Most of the metals are in the periodic table. L.p.vanova, teacher of chemistry at the Novinsky secondary school (Astrakhan region). General chemical properties of metals

1. What features of the structure of metal atoms determine their reducing properties?

The reducing properties of metals are determined by the ability to donate electrons from the outer layer. The easier an atom donates electrons to the outer layer, the stronger the reducing agent it is.

2. Name the chemical element that forms a simple substance - the most active metal. Justify your choice.

The most active metal is francium (Fr).

Francium most easily donates an electron to the outer layer. It has the largest atomic radius, so the energy of interaction of the atomic nucleus with the outer electron shell small.

3. How does the statement that metals exhibit only reducing properties and, therefore, oxidize at the same time, agree with the process that can be reflected using the equation: Name this process. In what forms of existence of the chemical element does copper appear? For what form of existence chemical elements Is the above statement correct?

Metals exhibit reducing properties in zero degree oxidation, i.e. the metal itself can only be a reducing agent. The above process is an example of the oxidation of Cu2+ to Cu0. In this example, copper acts as a cation.

Introduction

Metals are simple substances that under normal conditions have characteristic properties: high electrical and thermal conductivity, the ability to reflect light well (which causes their brilliance and opacity), the ability to take the desired shape under the influence of external forces (plasticity). There is another definition of metals - these are chemical elements characterized by the ability to donate external (valence) electrons.

Of all the known chemical elements, about 90 are metals. Most inorganic compounds are metal compounds.

There are several types of classification of metals. The most clear is the classification of metals in accordance with their position in the periodic system of chemical elements - chemical classification.

If, in the "long" version of the periodic table, a straight line is drawn through the elements boron and astatine, then metals will be located to the left of this line, and non-metals to the right of it.

From the point of view of the structure of the atom, metals are divided into intransitive and transitional. Non-transition metals are located in the main subgroups of the periodic system and are characterized by the fact that in their atoms there is a sequential filling of the electronic levels s and p. Non-transition metals include 22 elements of the main subgroups a: Li, Na, K, Rb, Cs, Fr, Be, Mg, Ca, Sr, Ba, Ra, Al, Ga, In, Tl, Ge, Sn, Pb, Sb, Bi, Po.

Transition metals are located in side subgroups and are characterized by the filling of d - or f-electronic levels. The d-elements include 37 metals of secondary subgroups b: Cu, Ag, Au, Zn, Cd, Hg, Sc, Y, La, Ac, Ti, Zr, Hf, Rf, V, Nb, Ta, Db, Cr, Mo , W, Sg, Mn, Tc, Re, Bh, Fe, Co, Ni, Ru, Rh, Pd, Os, Ir, Pt, Hs, Mt.

The f-elements include 14 lanthanides (Ce, Pr, Nd, Pm, Sm, Eu, Gd, Tb, Dy, Ho, Er, Tm, Yb, Lu) and 14 actinides (Th, Pa, U, Np, Pu, Am, Cm, Bk, Cf, Es, Fm, Md, No, Lr).

Among the transition metals, rare earth metals (Sc, Y, La and lanthanides), platinum metals (Ru, Rh, Pd, Os, Ir, Pt), transuranium metals (Np and elements with a higher atomic mass) are also distinguished.

In addition to chemical, there is also, although not generally accepted, but long established technical classification of metals. It is not as logical as chemical one - it is based on one or another practically important feature of the metal. Iron and alloys based on it are classified as ferrous metals, all other metals are non-ferrous. There are light (Li, Be, Mg, Ti, etc.) and heavy metals (Mn, Fe, Co, Ni, Cu, Zn, Cd, Hg, Sn, Pb, etc.), as well as groups of refractory (Ti, Zr , Hf, V, Nb, Ta, Cr, Mo, W, Re), precious (Ag, Au, platinum metals) and radioactive (U, Th, Np, Pu, etc.) metals. In geochemistry, scattered (Ga, Ge, Hf, Re, etc.) and rare (Zr, Hf, Nb, Ta, Mo, W, Re, etc.) metals are also distinguished. As you can see, there are no clear boundaries between groups.

Historical reference

Despite the fact that the life of human society without metals is impossible, no one knows exactly when and how a person first began to use them. The most ancient writings that have come down to us tell about primitive workshops in which metal was smelted and products were made from it. This means that man mastered metals earlier than writing. Excavating ancient settlements, archaeologists find tools of labor and hunting that people used in those distant times - knives, axes, arrowheads, needles, fish hooks and much more. The older the settlements, the rougher and more primitive were the products of human hands. The most ancient metal products were found during excavations of settlements that existed about 8 thousand years ago. These were mainly jewelry made of gold and silver and arrowheads and spears made of copper.

The Greek word "metallon" originally meant mines, mines, hence the term "metal" came from. In ancient times, it was believed that there were only 7 metals: gold, silver, copper, tin, lead, iron and mercury. This number correlated with the number of planets then known - the Sun (gold), the Moon (silver), Venus (copper), Jupiter (tin), Saturn (lead), Mars (iron), Mercury (mercury) (see figure). According to alchemical concepts, metals originated in the bowels of the earth under the influence of the rays of the planets and gradually improved, turning into gold.

Man first mastered native metals - gold, silver, mercury. The first artificially obtained metal was copper, then it was possible to master the production of an alloy of copper with salting - bronze, and only later - iron. In 1556, a book by the German metallurgist G. Agricola "On Mining and Metallurgy" was published in Germany - the first detailed guide to obtaining metals that has come down to us. True, at that time lead, tin and bismuth were still considered varieties of the same metal. In 1789, the French chemist A. Lavoisier, in his manual on chemistry, gave a list of simple substances, which included all the then known metals - antimony, silver, bismuth, cobalt, tin, iron, manganese, nickel, gold, platinum, lead, tungsten and zinc. With the development of chemical research methods, the number of known metals began to increase rapidly. In the 18th century 14 metals were discovered, in the 19th century. - 38, in the 20th century. - 25 metals. In the first half of the 19th century satellites of platinum were discovered, alkali and alkaline earth metals were obtained by electrolysis. In the middle of the century, cesium, rubidium, thallium and indium were discovered by spectral analysis. The existence of metals predicted by D. I. Mendeleev on the basis of his periodic law (these are gallium, scandium and germanium) was brilliantly confirmed. The discovery of radioactivity at the end of the 19th century. led to the search for radioactive metals. Finally, by the method of nuclear transformations in the middle of the 20th century. radioactive metals that do not exist in nature, in particular transuranium elements, were obtained.

Physical and chemical properties of metals.

All metals are solids (except mercury, which is liquid under normal conditions), they differ from non-metals in a special type of bond (metallic bond). Valence electrons are loosely bound to a particular atom, and inside every metal there is a so-called electron gas. Most metals have a crystalline structure, and a metal can be thought of as a "rigid" crystal lattice of positive ions (cations). These electrons can more or less move around the metal. They compensate for the repulsive forces between cations and thus bind them into a compact body.

All metals have high electrical conductivity (i.e., they are conductors, unlike non-dielectric non-metals), especially copper, silver, gold, mercury and aluminum; the thermal conductivity of metals is also high. A distinctive property of many metals is their ductility (ductility), as a result of which they can be rolled into thin sheets (foil) and drawn into wire (tin, aluminum, etc.), however, there are also quite brittle metals (zinc, antimony, bismuth).

In industry, not pure metals are often used, but their mixtures, called alloys. In an alloy, the properties of one component usually successfully complement the properties of another. So, copper has a low hardness and is of little use for the manufacture of machine parts, while copper-zinc alloys, called brass, are already quite hard and are widely used in mechanical engineering. Aluminum has good ductility and sufficient lightness (low density), but is too soft. On its basis, an alloy of ayuralumin (duralumin) is prepared, containing copper, magnesium and manganese. Duralumin, without losing the properties of its aluminum, acquires high hardness and is therefore used in aviation technology. Alloys of iron with carbon (and additions of other metals) are well-known cast iron and steel.

Metals vary greatly in density: for lithium it is almost half that of water (0.53 g/cm3), while for osmium it is more than 20 times higher (22.61 g/cm3). Metals also differ in hardness. The softest - alkali metals, they are easily cut with a knife; the hardest metal - chromium - cuts glass. The difference in melting points of metals is great: mercury is a liquid under normal conditions, cesium and gallium melt at the temperature of the human body, and the most refractory metal, tungsten, has a melting point of 3380 ° C. Metals whose melting point is above 1000 ° C are classified as refractory metals, below - as fusible. At high temperatures, metals are capable of emitting electrons, which is used in electronics and thermoelectric generators for the direct conversion of thermal energy into electrical energy. Iron, cobalt, nickel and gadolinium, after being placed in a magnetic field, are able to permanently maintain a state of magnetization.

Metals also have some chemical properties. Metal atoms give up valence electrons relatively easily and pass into positively charged ions. Therefore, metals are reducing agents. This, in fact, is their main and most common chemical property.

Obviously, metals as reducing agents will react with various oxidizing agents, among which there may be simple substances, acids, salts of less active metals, and some other compounds. Compounds of metals with halogens are called halides, with sulfur - sulfides, with nitrogen - nitrides, with phosphorus - phosphides, with carbon - carbides, with silicon - silicides, with boron - borides, with hydrogen - hydrides, etc. Many of these compounds found important applications in the new technology. For example, metal borides are used in radio electronics, as well as in nuclear technology as materials for regulating and protecting against neutron radiation.

Under the action of concentrated oxidizing acids, a stable oxide film is also formed on some metals. This phenomenon is called passivation. So, in concentrated sulfuric acid, metals such as Be, Bi, Co, Fe, Mg, and Nb are passivated (and do not react with it), and in concentrated nitric acid - metals Al, Be, Bi, Co, Cr, Fe, Nb, Ni, Pb, Th and U.

The more to the left of the metal in this row, the greater the reducing properties it has, i.e., it is more easily oxidized and goes into solution in the form of a cation, but it is more difficult to recover from the cation to the free state.

One non-metal, hydrogen, is placed in a series of voltages, since this makes it possible to determine whether this metal will react with acids - non-oxidizing agents in an aqueous solution (more precisely, it will be oxidized by hydrogen cations H +). For example, zinc reacts with hydrochloric acid, since in the series of voltages it is to the left (before) hydrogen. On the contrary, silver is not transferred into solution by hydrochloric acid, since it is in the series of voltages to the right (after) hydrogen. Metals behave similarly in dilute sulfuric acid. Metals that are in the series of voltages after hydrogen are called noble (Ag, Pt, Au, etc.)

An undesirable chemical property of metals is their electrochemical corrosion, i.e. active destruction (oxidation) of the metal upon contact with water and under the influence of oxygen dissolved in it (oxygen corrosion). For example, the corrosion of iron products in water is widely known.

Particularly corrosive can be the place of contact of two dissimilar metals - contact corrosion. Between one metal, such as Fe, and another metal, such as Sn or Cu, placed in water, a galvanic couple occurs. The flow of electrons goes from the more active metal, which is to the left in the voltage series (Fe), to the less active metal (Sn, Cu), and the more active metal is destroyed (corrodes).

It is because of this that the tinned surface of cans (tin-plated iron) rusts when stored in a humid atmosphere and carelessly handled (iron quickly collapses after even a small scratch appears, allowing contact of iron with moisture). On the contrary, the galvanized surface of an iron bucket does not rust for a long time, because even if there are scratches, it is not iron that corrodes, but zinc (a more active metal than iron).

Corrosion resistance for a given metal increases when it is coated with a more active metal or when they are fused; for example, coating iron with chromium or making alloys of iron with chromium eliminates the corrosion of iron. Chromium-plated iron and steels containing chromium (stainless steels) have high corrosion resistance.

General methods for obtaining metals:

Electrometallurgy, i.e., obtaining metals by electrolysis of melts (for the most active metals) or solutions of their salts;

Pyrometallurgy, i.e., the recovery of metals from their ores at high temperature (for example, the production of iron using a blast furnace process);

Hydrometallurgy, i.e., the isolation of metals from solutions of their salts by more active metals (for example, the production of copper from a CuSO4 solution by displacement of zinc, iron

or aluminium).

In nature, metals are sometimes found in free form, such as native mercury, silver and gold, and more often in the form of compounds (metal ores). The most active metals, of course, are present in the earth's crust only in bound form.

Lithium (from the Greek. Lithos - stone), Li, a chemical element of subgroup Ia of the periodic system; atomic number 3, atomic mass 6.941; belongs to the alkali metals.

The content of lithium in the earth's crust is 6.5-10-3% by weight. It was found in more than 150 minerals, of which about 30 are actually lithium. The main minerals are spodumene LiAl, lepidolite KLi1.5 Al1.5(F.0H)2 and petalite (LiNa). The composition of these minerals is complex; many of them belong to the class of aluminosilicates, which is very common in the earth's crust. Promising sources of raw materials for the production of lithium are brines (brine) of salt-bearing deposits and groundwater. The largest deposits of lithium compounds are in Canada, the USA, Chile, Zimbabwe, Brazil, Namibia and Russia.

Interestingly, the mineral spodumene occurs in nature in the form of large crystals weighing several tons. At the Etta mine in the United States, a needle-shaped crystal 16 m long and weighing 100 tons was found.

The first information about lithium dates back to 1817. The Swedish chemist A. Arfvedson, while analyzing the mineral petalite, discovered an unknown alkali in it. Arfvedson's teacher J. Berzelius gave it the name "lithion" (from the Greek liteos - stone), because, unlike potassium and sodium hydroxides, which were obtained from plant ash, a new alkali was found in the mineral. He also named the metal, which is the "basis" of this alkali, lithium. In 1818, the English chemist and physicist G. Davy obtained lithium by electrolysis of LiOH hydroxide.

Properties. Lithium is a silvery white metal; m.p. 180.54 °C, bp 1340 "C; the lightest of all metals, its density is 0.534 g / cm - it is 5 times lighter than aluminum and almost twice as light as water. Lithium is soft and ductile. Lithium compounds color the flame in a beautiful carmine red color. This very sensitive method is used in a qualitative analysis for the detection of lithium.

The configuration of the outer electron layer of the lithium atom is 2s1 (s-element). In compounds, it exhibits an oxidation state of +1.

Lithium is the first in the electrochemical series of voltages and displaces hydrogen not only from acids, but also from water. However, many chemical reactions of lithium are less vigorous than those of other alkali metals.

Lithium practically does not react with air components in the complete absence of moisture at room temperature. When heated in air above 200 °C, Li2O oxide forms as the main product (only traces of Li2O2 peroxide are present). In moist air it gives mainly Li3N nitride, at air humidity more than 80% - LiOH hydroxide and Li2CO3 carbonate. Lithium nitride can also be obtained by heating the metal in a stream of nitrogen (lithium is one of the few elements that combine directly with nitrogen): 6Li + N2 \u003d 2Li3N

Lithium easily alloys with almost all metals and is highly soluble in mercury. It combines directly with halogens (with iodine - when heated). At 500 °C, it reacts with hydrogen to form LiH hydride, when interacting with water, LiOH hydroxide, with dilute acids, lithium salts, and with ammonia, LiNH2 amide, for example:

2Li + H2 = 2LiH

2Li + 2H2O = 2LiOH + H2

2Li + 2HF = 2LiF + H2

2Li + 2NH3 = 2LiNH2 + H2

LiH hydride - colorless crystals; used in various fields of chemistry as a reducing agent. When interacting with water, it releases a large amount of hydrogen (2820 l of H2 are obtained from 1 kg of LiH):

LiH + H2O = LiOH + H2

This makes it possible to use LiH as a source of hydrogen for filling balloons and rescue equipment (inflatable boats, belts, etc.), as well as a kind of “warehouse” for storing and transporting flammable hydrogen (in this case, it is necessary to protect LiH from the slightest traces of moisture).

Mixed lithium hydrides are widely used in organic synthesis, for example, lithium aluminum hydride LiAlH4 is a selective reducing agent. It is obtained by the interaction of LiH with aluminum chloride A1C13

LiOH hydroxide is a strong base (alkali), its aqueous solutions destroy glass, porcelain; nickel, silver and gold are resistant to it. LiOH is used as an additive to the electrolyte of alkaline batteries, which increases their service life by 2-3 times and the capacity by 20%. Based on LiOH and organic acids (especially stearic and palmitic acids), frost- and heat-resistant greases (lithols) are produced to protect metals from corrosion in the temperature range from -40 to +130 "C.

Lithium hydroxide is also used as a carbon dioxide absorber in gas masks, submarines, aircraft, and spacecraft.

Receipt and application. The raw material for the production of lithium is its salts, which are extracted from minerals. Depending on the composition, the minerals are decomposed with sulfuric acid H2SO4 (acid method) or by sintering with calcium oxide CaO and its carbonate CaCO3 (alkaline method), with potassium sulfate K2SO4 (salt method), with calcium carbonate and its CaCl chloride (alkaline-salt method) . With the acid method, a solution of sulfate Li2SO4 is obtained [the latter is freed from impurities by treatment with calcium hydroxide Ca (OH) 2 and soda Na2 Co3]. Speck formed by other methods of decomposition of minerals is leached with water; at the same time, with the alkaline method, LiOH passes into the solution, with the saline method, Li 2SO4, and with the alkaline-salt method, LiCl. All these methods, except alkaline, provide for obtaining the finished product in the form of Li2CO3 carbonate. which is used directly or as a source for the synthesis of other lithium compounds.

Lithium metal is obtained by electrolysis of a molten mixture of LiCl and potassium chloride KCl or barium chloride BaCl2 with further purification from impurities.

Interest in lithium is huge. This is primarily due to the fact that it is a source of industrial production of tritium (a heavy hydrogen nuclide), which is the main component of the hydrogen bomb and the main fuel for thermonuclear reactors. A thermonuclear reaction is carried out between the nuclide 6Li and neutrons (neutral particles with a mass number of 1); reaction products - tritium 3H and helium 4He:

63Li + 10n= 31H +42He

A large amount of lithium is used in metallurgy. An alloy of magnesium with 10% lithium is stronger and lighter than magnesium itself. Aluminum and lithium alloys - scleron and aeron, containing only 0.1% lithium, in addition to lightness, have high strength, ductility, and increased resistance to corrosion; they are used in aviation. The addition of 0.04% lithium to lead-calcium bearing alloys increases their hardness and reduces the coefficient of friction.

Lithium halides and carbonate are used in the production of optical, acid-resistant and other special glasses, as well as heat-resistant porcelain and ceramics, various glazes and enamels.

Small crumbs of lithium cause chemical burns to wet skin and eyes. Lithium salts irritate the skin. When working with lithium hydroxide, precautions must be taken, as when working with sodium and potassium hydroxides.

Sodium (from Arabic, natrun, Greek nitron - natural soda, chemical element of subgroup Ia of the periodic system; atomic number 11, atomic mass 22.98977; belongs to alkali metals. It occurs in nature in the form of one stable nuclide 23 Na.

Even in ancient times, sodium compounds were known - table salt (sodium chloride) NaCl, caustic alkali (sodium hydroxide) NaOH and soda (sodium carbonate) Na2CO3. The last substance the ancient Greeks called "nitron"; hence the modern name of the metal - "sodium". However, in the UK, USA, Italy, France, the word sodium is preserved (from the Spanish word "soda", which has the same meaning as in Russian).

For the first time, the production of sodium (and potassium) was reported by the English chemist and physicist G. Davy at a meeting of the Royal Society in London in 1807. He managed to decompose the caustic alkalis of KOH and NaOH by the action of an electric current and isolate previously unknown metals with extraordinary properties. These metals oxidized very quickly in air, and floated on the surface of the water, releasing hydrogen from it.

distribution in nature. Sodium is one of the most abundant elements in nature. Its content in the earth's crust is 2.64% by weight. In the hydrosphere, it is contained in the form of soluble salts in an amount of about 2.9% (with a total salt concentration in sea water of 3.5-3.7%). The presence of sodium has been established in the solar atmosphere and interstellar space. Sodium is naturally found only in the form of salts. The most important minerals are halite (rock salt) NaCl, mirabilite (Glauber's salt) Na2SO4 *10H2O, thenardite Na2SO4, chelian nitrate NaNO3, natural silicates, e.g. albite Na, nepheline Na

Russia is exceptionally rich in deposits of rock salt (for example, Solikamsk, Usolye-Sibirskoye, etc.), large deposits of the mineral trona in Siberia.

Properties. Sodium is a silvery-white fusible metal, m.p. 97.86 °C, bp 883.15 °C. This is one of the lightest metals - it is lighter than water with a density of 0.99 g / cm3 at 19.7 ° C). Sodium and its compounds color the burner flame yellow. This reaction is so sensitive that it reveals the presence of the slightest traces of sodium everywhere (for example, in room or street dust).

Sodium is one of the most active elements in the periodic table. The outer electron layer of the sodium atom contains one electron (configuration 3s1, sodium is an s-element). Sodium easily donates its only valence electron and therefore always exhibits an oxidation state of +1 in its compounds.

In air, sodium is actively oxidized, forming, depending on the conditions, Na2O oxide or Na2O2 peroxide. Therefore, sodium is stored under a layer of kerosene or mineral oil. Reacts vigorously with water, displacing hydrogen:

2Na + H20 = 2NaOH + H2

Such a reaction occurs even with ice at a temperature of -80 ° C, and with warm water or at the contact surface it goes with an explosion (it’s not for nothing that they say: “If you don’t want to become a freak, don’t throw sodium into the water”).

Sodium directly reacts with all non-metals: at 200 °C it begins to absorb hydrogen, forming a very hygroscopic NaH hydride; with nitrogen in an electric discharge gives nitride Na3N or azide NaN3; ignites in fluorine atmosphere; in chlorine it burns at a temperature; reacts with bromine only when heated:

2Na + H2 = 2NaH

6Na + N2=2Na3N or 2Na+ 3Na2=2NaN3

2Na+ C12 = 2NaCl

At 800-900 °C, sodium combines with carbon, forming Na2C2 carbide; when triturated with sulfur gives Na2S sulfide and a mixture of polysulfides (Na2S3 and Na2S4)

Sodium easily dissolves in liquid ammonia, the resulting blue solution has metallic conductivity, with gaseous ammonia at 300-400 "C or in the presence of a catalyst when cooled to -30 C gives amide NaNH2.

Sodium forms compounds with other metals (intermetallic compounds), for example, with silver, gold, cadmium, lead, potassium, and some others. With mercury, it gives amalgams NaHg2, NaHg4, etc. Liquid amalgams, which are formed by the gradual introduction of sodium into mercury under a layer of kerosene or mineral oil, are of the greatest importance.

Sodium forms salts with dilute acids.

Receipt and application. The main method for obtaining sodium is the electrolysis of molten common salt. In this case, chlorine is released at the anode, and sodium is released at the cathode. To reduce the melting point of the electrolyte, other salts are added to common salt: KCl, NaF, CaCl2. Electrolysis is carried out in electrolyzers with a diaphragm; anodes are made of graphite, cathodes are made of copper or iron.

Sodium can be obtained by electrolysis of a NaOH hydroxide melt, and small amounts can be obtained by decomposition of NaN3 azide.

Sodium metal is used to reduce pure metals from their compounds - potassium (from KOH), titanium (from TiCl4), etc. An alloy of sodium and potassium is a coolant for nuclear reactors, since alkali metals absorb neutrons poorly and therefore do not prevent the fission of uranium nuclei. Sodium vapor, which has a bright yellow glow, is used to fill gas discharge lamps used to illuminate highways, marinas, train stations, etc. Sodium finds application in medicine: an artificially obtained nuclide 24Na is used for radiological treatment of certain forms of leukemia and for diagnostic purposes.

The use of sodium compounds is much more extensive.

Peroxide Na2O2 - colorless crystals, yellow technical product. When heated to 311-400 °C, it begins to release oxygen, and at 540 °C it rapidly decomposes. A strong oxidizing agent, due to which it is used to bleach fabrics and other materials. It absorbs CO2 in air, releasing oxygen and forming carbonate 2Na2O2+2CO2=2Na2Co3+O2). This property is the basis for the use of Na2O2 for air regeneration in enclosed spaces and insulating breathing devices (submarines, insulating gas masks, etc.).

NaOH hydroxide; the outdated name is caustic soda, the technical name is caustic soda (from Latin caustic - caustic, burning); one of the strongest bases. The technical product, in addition to NaOH, contains impurities (up to 3% Na2CO3 and up to 1.5% NaCl). A large amount of NaOH is used for the preparation of electrolytes for alkaline batteries, the production of paper, soap, paints, cellulose, and is used to purify oil and oils.

From sodium salts, chromate Na2CrO4 is used - in the production of dyes, as a mordant in dyeing fabrics and a tanning agent in the leather industry; sulfite Na2SO3 - a component of fixers and developers in photography; hydrosulfite NaHSO3 - bleach of fabrics, natural fibers, used for canning fruits, vegetables and vegetable feed; thiosulfate Na2S2O3 - to remove chlorine when bleaching fabrics, as a fixative in photography, an antidote for poisoning with mercury compounds, arsenic, etc., an anti-inflammatory agent; chlorate NaClO3 - oxidizing agent in various pyrotechnic compositions; triphosphate Na5P3O10 - additive in synthetic detergents for water softening.

Sodium, NaOH and its solutions cause severe burns of the skin and mucous membranes.

In appearance and properties, potassium is similar to sodium, but more reactive. Reacts vigorously with water and ignites hydrogen. It burns in air, forming an orange superoxide CO2. At room temperature, it reacts with halogens, with moderate heating - with hydrogen, sulfur. In moist air, it quickly becomes covered with a layer of KOH. Potassium is stored under a layer of gasoline or kerosene.

Potassium compounds - KOH hydroxide, KNO3 nitrate and K2CO3 carbonate - find the greatest practical application.

Potassium hydroxide KOH (technical name - caustic potash) - white crystals that spread in humid air and absorb carbon dioxide (K2CO3 and KHCO3 are formed). It dissolves very well in water with a high exo effect. The aqueous solution is strongly alkaline.

Potassium hydroxide is produced by electrolysis of a KCl solution (similar to the production of NaOH). The initial potassium chloride KCl is obtained from natural raw materials (minerals sylvin KCl and carnallite KMgC13 6H20). KOH is used for the synthesis of various potassium salts, liquid soap, dyes, as an electrolyte in batteries.

Potassium nitrate KNO3 (potassium nitrate mineral) - white crystals, very bitter in taste, low melting point (tmelt = 339 ° C). Let's well dissolve in water (hydrolysis is absent). When heated above the melting point, it decomposes into potassium nitrite KNO2 and oxygen O2, and exhibits strong oxidizing properties. Sulfur and charcoal ignite upon contact with the KNO3 melt, and the C + S mixture explodes (combustion of "black powder"):

2КNO3 + ЗС(coal) + S=N2 + 3CO2 + K2S

Potassium nitrate is used in the production of glass and mineral fertilizers.

Potassium carbonate K2CO3 (technical name - potash) is a white hygroscopic powder. It is very soluble in water, highly hydrolyzed by the anion and creates an alkaline environment in the solution. Used in the manufacture of glass and soap.

Obtaining K2CO3 is based on the reactions:

K2SO4 + Ca(OH)2 + 2CO = 2K(HCOO) + CaSO4

2K(HCOO) + O2 = K2C03 + H20 + CO2

Potassium sulfate from natural raw materials (minerals kainite KMg (SO4) Cl ZH20 and schenite K2Mg (SO4) 2 * 6H20) is heated with slaked lime Ca (OH) 2 in a CO atmosphere (under a pressure of 15 atm), potassium formate K (HCOO) is obtained , which is calcined in a stream of air.

Potassium is a vital element for plants and animals. Potash fertilizers are potassium salts, both natural and their processed products (KCl, K2SO4, KNO3); high content of potassium salts in the ashes of plants.

Potassium is the ninth most abundant element in the earth's crust. It is found only in bound form in minerals, sea water (up to 0.38 g of K + ions in 1 liter), plants and living organisms (inside cells). The human body has = 175 g of potassium, the daily requirement reaches ~ 4g. The radioactive isotope 40K (an admixture to the predominant stable isotope 39K) decays very slowly (half-life is 1 109 years), it, along with the isotopes 238U and 232Th, makes a large contribution to

Metals in the periodic system. The structure of metal atoms. General characteristics of metals.

Na(+11): 1S 2 2S 2 2p 6 3S 1

Ca(+20): 1S 2 2S 2 2p 6 3S 2 3p 6 3d 0 4S 2

Alkali metals (Group 1, main subgroup): ... nS 1. Alkaline earth (Group 2, main subgroup): ... nS 2. The properties of metal atoms are in periodic dependence on their location in the table D.I. Mendeleev. IN MAIN SUBGROUP:

does not change.

Atom radius increases

Electronegativity decreases.

Restorative properties intensify.

Metal properties intensify.

Charges of the nuclei of atoms increase.

Radii of atoms decrease.

Number of electrons on the outer layer increases.

Electronegativity increases.

Restorative properties decrease.

Metal properties weaken.

a - copper; b) magnesium; c) α-modification of iron

Metal atoms tend to donate their outer electrons. In a piece of metal, ingot or metal product, metal atoms donate external electrons and send them to this piece, ingot or product, turning into ions. The “torn off” electrons move from one ion to another, temporarily reconnect with them into atoms, break off again, and this process occurs continuously. Metals have a crystal lattice, in the nodes of which there are atoms or ions (+); between them are free electrons (electron gas). The connection scheme in metal can be displayed as follows:

M 0 ↔ nē + M n+,

atom - ion

Where n is the number of external electrons participating in the bond (y Na - 1, y Sa - 2 ē, y Al - 3 ē).This type of bond is observed in metals - simple substances-metals and in alloys. A metallic bond is a bond between positively charged metal ions and free electrons in the crystal lattice of metals. communication is based on the socialization of electrons (similarity), all atoms take part in the socialization of these electrons (difference). That is why crystals with a metallic bond are plastic, electrically conductive and have a metallic sheen. However, in the vapor state, metal atoms are bound together covalent bond, metal pairs consist of individual molecules (monatomic and diatomic). general characteristics metals

Electrochemical series of voltages of metals. Physical and chemical properties of metals

α-tin, or gray tin ("tin plague" - the transformation of ordinary β-tin into α-tin during low temperatures caused the death of R. Scott's expedition to South Pole, which has lost all the fuel, since it was stored in tanks sealed with tin), is stable at t<14°С, серый порошок. β-олово, или белое олово (t = 14 ― 161°С) очень мягкий металл, но тверже свинца, поддается литью и пайке. Используется в сплавах, например, для изготовления белой жести (луженого железа).

With halogens, metals form salts - halides:

With oxygen, metals form oxides:

Metals form salts with sulfur - sulfides:

With hydrogen, the most active metals form hydrides, for example:

many metals form carbides with carbon:

Metals at the beginning of a series of voltages (from lithium to sodium), under normal conditions, displace hydrogen from water and form alkalis, for example:

Metals located in a series of voltages up to hydrogen interact with dilute acids (HCl, H 2 SO 4, etc.), as a result of which salts are formed and hydrogen is released, for example:

Metals interact with solutions of salts of less active metals, as a result of which a salt of a more active metal is formed, and a less active metal is released in a free form, for example:

Metals in nature.

coal:

carbon monoxide:

hydrogen:

aluminum (aluminothermy):

magnesium:

2CuS + 3O 2 \u003d 2CuO + 2SO 2.

CuSO 4 + Fe = FeSO 4 + Cu.

General characteristics of metals of the IA group.

The metals of the main subgroup of the first group (IA-groups) include lithium (Li), sodium (Na), potassium (K), rubidium (Rb), cesium (Cs), francium (Fr). These metals are called alkali metals, since they and their oxides form alkalis when interacting with water. Alkali metals are s-elements. Metal atoms have one s-electron (ns 1) on the outer electron layer. Potassium, sodium - simple substances

Alkali metals in ampoules:
a - cesium; b - rubidium; c - potassium; g - sodium Basic information about the elements of group IA

Lesson objectives:

educational - consider the position of metals in the system of elements D.I. Mendeleev, to introduce students to the basic properties of metals, to find out what causes them, to introduce the concept of corrosion of metals

Educational - be able to find metals in the PSCE table, be able to compare metals and non-metals, explain the reasons for the chemical and physical properties of metals, develop students' theoretical thinking and their ability to predict the properties of metals based on their structure.

nurturing - to promote the development of students' cognitive interest in the study of chemistry

Lesson type: lesson learning new material.

Teaching methods : verbal and visual

During the classes:

Lesson timing.

Organizational moment (1 min.)

Knowledge update(3 min)

Learning new material

1.1. Position in the periodic system. (10 min)

1.2. Features of the electronic structure of atoms. (10 min)

1.3. Restorative properties of metals. (10 min)

2.1. Metal connection. (5 minutes)

4. Emotional release 2 min

2.2. Physical properties. (10 min)

3. Chemical properties. (17 min)

4. Corrosion of metals. (5 min)

Fixing (15 min)

Homework (3 min)

Summary of the lesson (1 min)

Organizing time

(Mutual greeting, fixation of those present).

Knowledge update. At the beginning of the lesson, the teacher focuses the students' attention on the significance of the new topic, determined by the role that metals play in nature and in all spheres of human activity.. Industry

The teacher reads the riddle:

I am hard, malleable and plastic,

Brilliant, everyone needs, practical.

I already gave you a hint

So who am I...? and offers to write down the answer in a notebook in the form of a lesson topic?

Learning new material

Lecture plan.

1. Characteristics of the metal element.

1.2. Features of the electronic structure of atoms.

1.3. Restorative properties of metals.

2. Characteristics of a simple substance.

2.1. Metal connection.

2.2. physical properties.

3. Chemical properties.

4. Corrosion of metals.

1.1. Position in the periodic system.

The conditional boundary between metal elements and non-metal elements runs along the diagonal B (boron) - (silicon) - Si (arsenic) - Te (tellurium) - As (astatine) (track it in the table of D. I. Mendeleev) ..

The initial elements formthe main subgroup of group I and are called alkali metals . They got their name from the name of their corresponding hydroxides, which are highly soluble in water - alkalis.

Of the elements of the main subgroups of the following groups, metals include: in group IV germanium, tin, lead (32.50.82) (the first two elements are carbon and silicon - non-metals), in group V antimony and bismuth (51.83) (the first three elements are non-metals), in group VI only the last element - polonium (84) - is a pronounced metal. In the main subgroups of groups VII and VIII, all elements are typical non-metals.

As for the elements of the secondary subgroups, they are all metals.

Alkali metal atoms contain only one electron at the external energy level, which they easily donate during chemical interactions, therefore they are the strongest reducing agents. It is clear that, in accordance with an increase in the radius of the atom, the reducing properties of alkali metals increase from lithium to francium.

Following the alkali metals, the elements that make upthe main subgroup of group II, are also typical metals with a strong reducing ability (their atoms contain two electrons at the outer level).Of these metals, calcium, strontium, barium and radium are called alkaline earth metals. . These metals got this name because their oxides, which the alchemists called "earths", form alkalis when dissolved in water.

Metals also include elementsthe main subgroup of group III, excluding boron.

Group 3 includes metals called the aluminum subgroup.

1.2 Features of the electronic structure of metals.

Based on the knowledge gained, students themselves formulate the definition of "metal"

Metals are chemical elements whose atoms donate electrons from the outer (and sometimes pre-outer) electron layer, turning into positive ions. Metals are reducing agents. This is due to the small number of electrons in the outer layer, the large radius of the atoms, as a result of which these electrons are weakly retained with the nucleus.Metal atoms have relatively large sizes (radii), therefore their outer electrons are also significantly removed from the nucleus and are weakly bound to it. And the second feature that is inherent in the atoms of the most active metals isthe presence of 1-3 electrons in the external energy level.
Metal atoms have similarities in the structure of the outer electron layer, which is formed by a small number of electrons (mostly no more than three).
This statement can be illustrated by the examples of Na, aluminum Al and zinc Zn. Drawing up diagrams of the structure of atoms, if desired, you can draw up electronic formulas and give examples of the structure of elements of large periods, such as zinc.

Due to the fact that the electrons of the outer layer of metal atoms are weakly bound to the nucleus, they can be “given away” to other particles, which happens during chemical reactions:

The property of metal atoms to donate electrons is their characteristic chemical property and indicates that metals exhibit reducing properties.

1.3 Reducing properties of metals.

How the oxidizing power of elements changesIIIperiod?

(Oxidative properties increase in periods, and reduction properties weaken. The reason for the change in these properties is an increase in the number of electrons in the last orbital.)

How do the oxidizing properties of the elements of the 4th group of the main subgroup change?(from bottom to top, oxidizing properties are enhanced. The reason for the change in these properties is a decrease in the radius of the atom (it is easier to accept than to give away)

Based on the position of metals in the Periodic system, what conclusion can be drawn about the redox properties of metal elements?

(Metals are reducing agents in chemical reactions, because they donate their valence electrons)

Students answer that the strength of the bond between valence electrons and the nucleus depends on two factors:the charge of the nucleus and the radius of the atom. .

(recording the conclusion in students' notebooks) in periods with an increase in the charge of the nucleus, the restorative properties decrease.

For elements - metals of secondary subgroups, the properties are slightly different.

The teacher offers to compare the activity of the elements of the secondary subgroup.Cu, Ag, Au – activityb elements - metals drops. This pattern is also observed in the elements of the second secondary subgroupZn, CD, hg.The increase in electrons at the outer level, so the reducing properties are weakened

For elements of secondary subgroups - these are elements of 4-7 periods 31-36, 49-54 - with an increase in the ordinal element, the radius of the atoms will change little, and the value of the charge of the nucleus increases significantly, therefore, the strength of the bond of valence electrons with the nucleus increases, reducing properties weaken.

2.1. Metal connection.

The metallic bond is carried out through the mutual attraction of atom-ions and relatively free electrons.

Picture 1.
The structure of the crystal lattice of metals

In metals, valence electrons are held by atoms extremely weakly and are able to migrate. Atoms left without external electrons acquire a positive charge. They form a metallic crystal lattice.

A set of socialized valence electrons (electron gas), negatively charged, holds positive metal ions at certain points in space - the nodes of the crystal lattice, for example, silver metal.

External electrons can move freely and randomly, therefore metals are characterized by high electrical conductivity (especially gold, silver, copper, aluminum).

Chemical bond involves a certain type of crystal lattice. The metallic chemical bond promotes the formation of crystals with a metallic crystal lattice. At the nodes of the crystal lattice are atom-ions of metals, and between them are freely moving electrons. The metallic bond differs from the ionic one, because no anions, although there are cations. It also differs from the covalent one, because no shared electron pairs are formed.

Emotional discharge

The absence of what metal was described by Academician A.E. Fersman?

There would be a horror of destruction on the streets: there would be no rails, no wagons, no steam locomotives, no cars, even pavement stones would turn into clay dust, and plants would begin to wither and die without this metal. Destruction by a hurricane would have passed over the entire Earth, and the death of mankind would have become inevitable. However, a person would not have lived up to this moment, because having lost three grams of this metal in his body and in his blood, he would have ceased to exist before the drawn events would have unfolded (Answer: All people would have died, having lost iron in their blood)

Name the counterfeiter's metal

The name of the metal was given by the Spanish conquistadors, who in the middle of the 16th century. first met in South America (on the territory of modern Colombia) with a new metal that looks like silver. The name of the metal literally means "little silver", "silver".

Such a dismissive name is explained by the exceptional refractoriness of the metal, which was not amenable to remelting, did not find application for a long time and was valued half as much as silver. They used this metal to make counterfeit coins.

Today, this metal, used as a catalyst and in jewelry, is one of the most expensive.

It does not exist in its pure form in nature. Native platinum is usually a natural alloy with other noble (palladium, iridium, rhodium, ruthenium, osmium) and base (iron, copper, nickel, lead, silicon) metals. To obtain it, nuggets are heated in boilers with "aqua regia" (a mixture of nitric and hydrochloric acid) and then "finished" by numerous chemical reactions, heating and melting.

Thus, the crystal lattice depends and is determined by the type of chemical bond, but at the same time is the cause for physical properties.

2.2. physical properties.

The teacher emphasizes that the physical properties of metals are determined precisely by their structure.

A)hardness All metals except mercury are solids under normal conditions. The mildest are sodium, potassium. They can be cut with a knife; hardest chrome - scratches glass

b)density. Metals are divided into soft (5g/cm³) and heavy (less than 5g/cm³).

V)fusibility. Metals are divided into fusible and refractory.

G)electrical conductivity, thermal conductivity metals is due to their structure. Chaotically moving electrons under the influence of an electric voltage acquire a directed movement, resulting in an electric current.

With an increase in temperature, the amplitude of the movement of atoms and ions located in the nodes of the crystal lattice increases sharply, and this interferes with the movement of electrons, and the electrical conductivity of metals decreases.

It should be noted that in some non-metals, with increasing temperature, the electrical conductivity increases, for example, in graphite, while with increasing temperature, some covalent bonds are destroyed, and the number of freely moving electrons increases.

e)metallic luster - Electrons filling the interatomic space reflect light rays, and do not transmit like glass. They fall on the nodes of the crystal lattice. Therefore, all metals in the crystalline state have a metallic luster. For most metals, all the rays of the visible part of the spectrum are equally scattered, so they have a silvery-white color. Only gold and copper absorb short wavelengths to a large extent and reflect long wavelengths of the light spectrum, so they have a yellow color. The most brilliant metals are mercury, silver, palladium. All metals in powder exceptAlAndmg, lose their luster and are black or dark gray in color.

e)plastic

The mechanical effect on a crystal with a metal lattice causes only a displacement of the layers of atoms and is not accompanied by bond breaking, and therefore the metal is characterized by high plasticity.

3. Chemical properties.

According to their chemical properties, all metals are reducing agents, they all give up valence electrons relatively easily, pass into positively charged ions, that is, they are oxidized . The reducing activity of a metal in chemical reactions occurring in aqueous solutions reflects its position in the electrochemical series of metal voltages (Discovered and compiled by Beketov)

The further to the left a metal is in the electrochemical series of metal voltages, the more powerful the reducing agent is, the strongest reducing agent is lithium metal, gold is the weakest, and, conversely, the gold (III) ion is the strongest oxidizing agent, lithium (I) is the most weak.

Each metal is able to restore from salts in solution those metals that are in a series of voltages after it, for example, iron can displace copper from solutions of its salts. However, it should be remembered that alkali and alkaline earth metals will interact directly with water.

Metals, standing in the series of voltages to the left of hydrogen, are able to displace it from solutions of dilute acids, while dissolving in them.

The reducing activity of a metal does not always correspond to its position in the periodic system, because when determining the place of a metal in a series, not only its ability to donate electrons is taken into account, but also the energy expended on the destruction of the metal crystal lattice, as well as the energy expended on the hydration of ions.

Interaction with simple substances

WITHoxygen most metals form oxides - amphoteric and basic:

4Li+O 2 = 2Li 2 O

4Al + 3O 2 = 2Al 2 O 3 .

Alkali metals, with the exception of lithium, form peroxides:

2Na+O 2 = Na 2 O 2 .

WITHhalogens metals form salts of hydrohalic acids, for example,

Cu + Cl 2 = CuCl 2 .

WITHhydrogen the most active metals form ionic hydrides - salt-like substances in which hydrogen has an oxidation state of -1.

2Na+H 2 = 2NaH.

WITHgray metals form sulfides - salts of hydrosulfide acid:

Zn + S = ZnS.

WITHnitrogen some metals form nitrides, the reaction almost always proceeds when heated:

3Mg+N 2 =Mg 3 N 2 .

WITHcarbon carbides are formed.

4Al + 3C = Al 3 C 4 .

WITHphosphorus - phosphides:

3Ca + 2P = Ca 3 P 2 .

Metals can interact with each other to formintermetallic compounds :

2Na + Sb = Na 2 sb,

3Cu + Au = Cu 3 Au.

Metals can dissolve in each other at high temperature without interaction, forming alloys.

The ratio of metals to acids.

Most often in chemical practice such strong acids as sulfuric H 2 SO 4 , hydrochloric HCl and nitric HNO 3 .

WithHCl

The hydrogen ions H formed in this process + act as an oxidizing agentmetals in the activity series to the left of hydrogen . The interaction proceeds according to the scheme:

Me + HCl - salt + H 2

2 Al + 6 HCl → 2 AlCl 3 + 3 H 2

2│Al 0 – 3 e - → Al 3+ - oxidation

3│2H + + 2 e - → H 2 - recovery

"Aqua regia" (formerly called vodka acids) is a mixture of one volume of nitric acid and three to four volumes of concentrated hydrochloric acid, which has a very high oxidative activity. Such a mixture is capable of dissolving some low-active metals that do not interact with nitric acid. Among them is the "king of metals" - gold. This effect of "aqua regia" is explained by the fact that nitric acid oxidizes hydrochloric acid with the release of free chlorine and the formation of nitrogen (III) chlorine oxide, or nitrosyl chloride - NOCl:

Gold oxidation reactions proceed according to the following equations:

Au + HNO3 + 4 HCl → H + NO + 2H2O

If acids can interact with bases and basic oxides, and the key element in their composition is a metal, then is it possible for metals to interact with acids. Let's check it experimentally.

Magnesium interacts with acid under normal conditions, zinc - when heated, copper - does not interact.

A range of voltages are used in practice for a comparative assessment of the chemical activity of metals in reactions with aqueous solutions of salts and acids and for the assessment of cathodic and anodic processes during electrolysis:

Metals to the left are stronger reducing agents, than the metals to the right:they displace the latter from salt solutions . Metals in the row to the left of hydrogen displace hydrogen when interacting with aqueous solutions of non-oxidizing acids; the most active metals (up to and including aluminum) - and when interacting with water.

Metals in the row to the right of hydrogen do not interact with aqueous solutions of non-oxidizing acids under normal conditions.

During electrolysis, metals to the right of hydrogen are released at the cathode; the reduction of metals of moderate activity is accompanied by the release of hydrogen; the most active metals (up to aluminum) cannot be isolated from aqueous solutions of salts under normal conditions.

4. Corrosion of metals – physical-chemical or chemical interaction between a metal (alloy) and a medium, leading to a deterioration in the functional properties of the metal (alloy), the medium or the technical system that includes them.

The word corrosion comes from the Latin "corrodo" - "gnaw" (Late Latin "corrosio" means "corrosion").

Corrosion is caused by the chemical reaction of a metal with environmental substances occurring at the interface between the metal and the medium. Most often, this is the oxidation of a metal, for example, with atmospheric oxygen or acids contained in solutions with which the metal comes into contact. Metals located in the voltage series (activity series) to the left of hydrogen, including iron, are especially susceptible to this.

As a result of corrosion, iron rusts. This process is very complex and includes several stages. It can be described by the overall equation:

4Fe + 6H 2 O (moisture) + 3O 2 (air) = 4Fe(OH) 3

Iron(III) hydroxide is very unstable, quickly loses water and turns into iron(III) oxide. This compound does not protect the iron surface from further oxidation. As a result, the iron object can be completely destroyed.

To slow down corrosion, varnishes and paints, mineral oils and grease are applied to the metal surface. Underground structures are covered with a thick layer of bitumen or polyethylene. The interior surfaces of steel pipes and tanks are protected with cheap cement coatings.

For steel products, so-called rust converters containing phosphoric acid (H 3 RO 4 ) and its salts. They dissolve the remains of oxides and form a dense and durable film of phosphates, which is able to protect the surface of the product for some time. Then the metal is coated with a primer layer, which should fit well on the surface and have protective properties (usually red lead or zinc chromate is used). Only then can varnish or paint be applied.

Fixing (15 min)

Teacher:

Now, to fix it, let's do a test.

Solve test tasks

1.Select a group of elements that contains only metals:

A) Al, As, P;B) Mg, Ca, Si;IN) K, Ca, Pb

2. Select a group in which there are only simple substances - non-metals:

A) K 2 Oh, SO 2 , SiO 2 ; B) H 2 , Cl 2 , I 2 ; IN)Ca, Ba, HCl;

3. Indicate what is common in the structure of K and Li atoms:

A) 2 electrons on the last electron layer;

B) 1 electron on the last electron layer;

C) the same number of electronic layers.

4. Metal calcium exhibits properties:

A) an oxidizing agent

B) reducing agent;

C) an oxidizing or reducing agent, depending on the conditions.

5. The metallic properties of sodium are weaker than those of -

A) magnesium; B) potassium; B) lithium.

6. Inactive metals include:

A) aluminum, copper, zinc; B) mercury, silver, copper;

C) calcium, beryllium, silver.

7. What physical property is not common to all metals:

A) electrical conductivity, B) thermal conductivity,

C) solid state of aggregation under normal conditions,

D) metallic luster

8. Metals, when interacting with non-metals, exhibit the following properties:

a) oxidizing;

b) recovery;

c) both oxidizing and reducing;

d) do not participate in redox reactions.

9. In the periodic system, typical metals are located

a) the top

b) bottom

in the upper right corner

d) lower left corner

Part B. The answer to the tasks of this part is a set of letters that should be written down

Set a match.

With an increase in the ordinal number of an element in the main subgroup of group II of the Periodic system, the properties of the elements and the substances they form change as follows:

1) the number of electrons in the outer level

A) increases

3) electronegativity

4) restorative properties

B) decreases

B) does not change

(Answers: 1 -D, 2 -A, 3 -C, 4-B, 5-D)

TASKS FOR REINFORCEMENT

№1. Finish the equations of practically feasible reactions, name the reaction products

Li + H 2 O=

Cu + H 2 O=Cu( Oh) 2 + H 2

Ba+H 2 O=

Mg + H 2 O=

Ca+HCl=

2 Na+2 H 2 SO 4 ( TO)= Na 2 SO 4 + SO 2 + 2H 2 O

HCl + Zn =

H 2 SO 4 ( To)+ Cu=CuSO 4 + SO 2 + H 2 O

H 2 S + Mg \u003d MgS + H 2

HCl + Cu =

Homework: notes in notebooks, reports on the use of metals.

Teacher Offers to create a syncwine on the topic.

Line 1: Noun (one per topic) (Metals)

2nd line: two adjectives

3rd line: three verbs

4 line: four words combined into a sentence

Line 5: a word expressing the essence of this topic.

Lesson summary

Teacher : And so, we examined the structure and physical properties of metals, their position in the periodic system of chemical elements D.I. Mendeleev.

B O Most of the known chemical elements form simple substances, metals.

Metals include all elements of secondary (B) subgroups, as well as elements of the main subgroups located below the diagonal "beryllium - astatine" (Fig. 1). In addition, the chemical elements metals form groups of lanthanides and actinides.

Rice. 1. The location of metals among the elements of subgroups A (highlighted in blue)

Compared to non-metal atoms, metal atoms have b O Larger sizes and fewer outer electrons, usually 1-2. Consequently, the outer electrons of metal atoms are weakly bound to the nucleus; metals easily give them away, exhibiting reducing properties in chemical reactions.

Consider the patterns of change in some properties of metals in groups and periods.

In periodsWith As the nuclear charge increases, the atomic radius decreases. The nuclei of atoms attract outer electrons more and more, therefore, the electronegativity of atoms increases, the metallic properties decrease. Rice. 2.

Rice. 2. Change in metallic properties in periods

In the main subgroups from top to bottom in metal atoms, the number of electron layers increases, therefore, the radius of atoms increases. Then the outer electrons will be weaker attracted to the nucleus, so there is a decrease in the electronegativity of atoms and an increase in metallic properties. Rice. 3.

Rice. 3. Change in metallic properties in subgroups

These regularities are also characteristic of the elements of secondary subgroups, with rare exceptions.

Atoms of metal elements tend to donate electrons. In chemical reactions, metals act only as reducing agents, they donate electrons and increase their oxidation state.

Electrons can be received from metal atoms by atoms that make up simple substances, non-metals, as well as atoms that are part of complex substances that are able to lower their oxidation state. For example:

2Na 0 + S 0 = Na +1 2 S -2

Zn 0 + 2H +1 Cl \u003d Zn +2 Cl 2 + H 0 2

Not all metals have the same chemical activity. Some metals under normal conditions practically do not enter into chemical reactions, they are called noble metals. The noble metals include: gold, silver, platinum, osmium, iridium, palladium, ruthenium, rhodium.

Noble metals are very rare in nature and are almost always found in the native state (Fig. 4). Despite the high resistance to corrosion-oxidation, these metals still form oxides and other chemical compounds, for example, silver chloride and nitrate salts are known to everyone.

Rice. 4. Nugget of gold

Summing up the lesson

In this lesson, you examined the position of the chemical elements of metals in the Periodic Table, as well as the structural features of the atoms of these elements, which determine the properties of simple and complex substances. You have learned why there are much more chemical elements of metals than non-metals.

Bibliography

1. Orzhekovsky P.A. Chemistry: 9th grade: textbook for general education. inst. / P.A. Orzhekovsky, L.M. Meshcheryakova, M.M. Shalashova. - M.: Astrel, 2013. (§28)
2. Rudzitis G.E. Chemistry: inorgan. chemistry. Organ. chemistry: textbook. for 9 cells. / G.E. Rudzitis, F.G. Feldman. - M.: Enlightenment, JSC "Moscow textbooks", 2009. (§34)
3. Khomchenko I.D. Collection of problems and exercises in chemistry for high school. - M.: RIA "New Wave": Publisher Umerenkov, 2008. (p. 86-87)
4. Encyclopedia for children. Volume 17. Chemistry / Chapter. ed. V.A. Volodin, leading. scientific ed. I. Leenson. - M.: Avanta +, 2003.

1. A single collection of digital educational resources (video experiences on the topic) ().
2. Electronic version of the journal "Chemistry and Life" ().

Homework

1. With. 195-196 No. 7, A1-A4 from the textbook by P.A. Orzhekovsky "Chemistry: 9th grade" / P.A. Orzhekovsky, L.M. Meshcheryakova, M.M. Shalashova. - M.: Astrel, 2013.
2. What properties (oxidizing or reducing) can an Fe 3+ ion have? Illustrate your answer with reaction equations.
3. Compare the atomic radius, electronegativity and reducing properties of sodium and magnesium.