General properties of metals. Metal connection

Interaction of metals with simple oxidizing agents. The ratio of metals to water, aqueous solutions of acids, alkalis and salts. The role of the oxide film and oxidation products. Interaction of metals with nitric and concentrated sulfuric acids.

Metals include all s-, d-, f-elements, as well as p-elements located in the lower part periodic table from the diagonal drawn from boron to astatine. In simple substances of these elements, a metallic bond is realized. Metal atoms have few electrons on the outer electron shell, in quantities of 1, 2, or 3. Metals exhibit electropositive properties and have low electronegativity, less than two.

Metals have characteristic characteristics. These are solid substances, heavier than water, with a metallic luster. Metals have high thermal and electrical conductivity. They are characterized by the emission of electrons under the influence of various external influences: irradiation with light, heating, rupture (exoelectronic emission).

The main characteristic of metals is their ability to donate electrons to atoms and ions of other substances. Metals are reducing agents in the vast majority of cases. And this is their characteristic chemical property. Let's consider the ratio of metals to typical oxidizing agents, which include simple substances - non-metals, water, acids. Table 1 provides information on the ratio of metals to simple oxidizing agents.

Table 1

Ratio of metals to simple oxidizing agents

All metals react with fluorine. The exceptions are aluminum, iron, nickel, copper, zinc in the absence of moisture. These elements, when reacting with fluorine at the initial moment, form fluoride films that protect the metals from further reaction.

Under the same conditions and reasons, iron is passivated in reaction with chlorine. In relation to oxygen, not all, but only a number of metals form dense protective films of oxides. When moving from fluorine to nitrogen (Table 1), oxidative activity decreases and therefore all larger number metals are not oxidized. For example, only lithium and alkaline earth metals react with nitrogen.

The ratio of metals to water and aqueous solutions of oxidizing agents.

IN aqueous solutions The reduction activity of a metal is characterized by the value of its standard redox potential. From the entire series of standard redox potentials, a series of metal voltages is distinguished, which is listed in Table 2.

table 2

Range of voltage metals

Oxidizer	Electrode Process Equation	Standard electrode potential φ 0, V	Reducing agent	Conditional activity of reducing agents
Li+	Li + + e - = Li	-3,045	Li	Active
Rb+	Rb + + e - = Rb	-2,925	Rb	Active
K+	K + + e - = K	-2,925	K	Active
Cs+	Cs + + e - = Cs	-2,923	Cs	Active
Ca2+	Ca 2+ + 2e - = Ca	-2,866	Ca	Active
Na+	Na + + e - = Na	-2,714	Na	Active
Mg 2+	Mg 2+ +2 e - = Mg	-2,363	Mg	Active
Al 3+	Al 3+ + 3e - = Al	-1,662	Al	Active
Ti 2+	Ti 2+ + 2e - = Ti	-1,628	Ti	Wed. activity
Mn 2+	Mn 2+ + 2e - = Mn	-1,180	Mn	Wed. activity
Cr 2+	Cr 2+ + 2e - = Cr	-0,913	Cr	Wed. activity
H2O	2H 2 O+ 2e - =H 2 +2OH -	-0,826	H 2 , pH=14	Wed. activity
Zn 2+	Zn 2+ + 2e - = Zn	-0,763	Zn	Wed. activity
Cr 3+	Cr 3+ +3e - = Cr	-0,744	Cr	Wed. activity
Fe 2+	Fe 2+ + e - = Fe	-0,440	Fe	Wed. activity
H2O	2H 2 O + e - = H 2 +2OH -	-0,413	H 2 , pH=7	Wed. activity
Cd 2+	Cd 2+ + 2e - = Cd	-0,403	Cd	Wed. activity
Co2+	Co 2+ +2 e - = Co	-0,227	Co	Wed. activity
Ni 2+	Ni 2+ + 2e - = Ni	-0,225	Ni	Wed. activity
Sn 2+	Sn 2+ + 2e - = Sn	-0,136	Sn	Wed. activity
Pb 2+	Pb 2+ + 2e - = Pb	-0,126	Pb	Wed. activity
Fe 3+	Fe 3+ +3e - = Fe	-0,036	Fe	Wed. activity
H+	2H + + 2e - =H 2		H 2 , pH=0	Wed. activity
Bi 3+	Bi 3+ + 3e - = Bi	0,215	Bi	Low active
Cu 2+	Cu 2+ + 2e - = Cu	0,337	Cu	Low active
Cu+	Cu + + e - = Cu	0,521	Cu	Low active
Hg 2 2+	Hg 2 2+ + 2e - = Hg	0,788	Hg 2	Low active
Ag+	Ag + + e - = Ag	0,799	Ag	Low active
Hg 2+	Hg 2+ +2e - = Hg	0,854	Hg	Low active
Pt 2+	Pt 2+ + 2e - = Pt	1,2	Pt	Low active
Au 3+	Au 3+ + 3e - = Au	1,498	Au	Low active
Au+	Au + + e - = Au	1,691	Au	Low active

This series of voltages also shows the values ​​of the electrode potentials of the hydrogen electrode in acidic (pH=0), neutral (pH=7), alkaline (pH=14) environments. The position of a particular metal in the stress series characterizes its ability to undergo redox interactions in aqueous solutions under standard conditions. Metal ions are oxidizing agents, and metals are reducing agents. The further a metal is located in the voltage series, the more powerful its ions are as an oxidizing agent in an aqueous solution. The closer the metal is to the beginning of the series, the stronger the reducing agent it is.

Metals are capable of displacing each other from salt solutions. The direction of the reaction is determined by their relative position in the series of stresses. It should be kept in mind that active metals displace hydrogen not only from water, but also from any aqueous solution. Therefore, the mutual displacement of metals from solutions of their salts occurs only in the case of metals located in the stress series after magnesium.

All metals are divided into three conditional groups, as reflected in the following table.

Table 3

Conventional division of metals

Interaction with water. The oxidizing agent in water is the hydrogen ion. Therefore, only those metals whose standard electrode potentials are lower than the potential of hydrogen ions in water can be oxidized by water. It depends on the pH of the environment and is equal to

φ = -0.059рН.

In a neutral environment (pH=7) φ = -0.41 V. The nature of the interaction of metals with water is presented in Table 4.

Metals from the beginning of the series, having a potential significantly more negative than -0.41 V, displace hydrogen from water. But magnesium displaces hydrogen only from hot water. Typically, metals located between magnesium and lead do not displace hydrogen from water. Oxide films are formed on the surface of these metals, which have a protective effect.

Table 4

Interaction of metals with water in a neutral environment

Interaction of metals with hydrochloric acid.

The oxidizing agent in hydrochloric acid is the hydrogen ion. The standard electrode potential of a hydrogen ion is zero. Therefore, all active and intermediate active metals must react with the acid. Passivation occurs only for lead.

Table 5

Interaction of metals with hydrochloric acid

Copper can be dissolved in very concentrated hydrochloric acid, despite the fact that it is a low-active metal.

The interaction of metals with sulfuric acid occurs differently and depends on its concentration.

Interaction of metals with dilute sulfuric acid. The interaction with dilute sulfuric acid is carried out in the same way as with hydrochloric acid.

Table 6

Reaction of metals with dilute sulfuric acid

Dilute sulfuric acid oxidizes with its hydrogen ion. It interacts with those metals whose electrode potentials are lower than those of hydrogen. Lead does not dissolve in sulfuric acid at a concentration below 80%, since the PbSO 4 salt formed during the interaction of lead with sulfuric acid is insoluble and creates a protective film on the metal surface.

Interaction of metals with concentrated sulfuric acid.

In concentrated sulfuric acid, sulfur in the oxidation state +6 acts as an oxidizing agent. It is part of the sulfate ion SO 4 2-. Therefore, concentrated acid oxidizes all metals whose standard electrode potential is less than that of the oxidizing agent. Highest value the electrode potential in electrode processes involving the sulfate ion as an oxidizing agent is 0.36 V. As a result, some low-active metals also react with concentrated sulfuric acid.

For metals of medium activity (Al, Fe), passivation occurs due to the formation of dense oxide films. Tin is oxidized to the tetravalent state to form tin(IV) sulfate:

Sn + 4 H 2 SO 4 (conc.) = Sn(SO 4) 2 + 2SO 2 + 2H 2 O.

Table 7

Reaction of metals with concentrated sulfuric acid

Lead is oxidized to the divalent state to form soluble lead hydrogen sulfate. Mercury dissolves in hot concentrated sulfuric acid to form mercury(I) and mercury(II) sulfates. Even silver dissolves in boiling concentrated sulfuric acid.

It should be borne in mind that the more active the metal, the deeper the degree of reduction of sulfuric acid. With active metals, the acid is reduced mainly to hydrogen sulfide, although other products are also present. For example

Zn + 2H 2 SO 4 = ZnSO 4 + SO 2 + 2H 2 O;

3Zn + 4H 2 SO 4 = 3ZnSO 4 + S↓ +4H 2 O;

4Zn +5H 2 SO 4 = 4ZnSO 4 = 4ZnSO 4 +H 2 S +4H 2 O.

Interaction of metals with dilute nitric acid.

In nitric acid, nitrogen acts as an oxidizing agent in the oxidation state +5. The maximum value of the electrode potential for the nitrate ion of a dilute acid as an oxidizing agent is 0.96 V. As a result of this of great importance, nitric acid is a stronger oxidizing agent than sulfuric acid. This can be seen from the fact that nitric acid oxidizes silver. The more active the metal and the more dilute the acid, the more deeply the acid is reduced.

Table 8

Reaction of metals with dilute nitric acid

Interaction of metals with concentrated nitric acid.

Concentrated nitric acid is usually reduced to nitrogen dioxide. The interaction of concentrated nitric acid with metals is presented in Table 9.

When using acid in deficiency and without stirring, active metals reduce it to nitrogen, and metals of medium activity to carbon monoxide.

Table 9

Reaction of concentrated nitric acid with metals

Interaction of metals with alkali solutions.

Metals cannot be oxidized by alkalis. This is due to the fact that alkali metals are strong reducing agents. Therefore, their ions are the weakest oxidizing agents and do not exhibit oxidizing properties in aqueous solutions. However, in the presence of alkalis, the oxidizing effect of water is manifested to a greater extent than in their absence. Due to this, in alkaline solutions, metals are oxidized by water to form hydroxides and hydrogen. If the oxide and hydroxide are amphoteric compounds, then they will dissolve in an alkaline solution. As a result, metals that are passive in pure water vigorously interact with alkali solutions.

Table 10

Interaction of metals with alkali solutions

The dissolution process is represented in two stages: metal oxidation with water and hydroxide dissolution:

Zn + 2HOH = Zn(OH) 2 ↓ + H 2 ;

Zn(OH) 2 ↓ + 2NaOH = Na 2.

There are technological, physical, mechanical and chemical properties of metals. Physical properties include color and electrical conductivity. The characteristics of this group also include thermal conductivity, fusibility and density of the metal.

Mechanical characteristics include plasticity, elasticity, hardness, strength, and toughness.

Chemical properties of metals include corrosion resistance, solubility, and oxidability.

Characteristics such as fluidity, hardenability, weldability, and malleability are technological.

Physical properties

1. Color. Metals do not transmit light through themselves, that is, they are opaque. In reflected light, each element has its own shade - color. Among technical metals, only copper and its alloys have color. The remaining elements are characterized by a shade ranging from silver-white to steel-gray.
2. Fusibility. This characteristic indicates the ability of an element to transform into a liquid state from a solid state under the influence of temperature. Fusibility is considered the most important property of metals. During the heating process, all metals change from a solid state to a liquid state. When the molten substance is cooled, a reverse transition occurs - from the liquid to the solid state.
3. Electrical conductivity. This characteristic indicates the ability of free electrons to transfer electricity. The electrical conductivity of metallic bodies is thousands of times greater than that of non-metallic bodies. As the temperature increases, the conductivity of electricity decreases, and as the temperature decreases, it increases accordingly. It should be noted that the electrical conductivity of alloys will always be lower than that of any metal that makes up the alloy.
4. Magnetic properties. Obviously magnetic (ferromagnetic) elements include only cobalt, nickel, iron, as well as a number of their alloys. However, when heated to a certain temperature, these substances lose their magnetism. Certain iron alloys at room temperature are not ferromagnetic.
5. Thermal conductivity. This characteristic indicates the ability of heat to transfer to a less heated body from a more heated body without visible movement of its constituent particles. High level thermal conductivity allows metals to be heated and cooled evenly and quickly. Among technical elements, copper has the highest indicator.

Metals occupy a special place in chemistry. The presence of appropriate characteristics allows the use of a particular substance in a certain area.

Chemical properties of metals

1. Corrosion resistance. Corrosion is the destruction of a substance as a result of electrochemical or chemical interaction with environment. The most common example is the rusting of iron. Corrosion resistance is one of the most important natural characteristics of a number of metals. In this regard, substances such as silver, gold, and platinum are called noble. Nickel has high corrosion resistance and other non-ferrous materials are subject to destruction faster and more severely than non-ferrous ones.
2. Oxidability. This characteristic indicates the ability of the element to react with O2 under the influence of oxidizing agents.
3. Solubility. Metals that have liquid state unlimited solubility; upon hardening, they can form solid solutions. In these solutions, atoms from one component are incorporated into another component only within certain limits.

It should be noted that the physical and chemical properties of metals are one of the main characteristics of these elements.

Metals vary greatly in their chemical activity. The chemical activity of a metal can be approximately judged by its position in.

The most active metals are located at the beginning of this row (on the left), the least active are at the end (on the right).
Reactions with simple substances. Metals react with nonmetals to form binary compounds. The reaction conditions, and sometimes their products, vary greatly for different metals.
For example, alkali metals actively react with oxygen (including in air) at room temperature to form oxides and peroxides

4Li + O 2 = 2Li 2 O;
2Na + O 2 = Na 2 O 2

Medium activity metals react with oxygen when heated. In this case, oxides are formed:

2Mg + O 2 = t 2MgO.

Low-active metals (for example, gold, platinum) do not react with oxygen and therefore practically do not change their luster in air.
Most metals, when heated with sulfur powder, form the corresponding sulfides:

Reactions with complex substances. Compounds of all classes react with metals - oxides (including water), acids, bases and salts.
Active metals react violently with water at room temperature:

2Li + 2H 2 O = 2LiOH + H 2;
Ba + 2H 2 O = Ba(OH) 2 + H 2.

The surface of metals such as magnesium and aluminum is protected by a dense film of the corresponding oxide. This prevents the reaction from occurring with water. However, if this film is removed or its integrity is disrupted, then these metals also actively react. For example, powdered magnesium reacts with hot water:

Mg + 2H 2 O = 100 °C Mg(OH) 2 + H 2.

At elevated temperatures, less active metals also react with water: Zn, Fe, Mil, etc. In this case, the corresponding oxides are formed. For example, when passing water vapor over hot iron filings, the following reaction occurs:

3Fe + 4H 2 O = t Fe 3 O 4 + 4H 2.

Metals in the activity series up to hydrogen react with acids (except HNO 3) to form salts and hydrogen. Active metals (K, Na, Ca, Mg) react with acid solutions very violently (at high speed):

Ca + 2HCl = CaCl 2 + H 2;
2Al + 3H 2 SO 4 = Al 2 (SO 4) 3 + 3H 2.

Low-active metals are often practically insoluble in acids. This is due to the formation of a film of insoluble salt on their surface. For example, lead, which is in the activity series before hydrogen, is practically insoluble in dilute sulfuric acid and hydrochloric acids due to the formation of a film of insoluble salts (PbSO 4 and PbCl 2) on its surface.

You need to enable JavaScript to vote

Metals (from Latin metallum - mine, mine) - a group of elements in the form of simple substances with characteristic metallic properties such as high thermal and electrical conductivity, positive temperature coefficient of resistance, high ductility and metallic luster.

Of the 118 chemical elements discovered in this moment(not all of them are officially recognized), metals include:

• 6 elements in the alkali metal group,
• 6 in the group of alkaline earth metals,
• 38 in the group of transition metals,
• 11 in the group of light metals,
• 7 in the group of semimetals,
• 14 in the group lanthanides + lanthanum,
• 14 in the actinides group ( physical properties not studied for all elements) + sea anemone,
• outside of certain groups beryllium and magnesium.

Thus, 96 of all discovered elements may be metals.

In astrophysics, the term "metal" can have a different meaning and mean everything chemical elements heavier than helium

Characteristic properties of metals

1. Metallic luster (characteristic not only of metals: non-metals iodine and carbon in the form of graphite also have it)
2. Good electrical conductivity
3. Possibility of easy machining
4. High density (usually metals are heavier than non-metals)
5. High melting point (exceptions: mercury, gallium and alkali metals)
6. Great thermal conductivity
7. They are most often reducing agents in reactions.

Physical properties of metals

All metals (except mercury and, conditionally, France) at normal conditions are situated in solid state, however, they have different hardness. Below is the hardness of some metals on the Mohs scale.

Melting points pure metals range from −39 °C (mercury) to 3410 °C (tungsten). Most metals (except alkalis) have a high melting point, but some "normal" metals, such as tin and lead, can be melted on a regular electric or gas stove.

Depending on the density, metals are divided into light (density 0.53 ÷ 5 g/cm³) and heavy (5 ÷ 22.5 g/cm³). The lightest metal is lithium (density 0.53 g/cm³). It is currently impossible to name the heaviest metal, since the densities of osmium and iridium - the two heaviest metals - are almost equal (about 22.6 g/cm³ - exactly twice the density of lead), and calculating their exact density is extremely difficult: for this you need completely clean the metals, because any impurities reduce their density.

Most metals plastic, that is, the metal wire can be bent without breaking. This occurs due to the displacement of layers of metal atoms without breaking the bond between them. The most ductile are gold, silver and copper. Gold can be used to make foil 0.003 mm thick, which is used for gilding products. However, not all metals are ductile. Wire made of zinc or tin crunches when bent; When deformed, manganese and bismuth hardly bend at all, but immediately break. Plasticity also depends on the purity of the metal; Thus, very pure chromium is very ductile, but, contaminated with even minor impurities, it becomes brittle and harder. Some metals such as gold, silver, lead, aluminum, osmium can grow together, but this can take decades.

All metals are good conduct electric current; this is due to the presence in their crystal lattices of mobile electrons moving under the influence electric field. Silver, copper and aluminum have the highest electrical conductivity; for this reason, the latter two metals are most often used as wire materials. Sodium also has very high electrical conductivity; in experimental equipment, attempts are known to use sodium conductors in the form of thin-walled stainless steel pipes filled with sodium. Due to the low specific gravity of sodium, with equal resistance, sodium “wires” are much lighter than copper and even somewhat lighter than aluminum.

The high thermal conductivity of metals also depends on the mobility of free electrons. Therefore, the series of thermal conductivities is similar to the series of electrical conductivities, and the best conductor of heat, as well as electricity, is silver. Sodium also finds use as a good conductor of heat; It is widely known, for example, that sodium is used in valves of automobile engines to improve their cooling.

Color Most metals are approximately the same - light gray with a bluish tint. Gold, copper and cesium are yellow, red and light yellow, respectively.

Chemical properties of metals

At the external electronic level, most metals do not a large number of electrons (1-3), therefore in most reactions they act as reducing agents (that is, they “give up” their electrons)

Reactions with simple substances

• All metals except gold and platinum react with oxygen. The reaction with silver occurs at high temperatures, but silver(II) oxide is practically not formed, since it is thermally unstable. Depending on the metal, the output may include oxides, peroxides, and superoxides:

lithium oxide

sodium peroxide

potassium superoxide

To obtain an oxide from peroxide, the peroxide is reduced with a metal:

With medium and low-active metals, the reaction occurs when heated:

• Only the most active metals react with nitrogen; at room temperature only lithium reacts, forming nitrides:

When heated:

• All metals except gold and platinum react with sulfur:

Iron reacts with sulfur when heated, forming sulfide:

• Only the most active metals, that is, metals of groups IA and IIA except Be, react with hydrogen. Reactions occur when heated, and hydrides are formed. In reactions, the metal acts as a reducing agent, the oxidation state of hydrogen is −1:

• Only the most active metals react with carbon. In this case, acetylenides or methanides are formed. When reacting with water, acetylenides give acetylene, methanides give methane.

This lesson is devoted to the study of the topic “General properties of metals. Metal connection" During the lesson, the general chemical properties of metals and the features of metallic chemical bonds will be discussed. The teacher will explain the similarities in the chemical and physical properties of metals using a model of their internal structure.

Topic: Chemistry of metals

Lesson: General properties of metals. Metal connection

Metals are characterized by common physical properties: they have a special metallic luster, high thermal and electrical conductivity, and ductility.

Metals also share some common chemical properties. It is important to remember that in chemical reactions metals act as reducing agents: they donate electrons and increase their oxidation state. Let's look at some reactions in which metals participate.

INTERACTION WITH OXYGEN

Many metals can react with oxygen. Usually the products of these reactions are oxides, but there are exceptions, which you will learn about in the next lesson. Let's consider the interaction of magnesium with oxygen.

Magnesium burns in oxygen to form magnesium oxide:

2Mg + O2 = 2MgO

Rice. 1. Combustion of magnesium in oxygen

Magnesium atoms donate their outer electrons to oxygen atoms: two magnesium atoms donate two electrons to two oxygen atoms. In this case, magnesium acts as a reducing agent, and oxygen acts as an oxidizing agent.

Metals react with halogens. The product of this reaction is a metal halide, such as chloride.

Rice. 2. Combustion of potassium in chlorine

Potassium burns in chlorine to form potassium chloride:

2K + Cl 2 = 2KCl

Two potassium atoms donate one electron each to a chlorine molecule. Potassium, increasing the oxidation state, plays the role of a reducing agent, and chlorine, decreasing the oxidation state, plays the role of an oxidizing agent.

Many metals react with sulfur to form sulfides. In these reactions, metals also act as reducing agents, while sulfur will be an oxidizing agent. Sulfur in sulfides is in the oxidation state -2, i.e. it lowers its oxidation number from 0 to -2. For example, when heated, iron reacts with sulfur to form iron (II) sulfide:

Rice. 3. Interaction of iron with sulfur

Metals can also react with hydrogen, nitrogen, and other nonmetals under certain conditions.

Only active metals, such as alkali and alkaline earth metals, react with water without heating. During these reactions, an alkali is formed and hydrogen gas is released. For example, calcium reacts with water to form calcium hydroxide and hydrogen, releasing a large amount of heat:

Ca + 2H 2 O = Ca(OH) 2 + H 2

Less reactive metals, such as iron and zinc, react with water only when heated to form metal oxide and hydrogen. For example:

Zn + H 2 O = ZnO + H 2

In these reactions, the oxidizing agent is the hydrogen atom contained in water.

Metals located in the voltage series to the right of hydrogen do not react with water.

You already know that metals in the voltage series to the left of hydrogen react with acids. In these reactions, metals donate electrons and act as a reducing agent. The oxidizing agent is hydrogen cations formed in acid solutions. For example, zinc reacts with hydrochloric acid:

Zn + 2HCl = ZnCl 2 + H 2

Otherwise, reactions of metals with nitric and concentrated sulfuric acids occur. In these reactions, practically no hydrogen is released. We'll talk about these interactions in future lessons.

A metal can react with a salt solution if it is more active than the metal contained in the salt. For example, iron replaces copper from copper(II) sulfate:

Fe + CuSO 4 = FeSO 4 + Cu

Iron is a reducing agent, copper cations are an oxidizing agent.

Let's try to explain why metals have common physical and chemical properties. To do this, consider a model of the internal structure of a metal.

Metal atoms have relatively large radii and a small number of outer electrons. These electrons are weakly attracted to the nucleus, so in chemical reactions metals act as reducing agents, donating electrons from the outer energy level.

At the nodes of the crystal lattice of metals there are not only neutral atoms, but also metal cations, because outer electrons move freely around crystal lattice. In this case, atoms, giving up electrons, become cations, and cations, accepting electrons, turn into electrically neutral atoms.

Rice. 4. Model of the internal structure of metal

A chemical bond that is formed as a result of the attraction of metal cations to freely moving electrons is called metal.

The electrical and thermal conductivity of metals is explained by the presence of free electrons, which can be carriers electric current and carriers of heat. The plasticity of the metal is explained by the fact that it does not tear under mechanical stress. chemical bond, because a chemical bond is established not between specific atoms and cations, but between all metal cations with all free electrons in the metal crystal.

1. Mikityuk A.D. Collection of problems and exercises in chemistry. 8-11 grades / A.D. Mikityuk. - M.: Publishing house. "Exam", 2009.

2. Orzhekovsky P.A. Chemistry: 9th grade: textbook. for general education establishment / P.A. Orzhekovsky, L.M. Meshcheryakova, L.S. Pontak. - M.: AST: Astrel, 2007. (§23)

3. Orzhekovsky P.A. Chemistry: 9th grade: general education. establishment / P.A. Orzhekovsky, L.M. Meshcheryakova, M.M. Shalashova. - M.: Astrel, 2013. (§6)

4. Rudzitis G.E. Chemistry: inorganic. chemistry. Organ. chemistry: textbook. for 9th grade. / G.E. Rudzitis, F.G. Feldman. - M.: Education, OJSC “Moscow Textbooks”, 2009.

5. Khomchenko I.D. Collection of problems and exercises in chemistry for high school. - M.: RIA “New Wave”: Publisher Umerenkov, 2008.

6. Encyclopedia for children. Volume 17. Chemistry / Chapter. ed. V.A. Volodin, Ved. scientific ed. I. Leenson. - M.: Avanta+, 2003.

Additional web resources

1. Single collection of digital educational resources(video experiments on the topic) ().

2. Electronic version magazine "Chemistry and Life" ().

Homework

p.41 Nos. A1, A2 from the Textbook of Orzhekovsky P.A. “Chemistry: 9th grade” (M.: Astrel, 2013).