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 of 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 illnesses. 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 a fairly 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.

Properties of alkaline earth metals

Physical properties

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

Chemical properties

1. Very reactive.

2. Have a positive valence of +2.

3. React with water at room temperature (except for Be) with evolution of hydrogen.

4. They have a high affinity for oxygen (reducing agents).

5. They form salt-like hydrides EH 2 with hydrogen.

6. Oxides have the general formula EO. The tendency towards the formation of peroxides is less pronounced than for alkali metals.

Being in nature

3BeO ∙ Al 2 O 3 ∙ 6SiO 2 beryl

mg

MgCO 3 magnesite

CaCO 3 ∙ MgCO 3 dolomite

KCl ∙ MgSO 4 ∙ 3H 2 O kainite

KCl ∙ MgCl 2 ∙ 6H 2 O carnallite

CaCO 3 calcite (limestone, marble, etc.)

Ca 3 (PO 4) 2 apatite, phosphorite

CaSO 4 ∙ 2H 2 O gypsum

CaSO 4 anhydrite

CaF 2 fluorspar (fluorite)

SrSO 4 celestine

SrCO 3 strontianite

BaSO 4 barite

BaCO 3 witherite

Receipt

Beryllium is obtained by reduction of fluoride:

BeF 2 + Mg═ t ═ Be + MgF 2

Barium is obtained by oxide reduction:

3BaO + 2Al═ t ═ 3Ba + Al 2 O 3

The remaining metals are obtained by electrolysis of chloride melts:

CaCl 2 \u003d Ca + Cl 2 ╜

cathode: Ca 2+ + 2ē = Ca 0

anode: 2Cl - - 2ē = Cl 0 2

MgO + C = Mg + CO

Metals of the main subgroup of group II are strong reducing agents; in compounds, they exhibit only the +2 oxidation state. The activity of metals and their reducing ability increases in the series: Be Mg Ca Sr Ba╝

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 to form hydroxides, which are strong bases:

Mg + 2H 2 O═ t ═ Mg (OH) 2 + H 2

Ca + 2H 2 O \u003d Ca (OH) 2 + H 2 ╜

2. Reaction with oxygen.

All metals form oxides RO, barium peroxide BaO 2:

2Mg + O 2 \u003d 2MgO

Ba + O 2 \u003d BaO 2

3. Binary compounds are formed with other non-metals:

Be + Cl 2 = BeCl 2 (halides)

Ba + S = BaS (sulfides)

3Mg + N 2 \u003d 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 metals dissolve in acids:

Ca + 2HCl \u003d CaCl 2 + H 2 ╜

Mg + H 2 SO 4 (razb.) \u003d MgSO 4 + H 2 ╜

Beryllium also dissolves in aqueous solutions of alkalis:

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

5. Qualitative reaction to alkaline earth metal cations - coloring of the flame in the following colors:

Ca 2+ - dark orange

Sr 2+ - dark red

Ba 2+ - light green

The Ba 2+ cation is usually opened by an exchange reaction with sulfuric acid or its salts:

Barium sulfate is a white precipitate, insoluble in mineral acids.

Alkaline earth metal oxides

Receipt

1) Oxidation of metals (except Ba, which forms a peroxide)

2) Thermal decomposition of nitrates or carbonates

CaCO 3 ═ t ═ CaO + CO 2 ╜

2Mg(NO 3) 2 ═ t ═ 2MgO + 4NO 2 ╜ + O 2 ╜

Chemical properties

Typical basic oxides. React with water (except BeO), acid oxides and acids

MgO + H 2 O \u003d Mg (OH) 2

3CaO + P 2 O 5 \u003d Ca 3 (PO 4) 2

BeO + 2HNO 3 \u003d Be (NO 3) 2 + H 2 O

BeO- amphoteric oxide, soluble in alkalis:

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

Alkaline earth metal hydroxides R(OH) 2

Receipt

Reactions of alkaline earth metals or their oxides with water: Ba + 2H 2 O \u003d Ba (OH) 2 + H 2

CaO (quicklime) + H 2 O \u003d Ca (OH) 2 (slaked lime)

Chemical properties

Hydroxides R(OH) 2 - white crystalline substances, are less soluble in water than alkali metal hydroxides (the solubility of hydroxides decreases with decreasing serial number; Be (OH) 2 is insoluble in water, soluble in alkalis). The basicity of R(OH) 2 increases with increasing atomic number:

Be(OH) 2 - amphoteric hydroxide

Mg(OH) 2 - weak base

other hydroxides - strong bases(alkali).

1) Reactions with acid oxides:

Ca(OH) 2 + SO 2 = CaSO 3 ¯ + H 2 O

Ba(OH) 2 + CO 2 = BaCO 3 ¯ + H 2 O

2) Reactions with acids:

Mg(OH) 2 + 2CH 3 COOH = (CH 3 COO) 2 Mg + 2H 2 O

Ba(OH) 2 + 2HNO 3 = Ba(NO 3) 2 + 2H 2 O

3) Exchange reactions with salts:

Ba(OH) 2 + K 2 SO 4 = BaSO 4 ¯+ 2KOH

4) The reaction of beryllium hydroxide with alkalis:

Be(OH) 2 + 2NaOH = Na 2

Hardness of water

Natural water containing Ca 2+ and Mg 2+ ions is called hard. Hard water when boiled forms a scale, it does not boil soft food products; detergents do not produce foam.

Carbonate (temporary) hardness is due to the presence of calcium and magnesium bicarbonates in water, non-carbonate (permanent) hardness - chlorides and sulfates.

The total hardness of water is considered as the sum of carbonate and non-carbonate.

Removal of water hardness is carried out by precipitation of Ca 2+ and Mg 2+ ions from the solution:

1) boiling:

Ca(HCO 3) 2 ═ t ═ CaCO 3 ¯ + CO 2 + H 2 O

Mg(HCO 3) 2 ═ t═ MgCO 3 ¯ + CO 2 + H 2 O

2) by adding milk of lime:

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

3) adding soda:

Ca(HCO 3) 2 + Na 2 CO 3 \u003d CaCO 3 ¯+ 2NaHCO 3

CaSO 4 + Na 2 CO 3 \u003d CaCO 3 ¯ + Na 2 SO 4

MgCl 2 + Na 2 CO 3 \u003d MgCO 3 ¯ + 2NaCl

All four methods are used to remove temporary stiffness, and only the last two are used for permanent hardness.

Thermal decomposition of nitrates.

E (NO3) 2 \u003d t \u003d EO + 2NO2 + 1 / 2O2

Features of the chemistry of beryllium.

Be(OH)2 + 2NaOH (g) = Na2

Al(OH)3 + 3NaOH (g) = Na3

Be + 2NaOH + 2H2O = Na2 + H2

Al + 3NaOH + 3H2O = Na3 + 3/2H2

Be, Al + HNO3 (Conc) = passivation

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, " lands", 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 the discussion of the general properties of the elements of this subgroup, radium will not be considered, since the corresponding physicochemical properties have not yet been sufficiently studied.

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 their 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 voltages, these 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 already 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 lesson will cover the topic “Metals and their properties. alkali metals. alkaline earth metals. Aluminum". You will learn general properties and patterns of alkali and alkaline earth elements, study separately the chemical properties of alkali and alkaline earth metals and their compounds. By using chemical equations the concept of water hardness will be considered. Get to know aluminum, its properties and alloys. You will learn what oxygen regenerating mixtures, ozonides, barium peroxide and oxygen production are.

Topic: Basic metals and non-metals

Lesson: Metals and their properties. alkali metals. alkaline earth metals. Aluminum

The main subgroup of group I Periodic system DI. Mendeleev are lithium Li, sodium Na, potassium K, rubidium Rb, cesium Cs and francium Fr. Elements of this subgroup belong to. Their common name is alkali metals.

Alkaline earth metals are in the main subgroup of group II of the Periodic Table of D.I. Mendeleev. These are magnesium Mg, calcium Ca, strontium Sr, barium Ba and radium Ra.

Alkali and alkaline earth metals, as typical metals, exhibit pronounced reducing properties. For elements of the main subgroups, the metallic properties increase with increasing radius. Especially strong reducing properties are manifested in alkali metals. So strong that it is practically impossible to carry out their reactions with dilute aqueous solutions, since the first reaction will be their interaction with water. The situation is similar for alkaline earth metals. They also interact with water, but much less intensely than alkali metals.

Electronic configurations valence layer of alkali metals - ns 1 , where n is the number of the electronic layer. They are referred to as s-elements. For alkaline earth metals - ns 2 (s-elements). Aluminum has valence electrons …3 s 2 3r 1(p-element). These elements form compounds with ion type connections. In the formation of compounds for them, the oxidation state corresponds to the group number.

Detection of metal ions in salts

Metal ions are easily identified by the color change of the flame. Rice. one.

Lithium salts - carmine-red flame color. Sodium salts - yellow. Potassium salts - violet through cobalt glass. Rubidium - red, cesium - violet-blue.

Rice. one

Salts of alkaline earth metals: calcium - brick red, strontium - carmine red and barium - yellowish green. Aluminum salts do not change the color of the flame. Salts of alkali and alkaline earth metals are used to create fireworks. And you can easily determine by color, which metal salts were used.

Metal properties

alkali metals are silvery-white substances with a characteristic metallic luster. They quickly tarnish in air due to oxidation. These are soft metals, Na, K, Rb, Cs are similar in softness to wax. They are easily cut with a knife. They are light. Lithium is the lightest metal with a density of 0.5 g/cm3.

Chemical properties alkali metals

1. Interaction with non-metals

Due to high reducing properties alkali metals react violently with halogens to form the corresponding halide. When heated, they react with sulfur, phosphorus and hydrogen to form sulfides, hydrides, and phosphides.

2Na + Cl 2 → 2NaCl

Lithium is the only metal that reacts with nitrogen already at room temperature.

6Li + N 2 = 2Li 3 N, the resulting lithium nitride undergoes irreversible hydrolysis.

Li 3 N + 3H 2 O → 3LiOH + NH 3

2. Interaction with oxygen

Lithium oxide is formed immediately with lithium.

4Li + O 2 \u003d 2Li 2 O, and when oxygen reacts with sodium, sodium peroxide is formed.

2Na + O 2 \u003d Na 2 O 2. When all other metals are burned, superoxides are formed.

K + O 2 \u003d KO 2

3. Interaction with water

By reacting with water, one can clearly see how the activity of these metals in the group changes from top to bottom. Lithium and sodium calmly interact with water, potassium - with a flash, and cesium - already with an explosion.

2Li + 2H 2 O → 2LiOH + H 2

4.

8K + 10HNO 3 (conc) → 8KNO 3 + N 2 O +5 H 2 O

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

Obtaining alkali metals

Due to the high activity of metals, they can be obtained by electrolysis of salts, most often chlorides.

Alkali metal compounds are widely used in various industries. See Table. one.

COMMON ALKALI METAL COMPOUNDS
	Caustic soda (caustic soda)
	Salt
	Chilean saltpeter
Na 2 SO 4 ∙10H 2 O	Glauber's salt
Na 2 CO 3 ∙10H 2 O	Crystal soda
	caustic potash
	Potassium chloride (sylvin)
	Indian saltpeter

Their name is due to the fact that the hydroxides of these metals are alkalis, and the oxides used to be called "earths". For example, barium oxide BaO is barium earth. Beryllium and magnesium are most often not classified as alkaline earth metals. We will not consider radium either, since it is radioactive.

Chemical properties of alkaline earth metals.

1. Interaction withnon-metals

Ca + Cl 2 → 2CaCl 2

Ca + H 2 CaH 2

3Сa + 2P Сa 3 P 2-

2. Interaction with oxygen

2Сa + O 2 → 2CaO

3. Interaction with water

Sr + 2H 2 O → Sr(OH) 2 + H 2 , but the interaction is calmer than with alkali metals.

4. Interaction with acids - strong oxidizing agents

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

Obtaining alkaline earth metals

Metallic calcium and strontium are obtained by electrolysis of molten salts, most often chlorides.

CaCl 2 Ca + Cl 2

High purity barium can be obtained by aluminothermic process from barium oxide

3BaO + 2Al 3Ba + Al 2 O 3

COMMON COMPOUNDS OF ALKALINE EARTH METALS

The most famous compounds of alkaline earth metals are: CaO - quicklime. Ca(OH)2 - slaked lime, or lime water. When passing carbon dioxide turbidity occurs through lime water, since insoluble calcium carbonate CaCO 3 is formed. But we must remember that with further transmission of carbon dioxide, an already soluble bicarbonate is formed and the precipitate disappears.

Rice. 2

СaO + H 2 O → Ca (OH) 2

Ca(OH) 2 + CO 2 → CaCO 3 ↓+ H 2 O

CaCO 3 ↓+ H 2 O + CO 2 → Ca(HCO 3) 2

Gypsum - this is CaSO 4 ∙ 2H 2 O, alabaster - CaSO 4 ∙ 0.5H 2 O. Gypsum and alabaster are used in construction, medicine and for the manufacture of decorative products. Rice. 2.

Calcium carbonate CaCO 3 forms many different minerals. Rice. 3.

Rice. 3

calcium phosphate Ca 3 (PO 4) 2 - phosphorite, phosphorus flour is used as a mineral fertilizer.

pure anhydrous calcium chloride CaCl 2 is a hygroscopic substance, therefore it is widely used in laboratories as a desiccant.

calcium carbide- CaC 2 . It can be obtained like this:

CaO + 2C → CaC 2 + CO. One of its uses is the production of acetylene.

CaC 2 + 2H 2 O → Ca (OH) 2 + C 2 H 2

Barium sulfate BaSO 4 - barite. Rice. 4. Used as a reference for white in some studies.

Rice. four

Hardness of water

Natural water contains calcium and magnesium salts. If they are contained in noticeable concentrations, then soap does not lather in such water due to the formation of insoluble stearates. When it boils, scale forms.

Temporary stiffness due to the presence of calcium and magnesium bicarbonates Ca(HCO 3) 2 and Mg(HCO 3) 2. This hardness of water can be eliminated by boiling.

Ca (HCO 3) 2 CaCO 3 ↓ + CO 2 + H 2 O

Permanent water hardness due to the presence of cations Ca 2+ ., Mg 2+ and anions H 2 PO 4 -, Cl - , NO 3 - and others. Constant water hardness is eliminated only due to ion exchange reactions, as a result of which magnesium and calcium ions will be transferred to the precipitate.

Homework

1. No. 3, 4, 5-a (p. 173) Gabrielyan O.S. Chemistry. Grade 11. A basic level of. 2nd ed., ster. - M.: Bustard, 2007. - 220 p.

2. What is the reaction of the environment water solution potassium sulfide? Support your answer with the hydrolysis reaction equation.

3. Determine the mass fraction of sodium in sea ​​water, which contains 1.5% sodium chloride.

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.

Topic quiz

Report Evaluation

Average rating: 4.6. Total ratings received: 294.