100 g of fried potatoes - up to 0.5 g of potassium. In intracellular enzymatic ATP reactions and ADP participate in the form of magnesium complexes.

Every day a person needs 300-400 mg of magnesium. It enters the body with bread (90 mg of magnesium per 100 g of bread), cereals (in 100 g of oatmeal up to 115 mg of magnesium), nuts (up to 230 mg of magnesium per 100 g of nuts). In addition to building bones and teeth based on hydroxyapatite Ca 10 (PO 4) 6 (OH) 2, calcium cations are actively involved in blood clotting, transmission of nerve impulses, and muscle contraction. Adults need to consume about 1 g of calcium per day. 100 g of hard cheeses contain 750 mg of calcium; in 100 g of milk - 120 mg of calcium; in 100 g of cabbage - up to 50 mg.