CHEMICAL PROPERTIES OF METALS

1 ) Active (alkali and alkaline earth metals, Mg, Al, Zn, etc.)

2) Metalsaverage activity (Fe, Cr, Mn, etc.);

3 ) Inactive (Cu, Ag)

4) noble metals – Au, Pt, Pd, etc.

In reactions - only reducing agents. Metal atoms easily donate electrons from the outer (and some of them from the pre-outer) electron layer, turning into positive ions. Possible oxidation states Me Lower 0,+1,+2,+3 Higher +4,+5,+6,+7,+8

1. INTERACTION WITH NON-METALS

1. WITH HYDROGEN

Metals of groups IA and IIA react when heated, except for beryllium. Solid unstable substances hydrides are formed, other metals do not react.

2K + H₂ = 2KH (potassium hydride)

Ca + H₂ = CaH₂

2. WITH OXYGEN

All metals react except gold and platinum. The reaction with silver occurs at high temperatures, but silver(II) oxide is practically not formed, since it is thermally unstable. Alkali metals at normal conditions form oxides, peroxides, superoxides (lithium - oxide, sodium - peroxide, potassium, cesium, rubidium - superoxide

4Li + O2 = 2Li2O (oxide)

2Na + O2 = Na2O2 (peroxide)

K+O2=KO2 (superoxide)

The remaining metals of the main subgroups under normal conditions form oxides with an oxidation state equal to the group number 2Сa + O2 = 2СaO

2Сa+O2=2СaO

The metals of the secondary subgroups form oxides under normal conditions and when heated, oxides of various degrees of oxidation, and iron iron scale Fe3O4 (Fe⁺²O∙Fe2⁺³O3)

3Fe + 2O2 = Fe3O4

4Cu + O₂ = 2Cu₂⁺¹O (red) 2Cu + O₂ = 2Cu⁺²O (black);

2Zn + O₂ = ZnO 4Cr + 3O2 = 2Cr2O3

3. WITH HALOGENS

halides (fluorides, chlorides, bromides, iodides). Alkaline under normal conditions with F, Cl, Br ignite:

2Na + Cl2 = 2NaCl (chloride)

Alkaline earth and aluminum react under normal conditions:

WITHa+Cl2=WITHaCl2

2Al+3Cl2 = 2AlCl3

Metals of secondary subgroups at elevated temperatures

Cu + Cl₂ = Cu⁺²Cl₂ Zn + Cl₂ = ZnCl₂

2Fe + ЗС12 = 2Fe⁺³Cl3 iron chloride (+3) 2Cr + 3Br2 = 2Cr⁺³Br3

2Cu + I₂ = 2Cu⁺¹I(there is no copper iodide (+2)!)

4. INTERACTION WITH SULFUR

when heated even with alkali metals, with mercury under normal conditions. All metals react except gold and platinum

Withgray – sulfides: 2K + S = K2S 2Li+S = Li2S (sulfide)

WITHa+S=WITHaS(sulfide) 2Al+3S = Al2S3 Cu + S = Cu⁺²S (black)

Zn + S = ZnS 2Cr + 3S = Cr2⁺³S3 Fe + S = Fe⁺²S

5. INTERACTION WITH PHOSPHORUS AND NITROGEN

leaks when heated (exception: lithium with nitrogen under normal conditions) :

with phosphorus - phosphides: 3Ca + 2 P=Ca3P2,

With nitrogen - nitrides 6Li + N2 = 3Li2N (lithium nitride) (n.o.) 3Mg + N2 = Mg3N2 (magnesium nitride) 2Al + N2 = 2A1N 2Cr + N2 = 2CrN 3Fe + N2 = Fe₃⁺²N₂¯³

6. INTERACTION WITH CARBON AND SILICON

flows when heated:

Carbides are formed with carbon. Only the most active metals react with carbon. From alkali metals, carbides form lithium and sodium, potassium, rubidium, cesium do not interact with carbon:

2Li + 2C = Li2C2, Ca + 2C = CaC2

Metals - d-elements form compounds of non-stoichiometric composition such as solid solutions with carbon: WC, ZnC, TiC - are used to obtain superhard steels.

with silicon - silicides: 4Cs + Si = Cs4Si,

7. INTERACTION OF METALS WITH WATER:

Metals that reach hydrogen in the electrochemical series of voltages react with water. Alkali and alkaline earth metals react with water without heating, forming soluble hydroxides (alkalis) and hydrogen, aluminum (after the destruction of the oxide film - amalgation), magnesium, when heated, form insoluble bases and hydrogen .

2Na + 2HOH = 2NaOH + H2
WITHa + 2HOH = Ca(OH)2 + H2

2Al + 6H2O = 2Al(OH)3 + ZH2

The remaining metals react with water only in a hot state, forming oxides (iron - iron scale)

Zn + H2O = ZnO + H2 3Fe + 4HOH = Fe3O4 + 4H2 2Cr + 3H₂O = Cr₂O₃ + 3H₂

8 WITH OXYGEN AND WATER

In air, iron and chromium easily oxidize in the presence of moisture (rusting)

4Fe + 3O2 + 6H2O = 4Fe(OH)3

4Cr + 3O2 + 6H2O = 4Cr(OH)3

9. INTERACTION OF METALS WITH OXIDES

Metals (Al, Mg, Ca), at high temperature reduce non-metals or less active metals from their oxides → non-metal or low-active metal and oxide (calciumthermy, magnesiumthermy, aluminothermy)

2Al + Cr2O3 = 2Cr + Al2O3 3Са + Cr₂O₃ = 3СаО + 2Cr (800 °C) 8Al + 3Fe3O4 = 4Al2O3 + 9Fe (thermite) 2Mg + CO2 = 2MgO + С Mg + N2O = MgO + N2 Zn + CO2 = ZnO + CO 2Cu + 2NO = 2CuO + N2 3Zn + SO2 = ZnS + 2ZnO

10. WITH OXIDES

Metals iron and chromium react with oxides, reducing the degree of oxidation

Cr + Cr2⁺³O3 = 3Cr⁺²O Fe+ Fe2⁺³O3 = 3Fe⁺²O

11. INTERACTION OF METALS WITH ALKALI

Only those metals interact with alkalis, the oxides and hydroxides of which have amphoteric properties ((Zn, Al, Cr (III), Fe (III), etc. MELT → metal salt + hydrogen.

2NaOH + Zn → Na2ZnO2 + H2 (sodium zincate)

2Al + 2(NaOH H2O) = 2NaAlO2 + 3H2
SOLUTION → complex metal salt + hydrogen.

2NaOH + Zn0 + 2H2O = Na2 + H2 (sodium tetrahydroxozincate) 2Al + 2NaOH + 6H2O = 2Na + 3H2

12. INTERACTION WITH ACIDS (EXCEPT HNO3 and H2SO4 (conc.)

Metals standing in the electrochemical series of voltages of metals to the left of hydrogen displace it from dilute acids → salt and hydrogen

Remember! Nitric acid never releases hydrogen when interacting with metals.

Mg + 2HC1 = MgCl2 + H2
Al + 2HC1 = Al⁺³Cl₃ + H2

13. REACTIONS WITH SALT

Active metals displace less active metals from salts. Recovery from solutions:

CuSO4 + Zn = ZnSO4 + Cu

FeSO4 + Cu =REACTIONSNO

Mg + CuCl2(pp) = MgCl2 +WITHu

Recovery of metals from melts of their salts

3Na+ AlCl₃ = 3NaCl + Al

TiCl2 + 2Mg = MgCl2 + Ti

Group B metals react with salts, lowering their oxidation state.

2Fe⁺³Cl3 + Fe = 3Fe⁺²Cl2

Metals are active reducing agents with a positive oxidation state. Due to their chemical properties, metals are widely used in industry, metallurgy, medicine, and construction.

Metal activity

In reactions, metal atoms donate valence electrons and are oxidized. The more energy levels and fewer electrons a metal atom has, the easier it is for it to donate electrons and enter into reactions. Therefore, metallic properties increase from top to bottom and from right to left in the periodic table.

Rice. 1. Change in metallic properties in the periodic table.

Activity simple substances shown in the electrochemical series of voltages of metals. To the left of hydrogen are active metals (activity increases towards the left edge), to the right - inactive.

Alkali metals in group I show the greatest activity. periodic table and standing to the left of hydrogen in the electrochemical series of voltages. They react with many substances already at room temperature. They are followed by alkaline earth metals, which are included in group II. They react with most substances when heated. Metals that are in the electrochemical series from aluminum to hydrogen (medium activity) require additional conditions for entering into reactions.

Rice. 2. Electrochemical series of voltages of metals.

Some metals exhibit amphoteric properties or duality. Metals, their oxides and hydroxides react with acids and bases. Most metals only react with certain acids to replace the hydrogen and form a salt. The most pronounced dual properties show:

• aluminum;
• lead;
• zinc;
• iron;
• copper;
• beryllium;
• chromium.

Each metal is capable of displacing another metal to the right of it in the electrochemical series from salts. Metals to the left of hydrogen displace it from dilute acids.

Properties

Features of the interaction of metals with different substances are presented in the table of chemical properties of metals.

Metals interact with each other and form intermetallic compounds - 3Cu + Au → Cu 3 Au, 2Na + Sb → Na 2 Sb.

Application

The general chemical properties of metals are used to create alloys, detergents, and are used in catalytic reactions. Metals are present in batteries, electronics, and load-bearing structures.

The main fields of application are indicated in the table.

Rice. 3. Bismuth.

What have we learned?

From the 9th grade chemistry lesson, we learned about the basic chemical properties of metals. The ability to interact with simple and complex substances determines the activity of metals. The more active the metal, the easier it reacts under normal conditions. Active metals react with halogens, non-metals, water, acids, salts. Amphoteric metals interact with alkalis. Inactive metals do not react with water, halogens, and most non-metals. Briefly reviewed the application areas. Metals are used in medicine, industry, metallurgy, and electronics.

Topic quiz

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By metals is meant a group of elements, which is presented in the form of the simplest substances. They have characteristic properties, namely high electrical and thermal conductivity, positive temperature coefficient of resistance, high ductility and metallic luster.

Note that of the 118 chemical elements that have been discovered so far, metals should include:

• among the group of alkaline earth metals 6 elements;
• among alkali metals 6 elements;
• among transition metals 38;
• in the group of light metals 11;
• among semimetals 7 elements,
• 14 among the lanthanides and lanthanum,
• 14 in the group of actinides and actiniums,
• Outside the definition are beryllium and magnesium.

Based on this, 96 elements belong to metals. Let's take a closer look at what metals react with. Since most metals have a small number of electrons from 1 to 3 on the external electronic level, they can act as reducing agents in most of their reactions (that is, they donate their electrons to other elements).

Reactions with the simplest elements

• In addition to gold and platinum, absolutely all metals react with oxygen. Note also that the reaction occurs with silver at high temperatures, but silver(II) oxide is not formed at normal temperatures. Depending on the properties of the metal, as a result of the reaction with oxygen, oxides, superoxides and peroxides are formed.

Here are examples of each of the chemical formations:

1. lithium oxide - 4Li + O 2 \u003d 2Li 2 O;
2. potassium superoxide - K + O 2 \u003d KO 2;
3. sodium peroxide - 2Na + O 2 \u003d Na 2 O 2.

In order to obtain oxide from peroxide, it must be reduced with the same metal. For example, Na 2 O 2 + 2Na \u003d 2Na 2 O. With low-active and medium metals, a similar reaction will occur only when heated, for example: 3Fe + 2O 2 \u003d Fe 3 O 4.

• Metals can react with nitrogen only with active metals, however, only lithium can interact at room temperature, forming nitrides - 6Li + N 2 \u003d 2Li 3 N, however, when heated, such a chemical reaction occurs 2Al + N 2 \u003d 2AlN, 3Ca + N 2 = Ca 3 N 2 .
• Absolutely all metals react with sulfur, as well as with oxygen, with the exception of gold and platinum. Note that iron can only interact when heated with sulfur, forming a sulfide: Fe+S=FeS
• Only active metals can react with hydrogen. These include metals of groups IA and IIA, except for beryllium. Such reactions can be carried out only when heated, forming hydrides.

  Since the oxidation state of hydrogen is considered? 1, then the metals in this case act as reducing agents: 2Na + H 2 \u003d 2NaH.

• The most active metals also react with carbon. As a result of this reaction, acetylenides or methanides are formed.

Consider which metals react with water and what do they give as a result of this reaction? Acetylenes, when interacting with water, will give acetylene, and methane will be obtained as a result of the reaction of water with methanides. Here are examples of these reactions:

1. Acetylene - 2Na + 2C \u003d Na 2 C 2;
2. Methane - Na 2 C 2 + 2H 2 O \u003d 2NaOH + C 2 H 2.

Reaction of acids with metals

Metals with acids can also react differently. With all acids, only those metals react that are in the series of the electrochemical activity of metals to hydrogen.

Let's give an example of a substitution reaction, which shows what metals react with. In another way, such a reaction is called a redox reaction: Mg + 2HCl \u003d MgCl 2 + H 2 ^.

Some acids can also interact with metals that are after hydrogen: Cu + 2H 2 SO 4 \u003d CuSO 4 + SO 2 ^ + 2H 2 O.

Note that such a dilute acid can react with a metal according to the following classical scheme: Mg + H 2 SO 4 \u003d MgSO 4 + H 2 ^.

Metal atoms give up valence electrons relatively easily and pass into positively charged ions. Therefore, metals are reducing agents. Metals interact with simple substances: Ca + C12 - CaC12. Active metals react with water: 2Na + 2H20 = 2NaOH + H2f. Metals standing in the series of standard electrode potentials up to hydrogen interact with dilute solutions of acids (except HN03) with hydrogen evolution: Zn + 2HC1 = ZnCl2 + H2f. Metals react with aqueous solutions of salts of less active metals: Ni + CuSO4 = NiSO4 + Cu J. Metals react with oxidizing acids: C. Methods for obtaining metals Modern metallurgy receives more than 75 metals and numerous alloys based on them. Depending on the methods of obtaining metals, pyrohydro- and electrometallurgy are distinguished. GD) Pyrometallurgy covers methods for obtaining metals from ores using reduction reactions carried out at high temperatures. Coal, active metals, carbon monoxide (II), hydrogen, methane are used as reducing agents. Cu20 + C - 2Cu + CO, t° Cu20 + CO - 2Cu + CO2, t° Cr203 + 2A1 - 2Cr + A1203, (aluminum) t° TiCl2 + 2Mg - Ti + 2MgCl2, (magnesium) t° W03 + 3H2 = W + 3H20. (hydrothermy) | C Hydrometallurgy is the production of metals from solutions of their salts. For example, when copper ore containing copper (I) oxide is treated with dilute sulfuric acid, copper goes into solution in the form of sulfate: CuO + H2S04 = CuS04 + H20. Then copper is extracted from the solution either by electrolysis or by displacement with iron powder: CuSO4 + Fe = FeSO4 + Cu. [h] Electrometallurgy is a method of obtaining metals from their molten oxides or salts using electrolysis: electrolysis 2NaCl - 2Na + Cl2. Questions and tasks for independent solution 1. Indicate the position of metals in periodic system D. I. Mendeleev. 2. Show the physical and chemical properties of metals. 3. Explain the reason for the commonality of the properties of metals. 4. Show the change in the chemical activity of metals of the main subgroups of groups I and II of the periodic system. 5. How do the metallic properties of the elements of II and III periods change? Name the most refractory and the most fusible metals. 7. Indicate which metals are found in nature in the native state and which are found only in the form of compounds. How can this be explained? 8. What is the nature of alloys? How does the composition of an alloy affect its properties? Show with specific examples. Specify the most important ways obtaining metals from ores. 10l Name the varieties of pyrometallurgy. What reducing agents are used in each specific method? Why? 11. Name the metals that are obtained using hydrometallurgy. What is the essence and what are the benefits this method in front of others? 12. Give examples of obtaining metals using electrometallurgy. In what case is this method used? 13. What are the modern methods of obtaining high purity metals? 14. What is "electrode potential"? Which of the metals has the largest and which - the smallest electrode potentials in an aqueous solution? 15. Describe a number of standard electrode potentials? 16. Is it possible to displace metallic iron from an aqueous solution of its sulfate using metallic zinc, nickel, sodium? Why? 17. What is the principle of operation of galvanic cells? What metals can be used in them? 18. What processes are corrosive? What types of corrosion do you know? 19. What is called electrochemical corrosion? What methods of protection against it do you know? 20. How does contact with other metals affect the corrosion of iron? What metal will be destroyed first on the damaged surface of tin-plated, galvanized and nickel-plated iron? 21. What process is called electrolysis? Write reactions that reflect the processes occurring at the cathode and anode during the electrolysis of sodium chloride melt, aqueous solutions of sodium chloride, copper sulfate, sodium sulfate, sulfuric acid. 22. What role does the electrode material play in the course of electrolysis processes? Give examples of electrolysis processes occurring with soluble and insoluble electrodes. 23. The alloy used to make copper coins contains 95% copper. Determine the second metal included in the alloy, if when processing a one-kopeck coin with an excess of hydrochloric acid 62.2 ml of hydrogen (n.a.) were released. aluminum. 24. A sample of metal carbide weighing 6 g was burned in oxygen. This formed 2.24 l of carbon monoxide (IV) (n.a.). Determine what metal was part of the carbide. 25. Show what products will be released during the electrolysis of an aqueous solution of nickel sulfate, if the process proceeds: a) with coal; b) with nickel electrodes? 26. During the electrolysis of an aqueous solution of copper sulphate, 2.8 liters of gas (n.a.) were released at the anode. What is this gas? What and in what quantity was released at the cathode? 27. Make a diagram of the electrolysis of an aqueous solution of potassium nitrate flowing on the electrodes. What is the amount of electricity passed if 280 ml of gas (n.a.) is released at the anode? What and in what quantity was released at the cathode?

Chemical properties metals: interaction with oxygen, halogens, sulfur and relation to water, acids, salts.

The chemical properties of metals are due to the ability of their atoms to easily give up electrons from an external energy level, turning into positively charged ions. Thus, in chemical reactions, metals act as energetic reducing agents. This is their main common chemical property.

The ability to donate electrons in atoms of individual metallic elements is different. The more easily a metal gives up its electrons, the more active it is, and the more vigorously it reacts with other substances. Based on the research, all metals were arranged in a row according to their decreasing activity. This series was first proposed by the outstanding scientist N. N. Beketov. Such a series of activity of metals is also called the displacement series of metals or the electrochemical series of metal voltages. It looks like this:

Li, K, Ba, Ca, Na, Mg, Al, Zn, Fe, Ni, Sn, Pb, H2, Cu, Hg, Ag, Рt, Au

Using this series, you can find out which metal is the active of the other. This series contains hydrogen, which is not a metal. Its visible properties are taken for comparison as a kind of zero.

Having the properties of reducing agents, metals react with various oxidizing agents, primarily with non-metals. Metals react with oxygen under normal conditions or when heated to form oxides, for example:

2Mg0 + O02 = 2Mg+2O-2

In this reaction, magnesium atoms are oxidized and oxygen atoms are reduced. The noble metals at the end of the row react with oxygen. Reactions with halogens actively occur, for example, the combustion of copper in chlorine:

Cu0 + Cl02 = Cu+2Cl-2

Reactions with sulfur most often occur when heated, for example:

Fe0 + S0 = Fe+2S-2

Active metals in the activity series of metals in Mg react with water to form alkalis and hydrogen:

2Na0 + 2H+2O → 2Na+OH + H02

Metals of medium activity from Al to H2 react with water under more severe conditions and form oxides and hydrogen:

Pb0 + H+2O Chemical properties of metals: interaction with oxygen Pb+2O + H02.

The ability of a metal to react with acids and salts in solution also depends on its position in the displacement series of metals. Metals to the left of hydrogen in the displacement series of metals usually displace (reduce) hydrogen from dilute acids, and metals to the right of hydrogen do not displace it. So, zinc and magnesium react with acid solutions, releasing hydrogen and forming salts, while copper does not react.

Mg0 + 2H+Cl → Mg+2Cl2 + H02

Zn0 + H+2SO4 → Zn+2SO4 + H02.

Metal atoms in these reactions are reducing agents, and hydrogen ions are oxidizing agents.

Metals react with salts in aqueous solutions. Active metals displace less active metals from the composition of salts. This can be determined from the activity series of metals. The reaction products are a new salt and a new metal. So, if an iron plate is immersed in a solution of copper (II) sulfate, after a while copper will stand out on it in the form of a red coating:

Fe0 + Cu+2SO4 → Fe+2SO4 + Cu0 .

But if a silver plate is immersed in a solution of copper (II) sulfate, then no reaction will occur:

Ag + CuSO4 ≠ .

To carry out such reactions, one should not take too active metals (from lithium to sodium), which are capable of reacting with water.

Therefore, metals are able to react with non-metals, water, acids and salts. In all these cases, the metals are oxidized and are reducing agents. To predict the course of chemical reactions involving metals, a displacement series of metals should be used.