Features of the structure of alkaline earth metals. Chemical properties of alkaline earth metals and their compounds

The chemical properties of alkali and alkaline earth metals are similar. Alkali metals have one electron in their outer energy level, while alkaline earth metals have two. During reactions, metals easily part with valence electrons, showing the properties of a strong reducing agent.

alkaline

Group I of the periodic table includes alkali metals:

• lithium;
• sodium;
• potassium;
• rubidium;
• cesium;
• francium.

Rice. 1. Alkali metals.

They are soft (can be cut with a knife), low melting and boiling points. These are the most active metals.

The chemical properties of alkali metals are presented in the table.

May react with organic acids and alcohols.

alkaline earth

Alkaline earth metals are in group II of the periodic table:

• beryllium;
• magnesium;
• calcium;
• strontium;
• barium;
• radium.

Rice. 2. Alkaline earth metals.

Unlike alkali metals, they are harder. Only strontium can be cut with a knife. The most dense metal is radium (5.5 g / cm 3).

Beryllium interacts with oxygen only when heated to 900°C. It does not react with hydrogen and water under any conditions. Magnesium oxidizes at 650°C and reacts with hydrogen under high pressure.

The table reflects the main chemical properties of alkaline earth metals.

Ions of alkali and alkaline earth metals in salts are easy to detect by changing the color of the flame. Sodium salts burn with a yellow flame, potassium - violet, rubidium - red, calcium - brick red, barium - yellow-green. Salts of these metals are used to create fireworks.

Rice. 3. Qualitative reaction.

What have we learned?

Alkali and alkaline earth metals are active elements of the periodic table that react with simple and complex substances. Alkali metals are softer, react violently with water and halogens, easily oxidize in air, forming oxides, peroxides, superoxides, interact with acids and ammonia. When heated, they react with non-metals. Alkaline earth metals react with non-metals, acids, water. Beryllium does not react with hydrogen and water, but reacts with alkalis and oxygen at high temperatures.

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The lesson will cover the topic “Metals and their properties. alkali metals. alkaline earth metals. Aluminum". You will learn the general properties and patterns of alkali and alkaline earth elements, study separately the chemical properties of alkali and alkaline earth metals and their compounds. With the help of chemical equations, such a concept as water hardness will be considered. Learn about 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 of the Periodic system D.I. 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 an ionic bond type. 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. 1.

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

Rice. 1

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 of alkali metals

1. Interaction with non-metals

Due to their 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. 1.

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 carbon dioxide is passed through lime water, turbidity occurs, since insoluble calcium carbonate CaCO 3 is formed. But we must remember that with further passage 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. 4

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 reaction of the medium does an aqueous solution of potassium sulfide have? Support your answer with the hydrolysis reaction equation.

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

Alkaline earth metals include metals of group IIA of the Periodic Table of D.I. Mendeleev - calcium (Ca), strontium (Sr), barium (Ba) and radium (Ra). In addition to them, the main subgroup of group II includes beryllium (Be) and magnesium (Mg). The outer energy level of alkaline earth metals has two valence electrons. The electronic configuration of the external energy level of alkaline earth metals is ns 2 . In their compounds, they exhibit a single oxidation state equal to +2. In OVR, they are reducing agents, i.e. donate an electron.

With an increase in the charge of the nucleus of atoms of elements that are part of the group of alkaline earth metals, the ionization energy of atoms decreases, and the radii of atoms and ions increase, the metallic signs of chemical elements increase.

Physical properties of alkaline earth metals

In the free state, Be is a steel-gray metal with a dense hexagonal crystal lattice, rather hard and brittle. In air, Be is covered with an oxide film, which gives it a matte tint and reduces its chemical activity.

Magnesium in the form of a simple substance is a white metal, which, like Be, acquires a matte hue when exposed to air due to the formation of an oxide film. Mg is softer and more ductile than beryllium. The crystal lattice of Mg is hexagonal.

Free Ca, Ba and Sr are silver-white metals. When exposed to air, they are instantly covered with a yellowish film, which is the products of their interaction with the constituent parts of the air. Calcium is a rather hard metal, Ba and Sr are softer.

Ca and Sr have a cubic face-centered crystal lattice, barium has a cubic body-centered crystal lattice.

All alkaline earth metals are characterized by the presence of a metallic type of chemical bond, which causes their high thermal and electrical conductivity. The boiling and melting points of alkaline earth metals are higher than those of alkali metals.

Obtaining alkaline earth metals

Getting Be is carried out by the reduction reaction of its fluoride. The reaction proceeds when heated:

BeF 2 + Mg = Be + MgF 2

Magnesium, calcium and strontium are obtained by electrolysis of molten salts, most often chlorides:

CaCl 2 \u003d Ca + Cl 2

Moreover, when Mg is obtained by electrolysis of a dichloride melt, NaCl is added to the reaction mixture to lower the melting temperature.

To obtain Mg in industry, metal- and carbon-thermal methods are used:

2(CaO×MgO) (dolomite) + Si = Ca 2 SiO 4 + Mg

The main way to obtain Ba is oxide reduction:

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

Chemical properties of alkaline earth metals

Since in n.a. the surface of Be and Mg is covered with an oxide film - these metals are inert with respect to water. Ca, Sr and Ba dissolve in water to form hydroxides exhibiting strong basic properties:

Ba + H 2 O \u003d Ba (OH) 2 + H 2

Alkaline earth metals are able to react with oxygen, and all of them, with the exception of barium, form oxides as a result of this interaction, barium - peroxide:

2Ca + O 2 \u003d 2CaO

Ba + O 2 \u003d BaO 2

Oxides of alkaline earth metals, with the exception of beryllium, exhibit basic properties, Be - amphoteric properties.

When heated, alkaline earth metals are capable of interacting with non-metals (halogens, sulfur, nitrogen, etc.):

Mg + Br 2 \u003d 2MgBr

3Sr + N 2 \u003d Sr 3 N 2

2Mg + 2C \u003d Mg 2 C 2

2Ba + 2P = Ba 3 P 2

Ba + H 2 = BaH 2

Alkaline earth metals react with acids - dissolve in them:

Ca + 2HCl \u003d CaCl 2 + H 2

Mg + H 2 SO 4 \u003d MgSO 4 + H 2

Beryllium reacts with aqueous solutions of alkalis - it dissolves in them:

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

Qualitative reactions

A qualitative reaction to alkaline earth metals is the coloring of the flame by their cations: Ca 2+ colors the flame dark orange, Sr 2+ dark red, Ba 2+ light green.

A qualitative reaction to the barium cation Ba 2+ are SO 4 2- anions, resulting in the formation of a white precipitate of barium sulfate (BaSO 4), insoluble in inorganic acids.

Ba 2+ + SO 4 2- \u003d BaSO 4 ↓

Examples of problem solving

EXAMPLE 1

The fresh surface of E quickly darkens due to the formation of an oxide film. This film is relatively dense - over time, the entire metal slowly oxidizes. The film consists of EO, as well as EO 2 and E 3 N 2 . The normal electrode potentials of the reactions E-2e = E 2+ are = -2.84V (Ca), = -2.89 (Sr). E are very active elements: they dissolve in water and acids, displace most metals from their oxides, halides, sulfides. Primarily (200-300 o C) calcium interacts with water vapor according to the scheme:

2Ca + H 2 O \u003d CaO + CaH 2.

The secondary reactions are:

CaH 2 + 2H 2 O \u003d Ca (OH) 2 + 2H 2 and CaO + H 2 O \u003d Ca (OH) 2.

In strong sulfuric acid, E almost do not dissolve due to the formation of a film of poorly soluble ESO 4 . With dilute mineral acids, E react violently with the release of hydrogen. When heated above 800 ° C, calcium reacts with methane according to the scheme:

3Ca + CH 4 \u003d CaH 2 + CaC 2.

When heated, they react with hydrogen, with sulfur and with gaseous ammonia. In terms of chemical properties, radium is closest to Ba, but it is more active. At room temperature, it noticeably combines with oxygen and nitrogen in the air. In general, its chemical properties are slightly more pronounced than those of its counterparts. All radium compounds slowly decompose under the action of their own radiation, while acquiring a yellowish or brown color. Radium compounds have the property of autoluminescence. As a result of radioactive decay, 1 g of Ra releases 553.7 J of heat every hour. Therefore, the temperature of radium and its compounds is always higher than the ambient temperature by 1.5 degrees. It is also known that 1 g of radium per day emits 1 mm 3 of radon (226 Ra = 222 Rn + 4 He), on which its use as a source of radon for radon baths is based.

hydrides E - white, crystalline salt-like substances. They are obtained directly from the elements by heating. The start temperatures of the reaction E + H 2 = EN 2 are 250 o C (Ca), 200 o C (Sr), 150 o C (Ba). Thermal dissociation of EN 2 begins at 600 o C. CaH 2 does not decompose in a hydrogen atmosphere at the melting point (816 o C). In the absence of moisture, alkaline earth metal hydrides are stable in air at ordinary temperatures. They do not react with halogens. However, when heated, the chemical activity of EN 2 increases. They are able to reduce oxides to metals (W, Nb, Ti, Ce, Zr, Ta), for example

2CaH 2 + TiO 2 \u003d 2CaO + 2H 2 + Ti.

The reaction of CaH 2 with Al 2 O 3 takes place at 750 o C:

3CaH 2 + Al 2 O 3 \u003d 3CaO + 3H 2 + 2Al,

CaH 2 + 2Al \u003d CaAl 2 + H 2.

CaH2 reacts with nitrogen at 600°C according to the scheme:

3CaH 2 + N 2 \u003d Ca 3 N 2 + 3H 2.

When EN 2 is ignited, they slowly burn out:

EN 2 + O 2 \u003d H 2 O + CaO.

Explosive when mixed with solid oxidizers. Under the action of water on EN 2, hydroxide and hydrogen are released. This reaction is highly exothermic: EN 2 wetted with water in air ignites spontaneously. EN 2 reacts with acids, for example, according to the scheme:

2HCl + CaH 2 \u003d CaCl 2 + 2H 2.

EN 2 is used to obtain pure hydrogen, as well as to determine traces of water in organic solvents. Nitride E are colorless refractory substances. They are obtained directly from the elements at elevated temperatures. They decompose in water according to the scheme:

E 3 N 2 + 6H 2 O \u003d 3E (OH) 2 + 2NH 3.

E 3 N 2 react when heated with CO according to the scheme:

E 3 N 2 + 3CO \u003d 3EO + N 2 + 3C.

The processes that occur when E 3 N 2 is heated with coal look like this:

E3N2 + 5C = ECN2 + 2ES2; (E = Ca, Sr); Ba3N2 + 6C = Ba(CN)2 + 2BaC2;

Strontium nitride reacts with HCl to give Sr and ammonium chlorides. Phosphides E 3 R 2 are formed directly from the elements or by calcining trisubstituted phosphates with coal:

Ca 3 (RO 4) 2 + 4C \u003d Ca 3 P 2 + 4CO

They are hydrolyzed by water according to the scheme:

E 3 R 2 + 6H 2 O \u003d 2RN 3 + 3E (OH) 2.

With acids, alkaline earth metal phosphides give the corresponding salt and phosphine. This is the basis for their use for the production of phosphine in the laboratory.

Complex ammonia composition E (NH 3) 6 - solids with a metallic luster and high electrical conductivity. They are obtained by the action of liquid ammonia on E. They ignite spontaneously in air. Without air access, they decompose into the corresponding amides: E (NH 3) 6 \u003d E (NH 2) 2 + 4NH 3 + H 2. When heated, they vigorously decompose according to the same pattern.

Carbides alkaline earth metals, which are obtained by calcining E with coal, are decomposed by water with the release of acetylene:

ES 2 + 2H 2 O \u003d E (OH) 2 + C 2 H 2.

The reaction with BaC 2 is so violent that it ignites on contact with water. The heats of formation of ES 2 from the elements for Ca and Ba are 14 and 12 kcalmol. When heated with nitrogen, ES 2 give CaCN 2 , Ba(CN) 2 , SrCN 2 . known silicides (ESi and ESi 2). They can be obtained by heating directly from the elements. They hydrolyze with water and react with acids to give H 2 Si 2 O 5 , SiH 4 , the corresponding E compound, and hydrogen. known borides EV 6 obtained from the elements when heated.

Oxides calcium and its analogues are white refractory (T bp CaO = 2850 o C) substances that actively absorb water. This is the basis for the use of BaO to obtain absolute alcohol. They react violently with water, releasing a lot of heat (except for SrO, the dissolution of which is endothermic). EOs dissolve in acids and ammonium chloride:

EO + 2NH 4 Cl \u003d SrCl 2 + 2NH 3 + H 2 O.

EO is obtained by calcining carbonates, nitrates, peroxides or hydroxides of the corresponding metals. The effective charges of barium and oxygen in BaO are 0.86. SrO at 700 o C reacts with potassium cyanide:

KCN + SrO = Sr + KCNO.

Strontium oxide dissolves in methanol to form Sr(OCH 3) 2 . During magnesium-thermal reduction of BaO, an intermediate oxide Ba 2 O can be obtained, which is unstable and disproportionate.

Hydroxides alkaline earth metals - white substances soluble in water. They are strong bases. In the Ca-Sr-Ba series, the basic nature and solubility of hydroxides increase. rPR(Ca(OH) 2) = 5.26, rPR(Sr(OH) 2) = 3.5, rPR(Ba(OH) 2) = 2.3. Ba(OH) 2 is usually isolated from hydroxide solutions. 8H 2 O, Sr (OH) 2. 8H 2 O, Ca (OH) 2. H 2 O. EOs add water to form hydroxides. The use of CaO in construction is based on this. A close mixture of Ca(OH) 2 and NaOH in a 2:1 weight ratio is called soda lime and is widely used as a CO 2 scavenger. Ca (OH) 2, when standing in air, absorbs CO 2 according to the scheme:

Ca(OH)2 + CO2 = CaCO3 + H2O.

About 400 o C Ca (OH) 2 reacts with carbon monoxide:

CO + Ca (OH) 2 \u003d CaCO 3 + H 2.

Barite water reacts with CS 2 at 100 o C:

CS 2 + 2Ba (OH) 2 \u003d BaCO 3 + Ba (HS) 2 + H 2 O.

Aluminum reacts with barite water:

2Al + Ba (OH) 2 + 10H 2 O \u003d Ba 2 + 3H 2. E(OH) 2

used to open carbonic anhydride.

E form peroxides white. They are significantly less stable than oxides and are strong oxidizers. Of practical importance is the most stable BaO 2, which is a white, paramagnetic powder with a density of 4.96 g1cm 3 etc. pl. 450°. BaO 2 is stable at normal temperature (it can be stored for years), it is poorly soluble in water, alcohol and ether, it dissolves in dilute acids with the release of salt and hydrogen peroxide. The thermal decomposition of barium peroxide is accelerated by oxides, Cr 2 O 3 , Fe 2 O 3 and CuO. Barium peroxide reacts when heated with hydrogen, sulfur, carbon, ammonia, ammonium salts, potassium ferricyanide, etc. Barium peroxide reacts with concentrated hydrochloric acid, releasing chlorine:

BaO 2 + 4HCl = BaCl 2 + Cl 2 + 2H 2 O.

It oxidizes water to hydrogen peroxide:

H 2 O + BaO 2 \u003d Ba (OH) 2 + H 2 O 2.

This reaction is reversible and in the presence of even carbonic acid the equilibrium is shifted to the right. ВаО 2 is used as a starting product for the production of Н 2 О 2 and also as an oxidizing agent in pyrotechnic compositions. However, BaO 2 can also act as a reducing agent:

HgCl 2 + BaO 2 \u003d Hg + BaCl 2 + O 2.

BaO 2 is obtained by heating BaO in air flow to 500 ° C according to the scheme:

2ВаО + О 2 = 2ВаО 2.

As the temperature rises, the reverse process takes place. Therefore, when Ba burns, only oxide is released. SrO 2 and CaO 2 are less stable. A common method for producing EO 2 is the interaction of E(OH) 2 with H 2 O 2 , whereby EO 2 is released. 8H 2 O. Thermal decomposition of EO 2 begins at 380 o C (Ca), 480 o C (Sr), 790 o C (Ba). When EO 2 is heated with concentrated hydrogen peroxide, yellow unstable substances, EO 4 superoxides, can be obtained.

E salts are usually colorless. Chlorides, bromides, iodides and nitrates are highly soluble in water. Fluorides, sulfates, carbonates and phosphates are poorly soluble. Ion Ba 2+ - toxic. Halides E are divided into two groups: fluorides and all the rest. Fluorides are almost insoluble in water and acids and do not form crystalline hydrates. On the contrary, chlorides, bromides, and iodides are highly soluble in water and are isolated from solutions in the form of crystalline hydrates. Some properties of EG 2 are presented below:

When obtained by exchange decomposition in solution, fluorides are released in the form of voluminous mucous precipitates, which quite easily form colloidal solutions. EG 2 can be obtained by acting with the corresponding halogens on the corresponding E. EG 2 melts are capable of dissolving up to 30% E. When studying the electrical conductivity of chloride melts of elements of the second group of the main subgroup, it was found that their molecular-ionic composition is very different. The degrees of dissociation according to the scheme ESl 2 = E 2+ + 2Cl- are equal: BeCl 2 - 0.009%, MgCl 2 - 14.6%, CaCl 2 - 43.3%, SrCl 2 - 60.6%, BaCl 2 - 80, 2%. Halides (except fluorides) E contain water of crystallization: CaCl 2 . 6H 2 O, SrCl 2. 6H 2 O and BaCl 2. 2H 2 O. X-ray diffraction analysis established the structure of E[(OH 2) 6 ]G 2 for Ca and Sr crystalline hydrates. With slow heating of EG 2 crystalline hydrates, anhydrous salts can be obtained. CaCl 2 readily forms supersaturated solutions. Natural CaF 2 (fluorite) is used in the ceramic industry, and is also used to produce HF and is a fluorine mineral. Anhydrous CaCl 2 is used as a desiccant due to its hygroscopicity. Calcium chloride hydrate is used for the preparation of refrigeration mixtures. BaCl 2 - used in cx and for opening

SO 4 2- (Ba 2+ + SO 4 2- \u003d BaSO 4).

Fusion of EG2 and EN2 hydrohalides can be obtained:

EG 2 + EN 2 = 2ENG.

These substances melt without decomposition but are hydrolyzed by water:

2ENG + 2H 2 O \u003d EG 2 + 2H 2 + E (OH) 2.

Solubility in water chlorates , bromates And iodates in water it decreases in the series Ca - Sr - Ba and Cl - Br - I. Ba (ClO 3) 2 - is used in pyrotechnics. Perchlorates E are highly soluble not only in water but also in organic solvents. The most important of the E(ClO 4) 2 is Ba(ClO 4) 2 . 3H 2 O. Anhydrous barium perchlorate is a good drying agent. Its thermal decomposition begins only at 400 o C. Hypochlorite calcium Ca (ClO) 2. nH 2 O (n=2.3.4) is obtained by the action of chlorine on milk of lime. It is an oxidizing agent and is highly soluble in water. bleach can be obtained by acting with chlorine on solid slaked lime. It decomposes with water and smells like chlorine in the presence of moisture. Reacts with CO 2 of air:

CO 2 + 2CaOCl 2 \u003d CaCO 3 + CaCl 2 + Cl 2 O.

Chlorine lime is used as an oxidizing agent, bleach and as a disinfectant.

For alkaline earth metals, azides E(N 3) 2 and thiocyanates E(CNS) 2 . 3H 2 O. Azides are much less explosive than lead azide. The thiocyanates easily lose water when heated. They are highly soluble in water and organic solvents. Ba(N 3) 2 and Ba(CNS) 2 can be used to obtain azides and thiocyanates of other metals from sulfates by an exchange reaction.

Nitrates calcium and strontium usually exist in the form of Ca(NO 3) 2 crystalline hydrates. 4H 2 O and Sr(NO 3) 2 . 4H 2 O. For barium nitrate, the formation of a crystalline hydrate is not characteristic. When heated, Ca (NO 3) 2. 4H 2 O and Sr(NO 3) 2 . 4H 2 O easily lose water. In an inert atmosphere, nitrates E are thermally stable up to 455 o C (Ca), 480 o C (Sr), 495 o C (Ba). The hydrated melt of calcium nitrate has an acidic environment at 75 ° C. A feature of barium nitrate is the low rate of dissolution of its crystals in water. Only barium nitrate exhibits a tendency to complex formation, for which an unstable complex K 2 is known. Calcium nitrate is soluble in alcohols, methyl acetate, acetone. Strontium and barium nitrates are almost insoluble there. The melting points of nitrates E are estimated at 600 o C, however, at the same temperature, decomposition begins:

E (NO 3) 2 \u003d E (NO 2) 2 + O 2.

Further decomposition occurs at a higher temperature:

E (NO 2) 2 \u003d EO + NO 2 + NO.

E nitrates have long been used in pyrotechnics. Highly volatile salts E color the flame in the appropriate colors: Ca - orange-yellow, Sr - red-carmine, Ba - yellow-green. Let's understand the essence of this using the example of Sr: Sr 2+ has two HAOs: 5s and 5p or 5s and 4d. We will inform the energy of this system - we will heat it. Electrons from orbitals closer to the nucleus will move to these HAOs. But such a system is not stable and will release energy in the form of a quantum of light. Just Sr 2+ emits quanta with a frequency corresponding to the lengths of red waves. When obtaining pyrotechnic compositions, it is convenient to use saltpeter, because. it not only colors the flame, but is also an oxidizing agent, releasing oxygen when heated. Pyrotechnic compositions consist of a solid oxidant, a solid reducing agent and some organic substances that bleach the flame of the reducing agent and act as a binding agent. Calcium nitrate is used as a fertilizer.

All phosphates And hydrophosphates E are poorly soluble in water. They can be obtained by dissolving an appropriate amount of CaO or CaCO 3 in phosphoric acid. They are also precipitated during exchange reactions such as:

(3-x) Ca 2+ + 2H x PO 4 - (3-x) \u003d Ca (3-x) (H x PO 4) 2.

Of practical importance (as a fertilizer) is monosubstituted calcium orthophosphate, which, along with Ca (SO 4), is part of superphosphate. It is received according to the scheme:

Ca 3 (PO 4) 2 + 2H 2 SO 4 \u003d Ca (H 2 PO 4) 2 + 2CaSO 4

Oxalates also slightly soluble in water. Of practical importance is calcium oxalate, which dehydrates at 200 ° C, and decomposes at 430 ° C according to the scheme:

CaC 2 O 4 \u003d CaCO 3 + CO.

Acetates E are isolated in the form of crystalline hydrates and are highly soluble in water.

WITH sulfates E - white, poorly soluble substances in water. Solubility CaSO 4 . 2H 2 O per 1000 g of water at normal temperature is 8. 10 -3 mol, SrSO 4 - 5. 10 -4 mol, BaSO 4 - 1. 10 -5 mol, RaSO 4 - 6. 10 -6 mol. In the Ca - Ra series, the solubility of sulfates decreases rapidly. Ba 2+ is a sulfate ion reagent. Calcium sulfate contains water of crystallization. Above 66 ° C, anhydrous calcium sulfate is released from the solution, below - CaSO 4 gypsum. 2H 2 O. Heating of gypsum above 170 ° C is accompanied by the release of hydrated water. When gypsum is mixed with water, this mass quickly hardens due to the formation of crystalline hydrate. This property of gypsum is used in construction. The Egyptians used this knowledge as early as 2000 years ago. The solubility of ESO 4 in strong sulfuric acid is much higher than in water (BaSO 4 up to 10%), which indicates complex formation. The corresponding complexes are ESO 4 . H 2 SO 4 can be obtained in the free state. Double salts with alkali metal and ammonium sulfates are known only for Ca and Sr. (NH 4) 2 is soluble in water and is used in analytical chemistry to separate Ca from Sr, because (NH 4) 2 is slightly soluble. Gypsum is used for the combined production of sulfuric acid and cement, because. when heated with a reducing agent (charcoal), gypsum decomposes:

CaSO 4 + C \u003d CaO + SO 2 + CO.

At a higher temperature (900 o C), sulfur is reduced even more according to the scheme:

CaSO 4 + 3C \u003d CaS + CO 2 + 2CO.

A similar decomposition of Sr and Ba sulfates begins at higher temperatures. BaSO 4 is non-toxic and is used in medicine and in the production of mineral paints.

Sulfides E are white solids that crystallize as NaCl. The heats of their formation and the energies of the crystal lattices are (kcalmol): 110 and 722 (Ca), 108 and 687 (Sr), 106 and 656 (Ba). They can be obtained by synthesis from elements when heated or by calcining sulfates with coal:

ES04 + 3C = ES + CO2 + 2CO.

Less soluble CaS (0.2 hl). ES enters into the following reactions when heated:

ES + H 2 O \u003d EO + H 2 S; ES + G 2 \u003d S + EG 2; ES + 2O 2 \u003d ESO 4; ES + xS = ES x+1 (x=2.3).

Sulfides of alkaline earth metals in a neutral solution are completely hydrolyzed according to the scheme:

2ES + 2H 2 O \u003d E (HS) 2 + E (OH) 2.

Acid sulfides can also be obtained in a free state by evaporating a solution of sulfides. They react with sulfur:

E (NS) 2 + xS \u003d ES x + 1 + H 2 S (x \u003d 2.3.4).

Of the crystalline hydrates, BaS are known. 6H 2 O and Ca(HS) 2 . 6H 2 O, Ba (HS) 2. 4H 2 O. Ca(HS) 2 is used to remove hair. ES are subject to the phenomenon of phosphorescence. known polysulfides E: ES 2, ES 3, ES 4, ES 5. They are obtained by boiling a suspension of ES in water with sulfur. In air, ES are oxidized: 2ES + 3O 2 \u003d 2ESO 3. By passing air through the CaS suspension, one can obtain thiosulfate Sa according to the scheme:

2CaS + 2O 2 + H 2 O \u003d Ca (OH) 2 + CaS 2 O 3

It is highly soluble in water. In the Ca - Sr - Ba series, the solubility of thiosulfates decreases. Tellurides E are slightly soluble in water and are also subject to hydrolysis, but to a lesser extent than sulfides.

Solubility chromates E in the series Ca - Ba falls as sharply as in the case of sulfates. These yellow substances are obtained by the interaction of soluble salts of E with chromates (or dichromates) of alkali metals:

E 2+ + CrO 4 2- = ECrO4.

Calcium chromate is released in the form of a crystalline hydrate - CaCrO 4 . 2H 2 O (rPR CaCrO 4 = 3.15). Even before the melting point, it loses water. SrCrO 4 and ВаCrO 4 do not form crystalline hydrates. pPR SrCrO 4 = 4.44, pPR BaCrO 4 = 9.93.

Carbonates E white, poorly soluble substances in water. When heated, ESO 3 pass into EO, splitting off CO 2 . In the Ca-Ba series, the thermal stability of carbonates increases. The most practically important of them is calcium carbonate (limestone). It is directly used in construction, and also serves as a raw material for the production of lime and cement. The annual world production of lime from limestone amounts to tens of millions of tons. Thermal dissociation of CaCO 3 is endothermic:

CaCO 3 \u003d CaO + CO 2

and requires an expenditure of 43 kcal per mole of limestone. Calcination of CaCO 3 is carried out in shaft furnaces. A by-product of roasting is valuable carbon dioxide. CaO is an important building material. When mixed with water, crystallization occurs due to the formation of hydroxide, and then carbonate according to the schemes:

CaO + H 2 O \u003d Ca (OH) 2 and Ca (OH) 2 + CO 2 \u003d CaCO 3 + H 2 O.

A colossally important practical role is played by cement, a greenish-gray powder consisting of a mixture of various silicates and calcium aluminates. When mixed with water, it hardens due to hydration. In its production, a mixture of CaCO 3 with clay is fired before sintering (1400-1500 about C). The mixture is then ground. The composition of cement can be expressed as a percentage of the components Cao, SiO 2, Al 2 O 3, Fe 2 O 3, with Cao being a base, and everything else being acid anhydrides. The composition of silicate (Portland) cement is composed mainly of Ca 3 SiO 5 , Ca 2 SiO 4 , Ca 3 (AlO 3) 2 and Ca (FeO 2) 2 . His grasp goes according to the schemes:

Ca 3 SiO 5 + 3H 2 O \u003d Ca 2 SiO 4. 2H 2 O + Ca (OH) 2

Ca 2 SiO 4 + 2H 2 O \u003d Ca 2 SiO 4. 2H 2 O

Ca 3 (AlO 3) 2 + 6H 2 O \u003d Ca 3 (AlO 3) 2. 6H 2 O

Ca (FeO 2) 2 + nH 2 O \u003d Ca (FeO 2) 2. nH2O.

Natural chalk is introduced into the composition of various putties. Fine-crystalline, precipitated from a solution of CaCO 3 is part of tooth powders. BaO is obtained from BaCO 3 by calcination with coal according to the scheme:

VaCO 3 + C \u003d BaO + 2CO.

If the process is carried out at a higher temperature in a stream of nitrogen, cyanide barium:

VaCO 3 + 4C + N 2 \u003d 3CO + Ba (CN) 2.

Ba(CN) 2 is highly soluble in water. Ba(CN) 2 can be used to produce other metal cyanides by exchange decomposition with sulfates. Bicarbonates E are soluble in water and can only be obtained in solution, for example, by passing carbon dioxide into a suspension of CaCO 3 in water:

CO 2 + CaCO 3 + H 2 O \u003d Ca (HCO 3) 2.

This reaction is reversible and shifts to the left when heated. The presence of calcium and magnesium bicarbonates in natural waters causes water hardness.

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 of the periodic table of Mendeleev. Full list of 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 at the external energy level, active metals exhibit powerful reducing agent properties, giving up external 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 of 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.

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. These are soft simple substances of gray or silver color with a low 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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