The third group of the periodic system. Physical properties of elements of the main subgroup of group III Physical properties of elements of group 3 of the main subgroup

The p-elements of group III of the periodic system of D. I. Mendeleev include: boron B, aluminum gallium indium and thallium Electronic configuration of atoms

Some constants characterizing the properties of atoms of the p-elements of the group under consideration and the corresponding metallic substances are compared below:

The properties of -elements of group III are affected by d-compression located in the periodic system in a small period III, in large periods immediately after the d-elements). So, from the atomic radius decreases somewhat, and the first ionization potential increases. The properties of thallium atoms are also affected by -compression. That is why the radius of the atom is close to the radius of the atom, and the ionization energy is somewhat higher.

Bor. In accordance with the electronic structure of the atom, boron can be monovalent (one unpaired electron on the energy sublevel). However, boron is most characteristic of compounds in which it is trivalent (when an atom is excited, there are three unpaired electrons in the energy and - sublevels).

The free -orbital in the excited boron atom determines the acceptor properties of many of its compounds, in which three covalent bonds are formed according to the exchange covalent mechanism (for example, These compounds are prone to the addition of particles with electron-donor properties, i.e., to the formation of another covalent bond along donor-acceptor mechanism.For example:

Two isotopes of boron are known: The nuclei of atoms of the isotope easily absorb neutrons:

The ability of boron to absorb neutrons determines its use in nuclear power engineering: control rods of nuclear reactors are made from boron-containing materials.

Black boron crystals; they are refractory (mp. 2300 ° C), diamagnetic, have semiconductor properties (band gap. The electrical conductivity of boron, like other metals, is small and increases somewhat with increasing temperature.

At room temperature, boron is chemically inert and interacts directly only with fluorine; when heated, boron is oxidized by chlorine, oxygen, and some other non-metals. For example:

In compounds with non-metals, the oxidation state of boron is that all these compounds are covalent.

Boron trioxide is a crystalline substance (mp 450°C, bp 2250°C) characterized by high enthalpy and Gibbs energy of formation. When interacting with water, it turns into boric acid:

A very weak monobasic acid. Electrolytic dissociation with the elimination of only one ion is explained by the previously described acceptor properties of boron: the free -orbital of the boron atom is provided to the electron donor formed during the dissociation of molecules. The process proceeds according to the scheme

The complex anion has a tetrahedral structure (-hybridization of electron orbitals).

The acceptor properties of boron in compounds with an oxidation state are also manifested in the chemistry of its halides. For example, the reactions

in which the chemical bond between and or is formed by the donor-acceptor mechanism. The property of boron halides to be electron acceptors determines their wide application as catalysts in the reactions of synthesis of organic compounds.

Boron does not interact directly with hydrogen, but forms borides with metals - usually non-stoichiometric compounds

Boron hydrides (boranes) are very poisonous and have a very unpleasant odor. They are obtained indirectly, most often

in the interaction of reactive borides with acids or boron halides with alkali metal hydrides:

The simplest compound of boron with hydrogen does not exist under normal conditions. -Hybridization of electron orbitals in the boron atom leads to the coordination unsaturation of the particle, as a result of which two such particles combine into a diborane molecule:

In diborane, boron is in the -hybridization state, with each boron atom having one of the four hybrid orbitals empty, while the other three are overlapped by -orbitals of hydrogen atoms. Bonds between groups in a molecule are formed in the form of a hydrogen bond due to the shift of the electron density from one hydrogen atom of the group to the empty orbital of another group Other boranes are also known, which can be represented by two rows

The β-metal borides are reactive and are often used to produce a mixture of boranes by treatment with acids. Most borides and heat resistant, very hard, chemically resistant. They are widely used directly in the form of alloys for the manufacture of jet engine parts, gas turbine blades. Some borides are used to make cathodes for electronic devices.

Aluminum. The electronic configuration of an aluminum atom is expressed by the formula There is one unpaired electron on the outer electron layer of the atom:

Therefore, aluminum can exhibit a valence equal to one. However, this valence is not typical for aluminum. In all stable compounds, the oxidation state of aluminum is equal to Valence, equal to three, corresponds to the excited state of the atom

In terms of its prevalence, aluminum ranks fourth among all elements (after O, H and Si) and is the most common metal in nature. The bulk of aluminum is concentrated in aluminosilicates: feldspars, micas, etc.

Aluminum is a silvery-white light and extremely ductile metal with high thermal and electrical conductivity.

Aluminum is reactive; it reacts with chlorine and bromine at room temperature, and with iodine - when heated or in the presence of water as a catalyst. At 800 °C, aluminum interacts with nitrogen, and at 2000 °C, with carbon. Aluminum exhibits a high chemical affinity for oxygen:

In air, aluminum is covered with a very strong thin oxide film, which somewhat weakens the metallic luster of aluminum. Thanks to the oxide film, the surface of aluminum acquires high corrosion resistance. This is primarily manifested in the indifference of aluminum to water and steam. Due to the formation of a protective film, aluminum is resistant to concentrated nitric and sulfuric acids. These acids passivate aluminum in the cold. The tendency to passivate allows you to increase the corrosion resistance of aluminum by treating its surface with strong oxidizing agents (for example,) or using anodic oxidation. In this case, the thickness of the oxide film increases to . At high temperatures, the strength of the protective film decreases sharply. If the oxide film is removed by mechanical action, aluminum becomes extremely reactive. It reacts vigorously with water and aqueous solutions of acids and alkalis, displacing hydrogen and forming cations or anions. The interaction of aluminum with acid solutions proceeds according to the reaction equation

and with alkali solutions

Aluminum cations and anions easily pass into each other when the pH of the solution changes:

Mixed compounds can also be formed in solution,

for example

The latter is easily (especially when heated) dehydrated and converted into hydroxide

The widest application of aluminum in engineering is based on its valuable physical and chemical properties and its high prevalence in the earth's crust. Due to its high electrical conductivity and low density, it

used for making electrical wires. The high ductility of aluminum makes it possible to produce the thinnest foil from it, which is used in capacitors and replaces lead in cable sheaths with aluminum. Due to non-magnetization, aluminum alloys are used in radio engineering.

The bulk of aluminum is used to produce light alloys - duralumin, the rest is silumin, the rest, etc. Aluminum is also used as an alloying additive to alloys to give them heat resistance. Aluminum and its alloys occupy one of the main places as structural materials in aircraft construction, rocketry, mechanical engineering, etc. The corrosion resistance of aluminum (especially anodized) significantly exceeds the corrosion resistance of steel. Therefore, its alloys are used as structural materials and in shipbuilding. With d-elements, aluminum forms chemical compounds - intermetallides (aluminides): etc., which are used as heat-resistant materials. Aluminum is used in aluminothermy to obtain a number of metals and for thermite welding. Aluminothermy is based on the high affinity of aluminum for oxygen. For example, in a reaction proceeding according to the equation

about 3500 kJ of heat is released and the temperature develops up to

Aluminum oxide is known in the form of several modifications. The most stable is This modification is found in the earth's crust in the form of the mineral corundum, from which grinding discs and emery powders are prepared. The use of corundum as an abrasive material is based on its high hardness, second only to the hardness of diamond, carborundum and borazone. Artificial rubies are obtained by fusion. They are used to make supporting stones in precise mechanisms. Recently, artificial rubies have been used in quantum generators (lasers). Products from are used as refractories and dielectrics.

Aluminum hydroxide is a polymer compound. It has a layered crystal lattice. Each layer consists of octahedrons (Fig. IX. 10); there is a hydrogen bond between the layers. The aluminum hydroxide obtained by the exchange reaction is a gelatinous white precipitate, readily soluble in acids and alkalis. When standing, the precipitate "ages" and loses its chemical activity. When calcined, the hydroxide loses water and turns into oxide. One of the forms of dehydrated hydroxide, aluminum gel, is used in technology as an adsorbent.

Compounds of great interest

Rice. IX. 10. Structure of the layer formed by the octahedral structural units of the compound

aluminum - zeolites related to aluminosilicates. Their composition can be expressed by the general formula where or (rarely ).

Aluminum is in the main subgroup of Group III of the Periodic Table. At the external energy level of the aluminum atom, there are free p-orbitals, which allows it to go into an excited state. In an excited state, the aluminum atom forms three covalent bonds or completely gives up three valence electrons, showing an oxidation state of +3.

aluminum is most common metal on earth : its mass fraction in the earth's crust is 8.8%. The bulk of natural aluminum is part of aluminosilicates - substances, the main components of which are oxides of silicon and aluminum.

Aluminum is a light silver-white metal, melts at 600°C, is very ductile, easily drawn into wire and rolled into sheets and foil. In terms of electrical conductivity, aluminum is second only to silver and copper.

Interaction with simple substances:

1) with halogens:

2Al + 3Cl 2 = 2AlCl 3

2) with oxygen:

4Al + 3O 2 \u003d 2Al 2 O 3

3) with sulfur:

2Al + 3S = Al 2 S 3

4) with nitrogen:

Aluminum does not directly react with hydrogen, but its AlH3 hydride was obtained indirectly.

Interaction with complex substances:

1) with acids:

2Al + 6HCl = 2AlCl 3 + 3H 2

2) with alkalis:

2Al + 2NaOH + 6H 2 O = 2Na + 3H 2

If NaOH is in the solid state:

2Al + 2NaOH + 6H2O = 2NaAlO 2 + 3H 2

3) with water:

2Al + 6H2O = 2Al(OH) 3 + 3H2

Properties of aluminum oxide and hydroxide: aluminum oxide, or alumina, Al 2 O 3 is a white powder. Aluminum oxide can be obtained by burning metal or by calcining aluminum hydroxide:

2Al(OH)3 = Al 2 O 3 + 3H 2 O

Aluminum oxide is practically insoluble in water. The hydroxide Al (OH) 3 corresponding to this oxide is obtained by the action of ammonium hydroxide or alkali solutions, taken in deficiency, on solutions of aluminum salts:

AlCl 3 + 3NH 3 H2O = Al(OH)3 + 3NH4Cl

The oxide and hydroxide of this metal are amphoteric those. exhibit both basic and acidic properties.

Basic properties:

Al 2 O 3 + 6HCl \u003d 2AlCl 3 + 3H 2 O

2Al(OH) 3 + 3H 2 SO 4 = Al 2 (SO 4) 3 + 6H 2 O

Acid properties:

Al 2 O 3 + 6KOH + 3H 2 O \u003d 2K 3

2Al(OH) 3 + 6KOH = K 3

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

Aluminum receive electrolytic method. It cannot be isolated from aqueous solutions of salts, because is a very active metal. Therefore, the main industrial method for obtaining metallic aluminum is the electrolysis of a melt containing aluminum oxide and cryolite.

Metallic aluminum is widely used in industry, in terms of production it ranks second after iron. The bulk of aluminum goes to the manufacture of alloys:

Duralumin - an aluminum alloy containing copper and a small amount of magnesium, manganese and other components. Duralumins are light, strong and corrosion-resistant alloys. Used in aircraft and mechanical engineering.

magnalin - an alloy of aluminum and magnesium. Used in aircraft and mechanical engineering, in construction. Resistant to corrosion in sea water, so it is used in shipbuilding. Silumin is an aluminum alloy containing silicon. Good for casting. This alloy is used in automotive, aircraft and mechanical engineering, the production of precision instruments. Aluminum is a ductile metal, so it is used to make thin foil used in the manufacture of radio engineering products and for packaging goods. Wires are made of aluminum, silver paints.

Chapter XI. THE THIRD GROUP OF THE PERIODIC

ELEMENT SYSTEMS

83. General characteristics of the elements of group III

Group III includes boron, aluminum, gallium, indium, thallium (the main subgroup), as well as scandium, yttrium, lanthanum and lanthanides, actinium and actinides (a side subgroup).

At the outer electronic level of the elements of the main subgroup, there are three electrons each (s 2 p 1). They easily donate these electrons or form three unpaired electrons due to the transition of one electron to the p level. For boron and aluminum, compounds are typical only with an oxidation state of +3. The elements of the gallium subgroup (gallium, indium, thallium) also have three electrons in the outer electronic level, forming the s 2 p 1 configuration, but they are located after the 18-electron layer. Therefore, unlike aluminum, gallium has clearly non-metallic properties. These properties in the series Ga, In, Tl weaken, and the metallic properties are enhanced.

The elements of the scandium subgroup also have three electrons in the outer electronic level. However, these elements are transitional d-elements, the electronic configuration of their valence layer is d 1 s 2 . These electrons donate all three elements rather easily. The elements of the lanthanide subgroup have a distinctive configuration of the outer electronic level: the 4f level builds up in them and the d level disappears. Starting with cerium, all elements, except for gadolinium and lutetium, have an electronic configuration of the outer electronic level 4f n 6s 2 (gadolinium and lutetium have 5d 1 electrons). The number n varies from 2 to 14. Therefore, s- and f-electrons take part in the formation of valence bonds. Most often, the oxidation state of lanthanides is +3, less often +4.

The electronic structure of the valence layer of actinides in many respects resembles the electronic structure of the valence layer of lanthanides. All lanthanides and actinides are typical metals.

All elements of group III have a very strong affinity for oxygen, and the formation of their oxides is accompanied by the release of a large amount of heat.

Elements of the III group find the most diverse application.

Boron was discovered by J. Gay-Lussac and L. Tenard in 1808. Its content in the earth's crust is 1.2·10 - 3%.

Boron compounds with metals (borides) have high hardness and heat resistance. Therefore, they are used to obtain superhard and heat-resistant special alloys. Boron carbide and boron nitride have high heat resistance. The latter is used as a high-temperature lubricant. The hydrated sodium tetraborate Na 2 B 4 O 7 10H 2 O (borax) has a constant composition, its solutions are used in analytical chemistry to determine the concentration of acid solutions. The reaction of borax with acid proceeds according to the equation

Na 2 B 4 O 7 + 2 HCl + 5 H 2 O \u003d 2 NaCl + 4 H 3 BO 3

Compounds of gallium with group VI elements (sulfur, selenium, tellurium) are semiconductors. High-temperature thermometers are filled with liquid gallium.

Indium was discovered by T. Richter and F. Reich in 1863. Its content in the earth chorus is 2.5 10 - 5%. The addition of indium to copper alloys increases the resistance of the latter to the action of sea water. The addition of this metal to silver increases the brilliance of silver and prevents it from tarnishing in air. Indium coatings protect metals from corrosion. It is part of some alloys used in dentistry, as well as some fusible alloys (an alloy of indium, bismuth, lead, tin and cadmium melts at 47 С). Compounds of indium with various non-metals have semiconductor properties.

Thallium was discovered by W. Crookes in 1861. Its content in the earth's crust is 10-4%. An alloy of thallium (10%) with tin (20%) and lead (70%) has a very high acid resistance, it withstands the action of a mixture of sulfuric, hydrochloric and nitric acids. Thallium increases the sensitivity of photocells to infrared radiation coming from heated objects. Thallium compounds are highly toxic and cause hair loss.

Gallium, indium and thallium are trace elements. Their content in ores, as a rule, does not exceed thousandths of a percent.

Compounds of scandium, yttrium, lanthanum and lanthanides were known as early as the beginning of the 19th century. Pure scandium was isolated by L.F. Nilson in 1879. The content of this element in the earth's crust is 10-3%. Yttrium was discovered by Yu. Gadolin in 1794. Its content in the earth's crust is 2.9·10 - 3%. The content of lanthanum in the earth's crust, discovered by K. G. Mosander in 1839, is 4.9 10 - 3%. These metals are mainly used to obtain special alloys with specific electrical and magnetic properties. In addition, lanthanides are used to prepare various pyrophoric compositions, cerium - to obtain aluminum alloys. The addition of cerium increases the electrical conductivity of aluminum and improves its mechanical properties, and facilitates the rolling of tungsten. Cerium dioxide is used in the grinding of optical glass.

The actinide family includes the heaviest elements, following actinium in the periodic system.

Of the actinides, uranium and thorium find practical application.

and plutonium.

Uranium was discovered by M. G. Klaproth in 1789. Its content in the earth's crust is 2.5 10 - 4%. In nature, uranium occurs in the form of three isotopes: 238 U - 99.285%, 235 U - 0.71%, 234 U - 0.005%. The isotope 235 U is capable of spontaneous decay. Therefore, uranium used in reactors as nuclear fuel is enriched in order to increase the content of the isotope-235 in it. For this isotope, there is the concept of a critical mass, upon reaching which a chain reaction begins and a nuclear explosion occurs. If the mass of 235 U is less than the critical one, the rate of the spontaneous decay reaction can be controlled. This property of 235 U is used in a nuclear reactor. Uranium compounds are also used as dyes in the printing and silicate industries.

Thorium dioxide was discovered by J. Ya. Berzelius in 1828, but metallic thorium was obtained relatively recently. The content of thorium in the earth's crust is 1.3·10 - 3%. Small additions of this metal to tungsten increase the service life of electric coils in incandescent lamps (thorium absorbs gases that contribute to the rapid wear of the tungsten filament). Thorium dioxide is used in medicine, as well as in the manufacture of some catalysts.

Plutonium was discovered by G. Seaborg, E. Macmillan, J. Kennedy and A. Wahl in 1940. Its content in the earth's crust is negligible. Plutonium is obtained from the decay products of fuel from nuclear reactors. It is used for the same purposes as uranium-235.

84. Aluminum

Aluminum was first obtained chemically by the Danish chemist H. K. Oersted in 1825. In 1854, the French chemist A. E. St. Clair Deville isolated it by the electrochemical method.

Being in nature. Aluminum is the most common metal in nature. Its content in the earth's crust is 8.05%. The most important natural aluminum compounds are aluminosilicates, bauxite, and corundum.

Aluminosilicates make up the bulk of the earth's crust. The product of their weathering is clay and feldspars (orthoclase, albite, anorthite). The basis of clay is kaolin Al 2 O 3 2SiO 2 2H 2 O.

Bauxite is a rock from which aluminum is obtained. Consists mainly of aluminum oxide hydrates Al 2 O 3 nH 2 O.

Physical properties. Aluminum is a silvery-white light metal that melts at 660°C. Very ductile, easily drawn into wire and rolled into sheets: it can be made into foil less than 0.01 mm thick. Aluminum has a very high thermal and electrical conductivity. Its alloys with various metals are strong and light.

Chemical properties. Aluminum is a very active metal. In a series of voltages, it is after the alkali and alkaline earth metals. However, it is quite stable in air, since its surface is covered with a very dense oxide film, which protects the metal from contact with air. If the protective oxide film is removed from the aluminum wire, then aluminum will begin to interact vigorously with oxygen and water vapor in the air, turning into a loose mass - aluminum hydroxide:

4 Al + 3 O 2 + 6 H 2 O \u003d 4 Al (OH) 3

This reaction is accompanied by the release of heat.

Purified from the protective oxide film, aluminum interacts with water with the release of hydrogen:

2 Al + 6 H 2 O \u003d 2 Al (OH) 3 + 3 H 2

Aluminum dissolves well in dilute sulfuric and hydrochloric acids:

2 Al + 6 Hcl \u003d 2 AlCl 3 + 3 H 2

2 Al + 3 H 2 SO 4 \u003d Al 2 (SO 4) 3 +3 H 2

Diluted nitric acid passivates aluminum in the cold, but when heated, aluminum dissolves in it with the release of nitrogen monoxide, nitrogen hemioxide, free nitrogen or ammonia, for example:

8 Al + 30 HNO 3 \u003d 8 Al (NO 3) 3 + 3 N 2 O + 15 H 2 O

Concentrated nitric acid passivates aluminum.

Since aluminum oxide and hydroxide have amphoteric

properties, aluminum is easily soluble in aqueous solutions of all alkalis, except for ammonium hydroxide:

2 Al + 6 KOH + 6 H 2 O \u003d 2 K 3 [Al (OH) 6] + 3 H 2

Powdered aluminum readily reacts with halogens, oxygen and all non-metals. To start the reactions, heating is necessary, then they proceed very intensively and are accompanied by the release of a large amount of heat:

2 Al + 3 Br 2 = 2 AlBr 3 (aluminum bromide)

4 Al + 3 O 2 \u003d 2 Al 2 O 3 (aluminum oxide)

2 Al + 3 S = Al 2 S 3 (aluminum sulfide)

2 Al + N 2 = 2 AlN (aluminum nitride)

4 Al + 3 C \u003d Al 4 C 3 (aluminum carbide)

Aluminum sulfide can only exist in solid form. In aqueous solutions, it undergoes complete hydrolysis with the formation of aluminum hydroxide and hydrogen sulfide:

Al 2 S 3 + 6 H 2 O \u003d 2 Al (OH) 3 + 3 H 2 S

Aluminum easily takes away oxygen and halogens from oxides and salts of other metals. The reaction is accompanied by the release of a large amount of heat:

8 Al + 3 Fe 3 O 4 \u003d 9 Fe + 4 Al 2 O 3

The process of reducing metals from their oxides with aluminum is called aluminothermy. Aluminothermy is used in the production of some rare metals that form a strong bond with oxygen (niobium, tantalum, molybdenum, tungsten, etc.), as well as for welding rails. If a mixture of fine aluminum powder and magnetic iron ore Fe 3 O 4 (termite) is ignited with a special fuse, then the reaction proceeds spontaneously with the mixture heated to 3500 С. Iron at this temperature is in a molten state.

Receipt. For the first time, aluminum was obtained by reduction from aluminum chloride with sodium metal:

AlCl 3 + 3 Na = 3 NaCl + Al

At present, it is obtained by electrolysis of molten salts in electrolytic baths (Fig. 46). The electrolyte is a melt containing 85-90% cryolite - complex salt 3NaF·AlF 3 (or Na 3 AlF 6) and 10-15% alumina - aluminum oxide Al 2 O 3 . This mixture melts at about 1000°C.

When dissolved in molten cryolite, alumina behaves like a salt of aluminum and aluminum acid and dissociates into aluminum cations and anions of the acid residue of aluminum acid:

AlAlO 3  Al 3  + AlO 3 3 

Cryolite also dissociates:

Na 3 AlF 6  3 Na  + AlF 6 3 

When an electric current is passed through the melt, aluminum and sodium cations move to the cathode - the graphite body of the bath, covered at the bottom with a layer of molten aluminum obtained in the electrolysis process. Since aluminum is less reactive than sodium, it is the first to be reduced. The reduced aluminum in the molten state is collected at the bottom of the bath, from where it is periodically withdrawn.

Anions AlO 3 3  and AlF 6 3  move to the anode - graphite rods or blanks. At the anode, the AlO 3 3 anion is first discharged -

4 AlO 3 3   12 e  \u003d 2 Al 2 O 3 + 3 O 2

Alumina consumption is replenished all the time. The amount of cryolite practically does not change, only insignificant losses occur due to the formation of carbon tetrafluoride СF 4 on the anode.

The electrolytic production of aluminum requires a lot of electricity (about 20,000 kWh of electricity is consumed to produce 1 ton of aluminum), so aluminum plants are built near power plants.

Application. Aluminum is used very widely. Foil is made from it, which is used in radio engineering and for packaging food products. Steel and cast iron products are coated with aluminum in order to protect them from corrosion: the products are heated to 1000 С in a mixture of aluminum powder (49%), aluminum oxide (49%) and aluminum chloride (2%). This process is called aluminizing.

Aluminized products withstand heating up to 1000 С without being corroded. Aluminum alloys, which are distinguished by their great lightness and strength, are used in the production of heat exchangers, aircraft construction and mechanical engineering.

Aluminum oxide Al 2 O 3 . It is a white substance with a melting point of 2050 С. In nature, aluminum oxide occurs in the form of corundum and alumina. Sometimes there are transparent crystals of corundum of a beautiful shape and color. Corundum dyed red with chromium compounds is called ruby, and blue dyed with titanium and iron compounds is called sapphire. Ruby and sapphire are precious stones. Currently, they are quite easily obtained artificially.

Aluminum oxide has amphoteric properties, but it does not dissolve in water, acids and alkalis. When boiled in a concentrated alkali solution, it only partially goes into solution. Aluminum oxide is converted into a soluble state by fusion with alkalis or potassium pyrosulfate:

AI 2 O 3 + 2 KOH \u003d 2 KAlO 2 + H 2 O

Al 2 O 3 + 3 K 2 S 2 O 7 \u003d 3 K 2 SO 4 + Al 2 (SO 4) 3

The resulting alloys are soluble in water. When aluminum oxide is fused with potash or soda, aluminates, which are readily soluble in water:

Al 2 O 3 + K 2 CO 3 \u003d 2 KAlO 2 + CO 2

Natural corundum is a very hard substance. It is used for the manufacture of emery wheels and grinding powders. Ruby is used to make bushings for clocks and other precision mechanisms.

Alumina is used as a raw material for the production of aluminium. Dehydrated aluminum oxide serves as an adsorbent in the purification and separation of organic substances by chromatography.

Aluminum hydroxide Al(OH) 3 . It is a white substance that, when heated, loses water, turning into aluminum oxide. Aluminum hydroxide has amphoteric properties. Freshly precipitated hydroxide is easily soluble in acids and alkalis (except ammonium hydroxide):

2 Al (OH) 3 + 3 H 2 SO 4 \u003d Al 2 (SO 4) 3 + 6 H 2 O

Al (OH) 3 + 3 KOH \u003d K 3 [Al (OH) 6]

Aluminum hydroxide is a weak base and an even weaker acid, so aluminum salts are in solution only in the presence of an excess of acid, and aluminates only in the presence of an excess of alkali. When solutions are diluted with water, these compounds are strongly hydrolyzed.

Dried aluminum hydroxide, having lost some of its water, does not dissolve in either acids or alkalis, and thus resembles aluminum oxide.

Aluminum hydroxide has the ability to absorb various substances, so it is used in water purification.

Chapter XII. SECOND GROUP OF PERIODIC

ELEMENT SYSTEMS

85. General characteristics of the elements of group II

Group II of the periodic system of elements includes beryllium, alkaline earth metals: magnesium, calcium, strontium, barium and radium (the main subgroup) and a subgroup of zinc: zinc, cadmium, riut (side subgroup). Alkaline earth metals owe their name to the fact that their oxides (earths) form alkaline solutions when dissolved in water.

At the external electronic level of the elements of the main and side subgroups, there are two electrons each (s 2), which they give away, forming compounds with an oxidation state of +2.

All elements of group II are characterized by a relatively low melting point and high volatility. In alkaline earth elements, the solubility of hydroxides increases from magnesium to barium: magnesium hydroxide is almost insoluble in water, calcium hydroxide is slightly soluble, and barium hydroxide is well. The solubility of many salts decreases from magnesium to radium. So, magnesium sulfate dissolves well in water, calcium sulfate is poorly soluble, and strontium, barium and radium sulfates are practically insoluble. The low solubility of radium sulfate is used to isolate radium from its concentrates.

In the zinc subgroup, the amphotericity of oxides decreases from zinc to mercury: zinc hydroxide dissolves well in alkalis, cadmium hydroxide is much worse, and mercury hydroxide is insoluble in alkalis. The activity of elements in this subgroup decreases as their atomic mass increases. So, zinc displaces cadmium and mercury from solutions of their salts, and cadmium displaces mercury.

Beryllium was discovered by L. N. Vauquelin in 1798. Its content in the earth's crust is 3.8·10 - 4%. Beryllium metal is used to make windows for X-ray machines, as it absorbs X-rays 17 times weaker than aluminum. The addition of beryllium to alloys increases their hardness and electrical conductivity. Beryllium compounds can cause very severe lung disease.

Strontium was first isolated in the form of oxide by A. Crawford in 1790, and in pure form was obtained by G. Davy in 1808. Its content in the earth's crust is 0.034%. Strontium nitrate is used in pyrotechnics, while its carbonate and oxide are used in the sugar industry. In nuclear explosions, strontium-90 is formed, the radiation of which is very dangerous, as it causes radiation sickness, leukemia and bone sarcoma.

Barium was discovered by K. V. Scheele in 1774 and G. Devi in ​​1808. Its content in the earth's crust is 0.065%. Of the compounds of barium, its hydroxide, peroxide and some salts are most widely used. Barium hydroxide and chloride are used in laboratory practice, barium peroxide - to produce hydrogen peroxide, nitrate and chlorate - in pyrotechnics, barium sulfate - in fluoroscopy of the digestive organs. Barium compounds are poisonous.

Radium was discovered by M. and P. Curie together with J. Belebn in 1898.

Its content in the earth's crust is 1·10 - 20%. Radium has a natural radioactivity: during its radioactive decay, α-particles, electrons are released and radon is formed. Salts of radium are used for research purposes, as well as to obtain radon, which has healing properties.

Cadmium was discovered by F. Stromeyer in 1817 and, independently of him, by K. Herman, K. Karsten and W. Meissner - in 1818. Its content in the earth's crust is 1.3·10 - 5%. Due to the ability of cadmium to be covered with a protective oxide film, it is used as a stable anti-corrosion coating. Cadmium compounds are poisonous.

86. Magnesium

Magnesium was discovered by G. Davy in 1808.

Being in nature. The magnesium content in the earth's crust is 1.87%. Its compounds are found in various minerals. Magnesium carbonate is part of dolomite CaCO 3 MgCO 3, and magnevit MgCO 3, chloride is part of carnallite KCl MgCl 2 6H 2 O, magnesium sulfate is part of kainite KCl MgSO 4 6H 2 O. A significant amount of magnesium salts found in sea water, giving it a bitter taste.

Physical properties. Magnesium is a silvery-white metal with a density of 1.74 g / cm 3, melts at 651 С, boils at 1110 С. In the cold, magnesium is covered with an oxide film, which protects it from further oxidation by atmospheric oxygen.

Chemical properties. Magnesium is an active metal. If the oxide film on its surface is destroyed, it is easily oxidized by atmospheric oxygen. When heated, magnesium interacts vigorously with halogens, sulfur, nitrogen, phosphorus, carbon, silicon and other elements:

2 Mg + O 2 \u003d 2 MgO (magnesium oxide)

Mg + Cl 2 = MgCl 2 (magnesium chloride)

3 Mg + N 2 = Mg 3 N 2 (magnesium nitride)

3 Mg + 2 P \u003d Mg 3 P 2 (magnesium phosphide)

2 Mg + Si \u003d Mg 2 Si (magnesium silicide)

Magnesium does not dissolve in water, however, when heated, it actively interacts with water vapor:

Mg + H 2 O \u003d MgO + H 2

Magnesium easily takes away oxygen and halogens from many metals, so it is used to obtain rare metals from their compounds:

3Mg + MoO 3 \u003d 3 MgO + Mo

2Mg + ZrCl 4 = 2 MgCl 2 + Zr

It burns in an atmosphere of carbon dioxide:

Mg + CO 2 \u003d MgO + CO

2 Mg + CO 2 \u003d 2 MgO + C

and is highly soluble in acids:

Mg + H 2 SO 4 = MgSO 4 + H 2

4 Mg + 10 HNO 3 \u003d 4 Mg (NO 3) 2 + N 2 O + 5 H 2 O

Receipt. Magnesium is obtained by electrolysis of melts of its salts. The electrolyte is pure dehydrated carnallite, the anode is a graphite rod, and the cathode is iron. The resulting liquid magic floats to the surface and is collected with scoops. During electrolysis, magnesium chloride is added to the electrolyte. Recently, magnesium has also been obtained by reducing it from oxide with calcium carbide, amorphous carbon, or silicon. The reduction process with carbide proceeds at a temperature of 1200, carbon - at 2000, and silicon - at 1200-1300 С. In order to avoid the interaction of metallic magnesium and SiO 2 formed during the reaction, not MgO is introduced into the reaction, but burnt dolomite - a mixture of calcium and magnesium oxides:

MgO + CaC 2 \u003d CaO + Mg + 2 C (1200 C)

MgO + C \u003d Mg + CO (2000 C)

2 MgO + CaO + Si \u003d CaSiO 3 + 2 Mg (1200-1300 C)

Application. Magnesium is used to produce many light alloys, in particular duralumin. The addition of magnesium to cast iron improves the mechanical properties of the latter. Magnesium is used as a reducing agent in the production of rare metals (Nb, Ta, Mo, W, Tl, Zr, Hf, etc.) and some non-metals (for example, Si).

Magnesium oxide MgO. White crystalline substance, insoluble in water. Melts at 2800°C. Has basic properties. It dissolves well in acids:

MgO + H 2 SO 4 \u003d MgSO 4 + H 2 O

when heated, it reacts with acid oxides:

MgO + SiO 2 = MgSO 4

In the laboratory, magnesium oxide can be obtained by burning magnesium metal or by calcining its hydroxide:

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

In industry, MgO is radiated by thermal decomposition of magnesium carbonate:

MgCO 3 \u003d MgO + CO 2

The bulk of magnesium oxide is consumed by the construction industry for the manufacture of magnesite cement and magnesite refractories.

Magnesium hydroxide Mg(OH) 2 . A white substance, insoluble in water, but readily soluble in acids:

Mg (OH) 2 + H 2 SO 4 \u003d MgSO 4 + 2 H 2 O

When carbon dioxide is passed through a suspension of magnesium hydroxide, the latter dissolves to form magnesium bicarbonate:

Mg (OH) 2 + CO 2 \u003d MgCO 3 + H 2 O

MgCO 3 + CO 2 + H 2 O \u003d Mg (HCO 3) 2

Magnesium hydroxide is obtained by the action of alkalis or ammonia on solutions of magnesium salts:

MgCl 2 + 2 KOH \u003d Mg (OH) 2 + 2 KCl

MgCl 2 + 2 NH 4 OH \u003d Mg (OH) 2 + 2 NH 4 Cl

If ammonium salts are added to a solution containing insoluble magnesium hydroxide, the precipitate dissolves. This is due to the fact that ammonium ions bind hydroxyl ions (slightly dissociated ammonium hydroxide is formed):

Mg (OH) 2 + 2 NH 4  \u003d Mg 2  + 2 NH 4 OH

In this way, magnesium can be kept dissolved in ammonia. This solution is called a magnesia mixture and is used for the qualitative and quantitative determination of phosphoric acid ions:

MgCl 2 + 3 NH 4 OH + H 3 RO 4 \u003d MgNH 4 RO 4 + 2 NH 4 Cl + 3 H 2 O

Nitrate, chloride, sulfate, perchlorate, magnesium acetate, as well as acid salts of polybasic acids are readily soluble in water. Other magnesium salts are poorly soluble in water.

87. Calcium

Calcium salts have been known to man for a very long time, but in the free state this metal was obtained by the English chemist G. Davy only in 1808.

Being in nature. The calcium content in the earth's crust is 3.3%. Its most common compounds are the mineral calcite CaCO 3 (the main component of limestone, chalk and marble) and a transparent variety of calcite - Icelandic spar. Calcium carbonate is also part of the mineral dolomite CaCO 3 ·MgCO 3 . Often there are deposits of calcium sulfate in the form of the mineral gypsum CaSO 4 2H 2 O, calcium phosphate - in the form of minerals phosphorite Ca 3 (PO 4) 2 and apatite 3Ca 3 (PO 4) 2 CaF 2 (or Ca 5 (PO 4) 3 F), calcium fluoride - in the form of the mineral fluorspar CaF 2 , and calcium nitrate - in the form of calcium or Norwegian saltpeter Ca(NO 3) 2 . Calcium is also a part of many aluminosilicates, in particular feldspars.

Physical properties. Calcium is a silvery-white malleable metal that melts at 850°C and boils at 1482°C. It is much harder than alkali metals.

Chemical properties. Calcium is an active metal. So, under normal conditions, it easily interacts with atmospheric oxygen and halogens:

2 Ca + O 2 \u003d 2 CaO (calcium oxide)

Ca + Br 2 \u003d CaBr 2 (calcium bromide)

With hydrogen, nitrogen, sulfur, phosphorus, carbon and other non-metals, calcium reacts when heated:

Ca + H 2 \u003d CaH 2 (calcium hydride)

3 Ca + N 2 \u003d Ca 3 N 2 (calcium nitride)

Ca + S = CaS (calcium sulfide)

3 Ca + 2 P \u003d Ca 3 P 2 (calcium phosphide)

Ca + 2 C \u003d CaC 2 (calcium carbide)

Calcium interacts slowly with cold water, and very vigorously with hot water:

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

Calcium can take away oxygen or halogens from oxides and halides of less active metals, i.e. it has reducing properties:

5 Ca + Nb 2 O 5 \u003d CaO + 2 Nb

5 Ca + 2 NbCl 5 \u003d 5 CaCl 2 + 2 Nb

Receipt. Calcium metal is obtained by electrolysis of its molten salts. The electrolyte is a molten mixture of CaCl 2 and CaF 2 in a ratio of 3: 1 by weight. Calcium fluoride is added to lower the melting point of the mixture. Application. Calcium is used in metallurgy to clean iron and steel from oxides, as well as in the production of many rare metals (Tl, Zr, Hf, Nb, Ta, etc.) as a reducing agent for these metals from their oxides and chlorides. An alloy of calcium and lead is used to make bearings and cable sheaths.

Calcium oxide CaO. White substance, melting at about 3000 С, with pronounced basic properties. Reacts well with water, acids and acid oxides:

CaO + H 2 O \u003d Ca (OH) 2

CaO + 2 Hcl \u003d CaCl 2 + H 2 O

CaO + CO 2 \u003d CaCO 3

Under laboratory conditions, calcium oxide can be obtained by oxidizing calcium, as well as by thermal decomposition of its carbonate. In industry, CaO is produced by burning limestone in shaft or rotary kilns at 1000–1100°C. Therefore it is also called burnt or quicklime.

Calcium oxide is used in the building materials industry as a binder.

Calcium hydroxide Ca(OH) 2 . A solid white substance, poorly soluble in water (1.56 g of Ca (OH) 2 dissolves in 1 liter of water at 20 ° C). When calcium oxide is treated with hot water, finely divided calcium hydroxide is obtained - fluff. A saturated aqueous solution of Ca(OH) 2 is called lime water. In air, it becomes cloudy due to interaction with carbon dioxide and the formation of calcium carbonate.

Calcium hydroxide is an alkali. It easily reacts with acids, acid oxides and salts:

Ca (OH) 2 + 2 HCl \u003d CaCl 2 + 2 H 2 O

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

3 Ca(OH) 2 +2 FeCl 3 = 3 CaCl 2 + 2 Fe(OH) 3

The process of interaction of calcium oxide with water is called quenching. Slaked lime, mixed with sand and water, forms a lime mortar used in construction: for bonding bricks when laying walls, for plastering, etc. In the air, slaked lime absorbs carbon dioxide and turns into calcium carbonate.

Zinc in alloys has been known since antiquity. In its purest form

it was received only at the end of the eighteenth century.

Being in nature. The content of zinc in the earth's crust is 8.3·10 - 3%. His connections are quite common. More often than others, the mineral zinc blende ZnS is found, less often - galmium ZnCO 3, silicon-zinc ore Zn 2 SiO 4 H 2 O, zinc spinel ZnO Al 2 O 3 and red zinc ore, or zincite, ZnO.

Physical properties. Zinc is a bluish-white metal with a metallic luster. In air, its surface is covered with an oxide film and tarnishes. Zinc melts at 419.5°C and boils at 913°C. The density of cast solid zinc is 7.13 g/cm 3 , the density of rolled zinc is somewhat higher. At the melting point, the density of zinc is 6.92 g/cm 3 . In the cold, zinc is rather brittle, but at a temperature of 100-150 С it is easy to roll and draw. Easily forms alloys with other metals.

Chemical properties. Zinc is a fairly active metal. It easily interacts with oxygen, halogens, sulfur and phosphorus:

2 Zn + O 2 \u003d 2 ZnO (zinc oxide)

Zn + Cl 2 = ZnCl 2 (zinc chloride)

Zn + S = ZnS (zinc sulfide)

3 Zn + 2 P \u003d Zn 3 P 2 (zinc phosphide)

When heated, it interacts with ammonia, resulting in the formation of zinc nitride:

3 Zn + 2 NH 3 \u003d Zn 2 N 3 + 3 H 2

and also with water:

Zn + H 2 O \u003d ZnO + H 2

and hydrogen sulfide:

Zn + H 2 S \u003d ZnS + H 2

The sulfide formed on the surface of zinc protects it from further interaction with hydrogen sulfide.

Zinc is highly soluble in acids and alkalis:

Zn + H 2 SO 4 \u003d ZnSO 4 + H 2

4 Zn + 10 HNO 3 \u003d 4 Zn (NO 3) 2 + NH 4 NO 3 + 3 H 2 O

Zn + 2 KOH + 2 H 2 O \u003d K 2 + H 2

Unlike aluminum, zinc dissolves in an aqueous solution of ammonia, as it forms a highly soluble ammonia:

Zn + 4 NH 4 OH \u003d (OH) 2 + H 2 + 2 H 2 O

Zinc displaces less active metals from solutions of their salts.

CuSO 4 + Zn = ZnSO 4 + Cu

СdSO 4 + Zn \u003d ZnSO 4 + Сd

Receipt. Zinc is obtained in two ways: pyrometallurgical and hydrometallurgical. In both methods, zinc ore is roasted to convert zinc sulfide to oxide:

2 ZnS + 3 O 2 \u003d 2 ZnO + 2 SO 2

ZnCO 3 \u003d ZnO + CO 2

The sulfur dioxide released is used in the production of sulfuric acid. When zinc is obtained by a pyrometallurgical layer, the resulting zinc cinder (a product of roasting zinc ore) is mixed with coke and heated to 1100-1200 С. Zinc is restored:

ZnO + C = Zn + CO

and at 913 "C it is distilled off.

To obtain zinc by hydrometallurgical slosobom, zinc cinder is dissolved in sulfuric acid, impurities are separated, and zinc is isolated by electrolysis of the sulfuric acid solution (aluminum serves as the cathode, and lead is the anode).

Application. Zinc is used for galvanizing iron in order to protect it from corrosion (zinc sheet), for the manufacture of galvanic cells. Zinc dust is used as a reducing agent for chemical processes. Zinc is a constituent of many alloys.

Zinc oxide ZnO. White powder. It melts at about 2000°C. Poorly soluble in water. It has amphoteric properties. Easily soluble in both acids and alkalis, forming salts of zinc and n to a t y:

ZnO + H 2 SO 4 \u003d ZnSO 4 + H 2 O

ZnO + 2 KOH + 2 H 2 O \u003d K 2

When fused, it interacts with basic and acidic oxides:

ZnO + CaO \u003d CaZnO 2

ZnO + SiO 2 \u003d ZnSiO 3

Zinc oxide is used as a catalyst in many chemical processes. It is also part of zinc white.

Zinc hydroxide Zn(OH) 2 . It has amphoteric properties, easily soluble in acids and alkalis:

Zn (OH) 2 + H 2 SO 4 \u003d ZnSO 4 + 2 H 2 O

Zn (OH) 2 + 2 KOH \u003d K 2

It also dissolves easily in ammonia - zinc ammonia is formed:

Zn (OH) 2 + 4 NH 4 OH \u003d (OH) 2 + 4 H 2 O

Zinc hydroxide is formed when a zinc salt is treated with alkali (but not ammonia) or zincate acid:

ZnSO 4 + 2 KOH \u003d K 2 SO 4 + Zn (OH) 2

K 2 + H 2 SO 4 \u003d K 2 SO 4 + Zn (OH) 2 + 2 H 2 O

Zinc salts. Zinc chloride ZpCl 2 is obtained by dissolving zinc or its oxide in hydrochloric acid. It is very soluble in water (spreads in air). A solution of zinc chloride in hydrochloric acid is used to treat the metal surface during soldering (etching). Zinc chloride forms with hydrochloric acid complex acid H 2 ZnCl 4 , which dissolves metal oxides, but not metals. Innki chloride is used in medicine as an antiseptic.

Zinc sulfide ZnS. Pale yellow powder, sparingly soluble in water. Melts at 1800-1900 С under pressure (sublimates at 1180 С). Easily soluble in acids:

ZnS + 2 HCl \u003d ZnCl 2, + H 2 S

It is part of lithopone, a mineral paint obtained by mixing barium sulfide with zinc sulfate:

BaS + ZnSO 4 = BaSO 4 + ZnS

Litonon is much cheaper than white lead, but less stable in the light. Under the action of ultraviolet and radioactive rays, zinc sulfide glows. Therefore, it is used as a phosphor in cathode ray tubes. Finely ground zinc sulfide (zinc sulfide gray paint) is used for coating metal structures of bridges and machine parts.

Zinc sulfate ZnSO 4 is used in medicine as an antiseptic.

89. Mercury

Mercury was known to the ancient Greeks.

Being in nature. Its content in the earth's crust is

8.3 10 - 6%. Native mercury occurs as inclusions in the rock. There is also mercury sulfide HgS, called cinnabar.

Physical properties. Mercury is a silvery-white liquid metal that solidifies at -38.84°C and boils at 356.95°C. In the solid state, it has good malleability and elasticity. Mercury dissolves many metals, forming amalyams. In them, metals behave as in the free state, but become less active (formation of amalgam reduces activity similarly to dilution). Mercury vapor is highly toxic. Mercury is not excreted from the human body.

Chemical properties. Mercury is an inactive metal. It interacts with oxygen only when heated:

2 Hg + O 2 \u003d 2 HgO

Mercury reacts with chlorine in the cold, forming mercury chloride, or sublimate:

Hg + Cl 2 = HgCl 2

Mercury easily interacts with powdered sulfur, forming a very strong compound - mercury sulfide:

This reaction is used to bind spilled mercury: a place where spilled mercury is suspected is sprinkled with sulfur powder.

Mercury does not dissolve in water and alkalis. It dissolves in oxidizing acids; in concentrated sulfuric acid when heated, and in nitric acid - in the cold. Depending on the amount of mercury, mercury salts are formed in the oxidation state +1 and +2:

Hg + 2 H 2 SO 4 \u003d HgSO 4 + SO 2 + 2 H 2 O

3 Hg + 8 HNO 3 \u003d 3 Hg (NO 3) 2 + 2 NO + 4 H 2 O

Hg + Hg (NO 3) 2 = Hg 2 (NO 3) 2

Mercury (II) in HgCl 2 chloride is reduced by metallic mercury to mercury (I):

HgCl 2 + Hg \u003d Hg 2 Cl 2 (calomel)

Receipt. Mercury is obtained from cinnabar by calcining it in air or heating it with iron or calcium oxide:

HgS + O 2 = Hg + SO 2

HgS + Fe = Hg + FeS

4 HgS + 4 CaO \u003d 4 Hg + 3 CaS + CaSO 4

Mercury is easily distilled off.

Application. Metallic mercury is used in various devices, such as pressure regulators, quartz lamps, thermometers, diffusion vacuum pumps, etc. It is also used to obtain paints, mercury fulminate, mercury ointments against skin diseases. Mercury amalgams are used as reducing agents. Significant amounts of mercury are used in the electrochemical industry (mercury cathodes) and polarography.

Mercury oxide HgO. The substance is yellow or red. When heated, it easily decomposes into oxygen and mercury. Mercury oxide has only basic properties. It can dissolve in acids, with which mercury forms easily soluble salts:

HgO + 2 Hcl \u003d HgCl 2 + H 2 O

HgO + 2 HNO 3 \u003d Hg (NO 3) 2 + H 2 O

Mercury oxide does not dissolve in water and, under the action of alkalis on solutions of mercury (II) salts, precipitates:

HgCl 2 + 2 KOH \u003d 2 KCl + HgO + H 2 O

Hg (NO 3) 2 + 2 KOH \u003d 2 KNO 3 + HgO + H 2 O

In mercury compounds with an oxidation state of +1, two mercury atoms are joined together by a covalent bond. Under the action of alkalis on mercury (I) salts, metallic mercury and mercury oxide (II) precipitate:

Hg (NO 3) 2 + 2 KOH \u003d 2 KNO 3 + HgO + Hg + H 2 O

Mercury salts are mainly used as catalysts for many chemical processes. So, sublimate HgCl 2 catalyzes the hydrochlorination of acetylene:

HCCH + HCl -- H 3 C = CHCl

Mercury sulfate HgSO 4 used as a catalyst in the hydration of acetylene according to the Kucherov reaction:

HCCH + H 2 O -- CH 3 CHO

Sparingly soluble Hg 2 Cl 2 calomel is used in the manufacture of standard electrodes for electrometric instruments.

Table 19 - Characteristics of elements 3Ap / group

Aluminum is in the main subgroup of Group III of the Periodic Table. The atoms of the elements of the subgroup in the ground state have the following structure of the outer electron shell: ns 2 np 1 . At the outer energy level of atoms, there are free p-orbitals, which allows atoms to pass into an excited state. In an excited state, the atoms of these elements form three covalent bonds or completely donate three valence electrons, showing an oxidation state of +3.

Aluminum is the most common metal on Earth: its mass fraction in the earth's crust is 8.8%. The bulk of natural aluminum is part of aluminosilicates - substances, the main components of which are oxides of silicon and aluminum. Aluminosilicates are part of many rocks and clays.

Properties: Al is a silvery white metal, It is a fusible and light metal. It has high plasticity, good electrical and thermal conductivity. Al is a reactive metal. However, its activity under normal conditions is somewhat reduced due to the presence of a thin oxide film that forms on the surface of the metal when it comes into contact with air.

1. Interaction with non-metals. Under normal conditions, aluminum reacts with chlorine and bromine:

2Al + 3Cl 2 = 2AlCl 3

When heated, aluminum interacts with many non-metals:

4Al + 3O 2 \u003d 2Al 2 O 3

2Al + 3I 2 = 2AlI 3

2Al + N 2 = 2AlN

4Al + 3C \u003d Al 4 C 3

2. Interaction with water. Due to the protective oxide film on the surface, aluminum is stable in water. However, when this film is removed, an energetic interaction occurs:

2Al + 6H 2 O \u003d 2Al (OH) 3 + 3H 2

2. Interaction with acids. Aluminum interacts with hydrochloric and dilute sulfuric acids:

2Al + 6HCl = 2AlCl 3 + 3H 2

2Al + 3H 2 SO 4 \u003d Al 2 (SO 4) 3 + 3H 2

Nitric and concentrated sulfuric acids passivate aluminum: under the action of these acids, the thickness of the protective film on the metal increases, and it does not dissolve.

4. Interaction with alkalis. Aluminum interacts with alkali solutions with the release of hydrogen and the formation of a complex salt:

2Al + 6NaOH + 6H 2 O = 2Na 3 + 3H 2

5. Recovery of metal oxides. Aluminum is a good reducing agent for many metal oxides:

2Al + Cr 2 O 3 \u003d Al 2 O 3 + 2Cr

8Al + 3Fe 3 O 4 = 4Al 2 O 3 + 9Fe

aluminum oxide and hydroxide. Aluminum oxide, or alumina, Al 2 O 3 is a white powder. Aluminum oxide can be obtained by burning metal or by calcining aluminum hydroxide:

2Al(OH) 3 \u003d Al 2 O 3 + 3H 2 O

Aluminum oxide is practically insoluble in water. The hydroxide Al (OH) 3 corresponding to this oxide is obtained by the action of ammonium hydroxide or alkali solutions, taken in deficiency, on solutions of aluminum salts:

AlCl 3 + 3NH 3 ∙ H 2 O \u003d Al (OH) 3 ↓ + 3NH 4 Cl

The oxide and hydroxide of this metal are amphoteric, i.e. exhibit both basic and acidic properties.

Basic properties:

Al 2 O 3 + 6HCl \u003d 2AlCl 3 + 3H 2 O

2Al(OH) 3 + 3H 2 SO 4 = Al 2 (SO 4) 3 + 6H 2 O

Acid properties:

Al 2 O 3 + 6KOH + 3H 2 O \u003d 2K 3

2Al(OH) 3 + 6KOH = K 3

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

Production. Aluminum is produced by the electrolytic method. It cannot be isolated from aqueous solutions of salts, because is a very active metal. Therefore, the main industrial method for obtaining metallic aluminum is the electrolysis of a melt containing aluminum oxide and cryolite.

Application. Metallic aluminum is widely used in industry, in terms of production it ranks second after iron. The bulk of aluminum goes to the manufacture of alloys:

Duralumin is an aluminum alloy containing copper and small amounts of magnesium, manganese and other components. Duralumins are light, strong and corrosion-resistant alloys. Used in aircraft and mechanical engineering.

Magnalin is an alloy of aluminum and magnesium. Used in aircraft and mechanical engineering, in construction. Resistant to corrosion in sea water, so it is used in shipbuilding.

Silumin is an aluminum alloy containing silicon. Good for casting. This alloy is used in automotive, aircraft and mechanical engineering, the production of precision instruments.

Aluminum is a ductile metal, so it is used to make thin foil used in the manufacture of radio engineering products and for packaging goods. Wires are made of aluminum, silver paints.

Assignments with a professional focus

1. To clean the root crops from the skin after washing, they are scalded with a boiling solution of soda ash (W = 4%). With an excess of hydrochloric acid in the gastric juice, the animals are fed with a solution of baking soda. Write the formulas for these substances. Name other areas of application of sodium and potassium salts in agricultural practice, in everyday life.

2. Potassium iodide is widely used to feed animals with microelements and to remove excess flowers on apple trees. Write an equation for the reaction of obtaining potassium iodide, indicate the oxidizing agent and reducing agent.

3. Why is wood ash (the ash contains potassium ions K + and carbonate - CO 3 2- ions) used for fertilizing fields recommended to be stored indoors or under a canopy? Write the equations for the reactions that occur when the ash is moistened.

4. Too much acidity of the soil has a detrimental effect on the plant. In this case, it is necessary to carry out liming of the soil. The introduction of CaCO 3 limestone into the soil reduces acidity. Write the equation for the reaction taking place in this case.

5. The acidity of the soil does not change from the introduction of superphosphate. However, the acidity of superphosphate containing excess phosphoric acid is harmful to plants. CaCO 3 is added to neutralize it. It is impossible to add Ca (OH) 2, because superphosphate will turn into a compound that is difficult for plants to digest. Write equations for the corresponding reactions.

6. To combat pests of grain, fruits and vegetables, chlorine is used at the rate of 35 g per 1 m 3 of the room. Calculate the mass of sodium chloride sufficient to treat 300 m 3 of the room with chlorine obtained by electrolysis of a melt of this salt.

7. For every 100 quintals of root crops and sugar beet tops, approximately 70 kg of potassium oxide is removed from the soil. What mass of sylvinite KCl Na Cl containing potassium chloride with a mass fraction of 0.56 can compensate for these losses?

8. For feeding potatoes, a solution of potassium chloride with a mass fraction of 0.04 is used. Calculate the mass of potash fertilizer (KCl) required to obtain 20 kg of such a solution.

9. When obtaining a nutrient solution for plant nutrition, 1 g of KNO 3, 1 g of MgSO 4, 1 g of KH 2 PO 4, 1 g of Ca (NO 3) 2 are taken per 400 ml of water. Calculate the mass fraction (in%) of each substance in the resulting solution.

10. To preserve wet grain from decay, it is treated with sodium hydrosulfate NaHSO 4. Calculate the mass of sodium hydrosulfate, which is obtained by reacting 120 g of sodium hydroxide with a solution of sulfuric acid.

11. Which fertilizer contains more potassium: potassium nitrate (KNO 3), potash (K 2 CO 3) or potassium chloride (KCl)?

12. Calcium cyanamide is used for pre-harvest deleafing of cotton during its mechanical harvesting. Find the formula of this compound, knowing that the mass fractions of calcium, carbon and nitrogen are 0.5, respectively; 0.15; 0.35.

13. When analyzing wood ash used in animal husbandry as a feed for livestock, it was found that 70 g of ash contains 18.4 g of calcium, 0.07 g of phosphorus and 2.3 g of sodium. Calculate the mass fraction (in%) of each element in the specified top dressing.

14. How much limestone containing 90% calcium carbonate should be applied per 30 hectares if liming is carried out at the rate of 4 tons of CaO per hectare.

15. There are: a) pure ammonium nitrate, b) technical sylvinite containing 33% potassium. By mixing these materials, it is necessary to obtain one ton of nitrogen-potassium fertilizer containing 15% nitrogen. What quantities of both materials should be mixed and what percentage of potassium will such a mixture contain?

4.9 Section: Major transition metals

Purpose: To study the properties of metals of secondary subgroups and their compounds

Transition metals are elements of secondary subgroups of the periodic system.

8950 0

Group 14 includes C, Si, Ge, Sn, Pb (Tables 1 and 2). Like the elements of the 3A subgroup, these are p-elements with a similar electronic configuration of the outer shell - s 2 p 2. As you move down the group, the atomic radius increases, causing the bond between the atoms to weaken. Due to the increasing delocalization of the electrons of the outer atomic shells, the electrical conductivity increases in the same direction, so the properties of the elements change from non-metallic to metallic. Carbon (C) in the form of a diamond is an insulator (dielectric), Si and Ge are semimetals, Sn and Pb are metals and good conductors.

Table 1. Some physical and chemical properties of metals of group 14

	Name		Refers, at. weight	Electronic formula	Radius, pm	Main isotopes (%)
	Carbon Carbon [from lat. carbo - coal]				covalent 77 with double bond 67, with triple bond 60	14 C (traces)
	Silicon Silicon [from lat. silicis - flint]				atomic 117, covalent 117
	Germanium Germanium [from lat. Germany]			3d 10 4s 2 4p 2	atomic 122.5, covalent 122
	Tin Tin [from the Anglo-Sax. tin, lat. stannum]			4d 10 5s 2 5p 2	atomic 140.5, covalent 140
	Lead Lead [from Anglo-Sax. lead, lat. plumbum]			4f 14 5d 10 6s 2 6r 2	atomic 175, covalent 154

All elements of this group form compounds with an oxidation state of +4. The stability of these compounds decreases when moving to the lower part of the group, while, as in divalent compounds, it, on the contrary, increases with such a movement. All elements except Si, also form compounds with a valence of +2, which is due to " inert pair effect»: by pulling in a pair of external s-elements into the inner electron shell due to worse shielding of outer electrons d- and f-electrons compared to s- and R-electrons of the inner shells of large atoms of the lower members of the group.

The properties of the elements of this group made it possible to use them as anti-algae coatings (AP) for ships. The first such coatings used Pb, then began to apply sn(in the form of a bis-tributyl organotin radical associated with a carbon polymer). For environmental reasons, in 1989 the use of these, as well as other toxic metals in PP ( Hg, Cd, As) was banned, replaced by PP based on organosilicon polymers.

Table 2. The content in the body, toxic (TD) and lethal doses (LD) of metals of the 14th group

	In the earth's crust (%)	In the ocean (%)	In the human body
			Average (with a body weight of 70 kg)			Blood (mg/l)
							usually non-toxic, but in the form of CO and CN cyanides it is very toxic
		(0.03-4.09)x10 -4					Non-toxic
		(0.07-7)x10 -10					Non-toxic
		(2.3-8.8)x10 -10			(0.33-2.4)x10 -4		TD 2 g, LD nd, some organotin. compounds are highly toxic
					(0.23-3.3)x10 -4		TD 1 mg, LD 10 g

Carbon (C) - different from all other elements of the so-called catenation, that is, the ability to form compounds in which its atoms are linked to each other in long chains or rings. This property explains the formation of millions of compounds called organic, which is devoted to a separate section of chemistry - organic chemistry.

The ability of carbon to catenation is explained by several features:

Firstly, strength connections C - C. Thus, the average enthalpy of this bond is about 350 kJ/mol, while the enthalpy of bond Si - Si— only 226 kJ/mol.

Secondly, the unique ability of carbon atoms to hybridization: education 4 sp 3 orbitals with a tetrahedral orientation (ensuring the formation of simple covalent bonds), or 3 sp 2 orbitals oriented in the same plane (providing the formation of double bonds), or 2 sp-orbitals with a linear orientation (providing the formation of triple bonds).

Thus, carbon can form 3 types of coordination environment: linear for two- and three-atomic molecules, when the CN of the element is 2, plane triangular for graphite molecules, fullerenes, alkenes, carbonyl compounds, benzene ring, when the CN is 3, and tetrahedral for alkanes and their derivatives with CN = 4.

In nature, carbon occurs in the form of allotropic, that is, various structural forms (graphite, diamond, fullerenes), as well as in the form of limestone and hydrocarbon raw materials (coal, oil and gas). It is used in the form of coke in steel smelting, carbon black in printing, activated carbon in the purification of water, sugar, etc.

In 2010, the Nobel Prize in Physics was awarded for the study of a unique shape FROM- graphene. The laureates - immigrants from Russia - A. Geim and K. Novoselov managed to obtain this material from graphite. It is a two-dimensional crystal, that is, it looks like a grid of C atoms one atom thick, wave-like structure, which ensures the stability of the crystal. Its properties are very promising: it is the thinnest transparent material of all currently known, moreover, it is extremely strong (about 200 times stronger than steel), has electrical and thermal conductivity. At room temperature, its electrical resistance is the lowest among all known conductors. In the near future, graphene will be used to create ultra-high-speed computers, flat-panel screens and solar panels, as well as sensitive gas detectors that respond to several gas molecules. Other areas of its use are not excluded.

In the form of an oxide ( SO) and cyanides ( CN-) carbon is very toxic because it disrupts the processes of respiration. The mechanisms of biological action of these compounds are different. Cyanide inhibits the respiratory enzyme cytochrome oxidase quickly contacting Xi- the active center of the enzyme, blocking the electron flow at the end of the respiratory chain. SO, being a Lewis base, binds to an atom Fe in the hemoglobin molecule is stronger than O 2 , forming carbonylhemoglobin devoid of the ability to bind and carry O 2. Ability SO form links with d-metals in low oxidation states leads to the formation of diverse carbonyl compounds. For example, Fe in a very toxic substance - psitacarbopile Fe(CO) 5 has a zero oxidation state, and in the complex [ Fe(CO) 4 ] 2- is the oxidation state -2 (Fig. 1).

Rice. one.

Stabilization of a metal atom in a low oxidation state in complexes with SO due to the ability of carbon to protrude due to the structure of low-lying R*-orbitals in the role acceptor ligand. These orbitals overlap with the occupied orbitals of the metal, forming a coordination R-bond in which the metal acts donor electrons. This is one of the few exceptions to the general rule for the formation of CSs, where the electron acceptor is a metal.

It makes no sense to describe the properties of carbon in more detail, since, as a rule, it is not only not determined in multielement analysis, but its admixture in the sample is also considered undesirable and subject to maximum removal during sample preparation. In optical emission analysis, it gives a very wide spectrum, increasing the noise background and thereby reducing the sensitivity limit for detecting the elements being determined. In mass spectrometry, organic molecules form a large number of fragments of molecules with different molecular weights, which give significant interference in the analysis. Therefore, in the vast majority of cases, all carbon-containing substances are removed during sample preparation.

Silicon (Si) - semimetal. When silica is reduced ( SiO 2) black amorphous is formed by carbon Si. crystals Si high purity resemble a gray-blue metal. Silicon is used in semiconductors, alloys and polymers. It is important for some forms of life, for example, for building shells in diatoms; possibly important for the human body. Some silicates are carcinogenic, some cause silicosis.

In all connections Si tetravalent, forms chemical bonds of a covalent nature. The most common oxide SiO 2. Despite the chemical inertness and insolubility in water, when ingested, it can form silicic acids and organosilicon compounds with implicit biological properties. Toxicity SiO 2 depends on the dispersion of particles: the smaller they are, the more toxic, although the correlations between the solubility of various forms SiO 2 and silicogenicity is not observed. The relationship of the toxicity of silicic acids with Si proves the complete inertness of diamond dust of the same fineness.

Recently, it has been noted that in biological media, silicic acids are involved in the formation hydroxylaluminosilicates, and this phenomenon cannot be explained by the relation Si-C, no connection Si-O-C. As industrial use expands Al and its compounds via aluminosilicates Al increasingly involved in many biochemical reactions. In particular, functional oxygen- and fluorine-containing groups easily form highly stable complex compounds with Al perverting their metabolism.

The most studied among organosilicon compounds silicones- polymers, the skeleton of the molecule of which consists of alternating interconnected atoms Si and O 2. To atoms Si in silicones, alkyl or aryl groups are attached. Availability Si in organosilicon compounds, it radically changes the properties of substances when they do not contain it. For example, common polysaccharides can be isolated and purified using strong ethanol, which precipitates the polysaccharide out of solution. Silicon-containing carbohydrates, on the other hand, do not precipitate even in 90% ethanol. The classification of organosilicon compounds is presented in Table. 3.

Table 3 Silicone polymers

	Name and structure	Note
		Consist only of Si. The binding energy of a carbon chain C - C is 58.6, and Si - Si 42.5 kcal/mol, and therefore polyorganosilanes are unstable.
		Bond energy Si-O 89.3 kcal/mol. Therefore, these polymers are strong, resistant to temperature and oxidative degradation. This class of polymers is very diverse in structure. Linear polysiloxanes are widely used as synthetic elastic and heat-resistant rubbers.
		Atoms in the main chain Si separated by chains of carbon atoms.
		The main chain contains siloxane groups separated by carbon chains.
		The backbone is made up of atoms FROM, and the atoms Si contained in side groups or offshoots.
		Macromolecular chains include atoms Si, O and metals, where M = Al, Ti, Sb, Sn, B.

The most likely development mechanism silicosis consider the destruction of phagocytes that have captured particles SiO 2. When interacting with lysosomes, silicon particles destroy lysosomes and the phagocyte cell itself, causing the release of enzymes and fragments of organelle molecules. They interact with other phagocytes, that is, a chain process of phagocyte death is launched. If there is a certain amount of silicic acids in the cell, this process is accelerated. The accumulation of dead macrophages initiates the production of collagen in the surrounding fibroblasts, as a result of which sclerosis develops in the focus.

Colloidal silicic acid is a powerful hemolytic, changes the ratio of serum proteins, inhibits a number of respiratory and tissue enzymes, disrupts the metabolism of many substances, including phosphorus. Recently, much attention has been paid silylium ions (R 3 Si+). They show the unique ability of the atom Si to expand its coordination sphere, in the form of increasing its electrophilicity. It interacts with any nucleophiles, including ions of opposite charge (including reactive metabolic intermediates) and solvent molecules. Therefore, in condensed phases, they become “elusive” and it is difficult to detect them (Kochina et al., 2006).

Organosilicon polymers (OSPs) were first used as anti-algae self-polishing ship hull coatings (Tsukerman, Rukhadze, 1996). However, then various methods were proposed for the use of COPs in other sectors of the national economy, in particular, in medicine as strong bone prostheses.

Germanium (Ge) — amphoteric semimetal; at ultra-high purity, it appears as brittle silver-white crystals. It is used in semiconductors, alloys and special glasses for infrared optics. It is considered a biological stimulant. In compounds, it exhibits an oxidation state of +2 and +4.

Absorption of dioxide and halides Ge weak in the intestine, but in the form of germanates M 2 GeO 4 is somewhat improved. Germanium does not bind to plasma proteins, and is distributed between erythrocytes and plasma in a ratio of approximately 2:1. Quickly (half-life of about 36 hours) is excreted from the body. Generally low toxicity.

Tin (Sn) - soft, ductile metal. It is used in lubricants, alloys, solder, as an additive to polymers, in the composition of paints for antifouling coatings, in the composition of volatile organotin compounds highly toxic to lower plants and animals. In the form of inorganic compounds, it is non-toxic.

Has two enantiotrope, "gray" (b) and "white" (c) tin, that is, different allotropic forms that are stable in a certain range of conditions. The transition temperature between these forms at a pressure of 1 atm. equal to 286.2°K (13.2°C). White tin has a distorted gray modification structure with CN = 6 and a density of 7.31 g/cm 3 . It is stable under normal conditions, and at low temperature it slowly transforms into a form having a diamond-like structure with CN = 4 and a density of 5.75 g/cm 3 . Such a change in the density of the metal depending on the temperature of the medium is extremely rare and can cause dramatic consequences. For example, in the conditions of cold winters, tin buttons on the uniforms of soldiers were destroyed, and in 1851, in the church of Seitz, the tin pipes of the organ turned into powder.

In the body it is deposited in the liver, kidneys, bones, muscles. With tin poisoning, erythropoiesis decreases, which is manifested by a decrease in hematocrit, hemoglobin, and the number of red blood cells. There has also been an inhibition 5-aminolevulinate dehydratase, one of the enzymes in the heme biosynthesis chain, as well as liver enzymes glutathione reductase and dehydrogenase glucose-6-phosphate, lactate and succinate. Apparently sn excreted from the body as part of complexes with SH containing substrates.

Lead (Pb) - soft, malleable, ductile metal. In moist air it is covered with an oxide film, resistant to oxygen and water. Used in batteries, cables, paints, glass, lubricants, gasoline and radiation protection products. It is a toxic metal of hazard group 1, as it accumulates in the body in bone tissue with impaired renal function and the cardiovascular system. In developed countries, its content is controlled with mandatory medical examination of the population. Causes various diseases.

Medical bioinorganics. G.K. Barashkov