Possible reactions of dem water with ferrous metal. What do metals react with? IV. Displacement of less active metals by more active metals from solutions of their salts

Metals mean a group of elements, which are 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, the following should be classified as metals:

• among the group of alkaline earth metals there are 6 elements;
• among alkali metals there are 6 elements;
• among transition metals 38;
• in the group of light metals 11;
• There are 7 elements among semimetals,
• 14 among lanthanides and lanthanum,
• 14 in the group of actinides and sea anemones,
• Beryllium and magnesium are outside the definition.

Based on this, 96 elements are classified as 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 at the outer electronic level, in most of their reactions they can act as reducing agents (that is, they give up their electrons to other elements).

Reactions with the simplest elements

• Except for 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, oxides, superoxides and peroxides are formed as a result of reaction with oxygen.

Here are examples of each chemical education:

1. lithium oxide – 4Li+O 2 =2Li 2 O;
2. potassium superoxide – K+O 2 =KO 2;
3. sodium peroxide – 2Na+O 2 =Na 2 O 2.

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

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

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

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

Let's consider what metals react with water and what do they produce as a result of this reaction? Acetylenes, when interacting with water, will produce 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= Na 2 C 2 ;
2. Methane - Na 2 C 2 +2H 2 O=2NaOH+C 2 H 2.

Reaction of acids with metals

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

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

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

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

Goal of the work: practically become familiar with the characteristic chemical properties of metals of various activities and their compounds; study the features of metals with amphoteric properties. Redox reactions are equalized using the electron-ion balance method.

Theoretical part

Physical properties of metals. Under normal conditions, all metals, except mercury, are solid substances that differ sharply in the degree of hardness. Metals, being conductors of the first kind, have high electrical and thermal conductivity. These properties are associated with the structure of the crystal lattice, in the nodes of which there are metal ions, between which free electrons move. The transfer of electricity and heat occurs due to the movement of these electrons.

Chemical properties of metals . All metals are reducing agents, i.e. During chemical reactions they lose electrons and become positively charged ions. As a result, most metals react with typical oxidizing agents, such as oxygen, forming oxides, which in most cases cover the surface of metals in a dense layer.

Mg° +O 2 °=2Mg +2 O- 2

Mg-2=Mg +2

ABOUT 2 +4 =2О -2

The reducing activity of metals in solutions depends on the position of the metal in the voltage series or on the value of the electrode potential of the metal (table). The lower the electrode potential of a given metal, the more active a reducing agent it is. All metals can be divided into 3 groups :

Active metals – from the beginning of the stress series (i.e. from Li) to Mg;

Intermediate activity metals from Mg to H;

Low-active metals – from H to the end of the voltage series (to Au).

Metals of group 1 interact with water (this includes mainly alkali and alkaline earth metals); The reaction products are hydroxides of the corresponding metals and hydrogen, for example:

2К°+2Н 2 O=2KOH+H 2 ABOUT

K°-=K + | 2

2H + +2 =H 2 0 | 1

Interaction of metals with acids

All oxygen-free acids (hydrochloric HCl, hydrobromic HBr, etc.), as well as some oxygen-containing acids (dilute sulfuric acid H 2 SO 4, phosphoric acid H 3 PO 4, acetic acid CH 3 COOH, etc.) react with metals 1 and 2 groups standing in the voltage series up to hydrogen. In this case, the corresponding salt is formed and hydrogen is released:

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

Zn 0 -2 = Zn 2+ | 1

2H + +2 =H 2 ° | 1

Concentrated sulfuric acid oxidizes metals of groups 1, 2 and partially 3 (up to Ag inclusive) while being reduced to SO 2 - a colorless gas with a pungent odor, free sulfur precipitated in the form of a white precipitate or hydrogen sulfide H 2 S - a gas with a rotten odor eggs The more active the metal, the more sulfur is reduced, for example:

| 1

| 8

Nitric acid of any concentration oxidizes almost all metals, resulting in the formation of nitrate of the corresponding metal, water and the reduction product N +5 (NO 2 - brown gas with a pungent odor, NO - colorless gas with a pungent odor, N 2 O - gas with a narcotic odor, N 2 is an odorless gas, NH 4 NO 3 is a colorless solution). The more active the metal and the more dilute the acid, the more nitrogen is reduced in nitric acid.

React with alkalis amphoteric metals belonging mainly to group 2 (Zn, Be, Al, Sn, Pb, etc.). The reaction proceeds by fusing metals with alkali:

Pb+2 NaOH= Na 2 PbO 2 +H 2

Pb 0 -2 = Pb 2+ | 1

2H + +2 =H 2 ° | 1

or when interacting with a strong alkali solution:

Be + 2NaOH + 2H 2 ABOUT = Na 2 +H 2

Be°-2=Be +2 | 1

Amphoteric metals form amphoteric oxides and, accordingly, amphoteric hydroxides (reacting with acids and alkalis to form salts and water), for example:

or in ionic form:

or in ionic form:

Practical part

Experience No. 1.Interaction of metals with water .

Take a small piece of alkali or alkaline earth metal (sodium, potassium, lithium, calcium), which is stored in a jar of kerosene, dry it thoroughly with filter paper, and add it to a porcelain cup filled with water. At the end of the experiment, add a few drops of phenolphthalein and determine the medium of the resulting solution.

When magnesium reacts with water, heat the reaction tube for some time on an alcohol lamp.

Experience No. 2.Interaction of metals with dilute acids .

Pour 20 - 25 drops of 2N solution of hydrochloric, sulfuric and nitric acids into three test tubes. Drop metals in the form of wires, pieces or shavings into each test tube. Observe the phenomena taking place. Heat the test tubes in which nothing happens in an alcohol lamp until the reaction begins. Carefully sniff the test tube containing nitric acid to determine the gas released.

Experience No. 3.Interaction of metals with concentrated acids .

Pour 20 - 25 drops of concentrated nitric and sulfuric (carefully!) acids into two test tubes, lower the metal into them, and observe what happens. If necessary, the test tubes can be heated in an alcohol lamp before the reaction begins. To determine the gases released, carefully sniff the tubes.

Experiment No. 4.Interaction of metals with alkalis .

Pour 20 - 30 drops of a concentrated alkali solution (KOH or NaOH) into a test tube and add the metal. Warm the test tube slightly. Observe what is happening.

Experience№5. Receipt and properties metal hydroxides.

Pour 15-20 drops of salt of the corresponding metal into a test tube, add alkali until a precipitate forms. Divide the sediment into two parts. Pour a hydrochloric acid solution to one part, and an alkali solution to the other. Note the observations, write equations in molecular, full ionic and short ionic forms, and draw conclusions about the nature of the resulting hydroxide.

Design of the work and conclusions

Write electron-ion balance equations for redox reactions, write ion exchange reactions in molecular and ion-molecular forms.

In your conclusions, write which activity group (1, 2 or 3) the metal you studied belongs to and what properties - basic or amphoteric - its hydroxide exhibits. Justify your conclusions.

Laboratory work No. 11

The structure of metal atoms determines not only the characteristic physical properties of simple substances - metals, but also their general chemical properties.

With great diversity, all chemical reactions of metals are redox and can be of only two types: combination and substitution. Metals are capable of donating electrons during chemical reactions, that is, being reducing agents and exhibiting only a positive oxidation state in the resulting compounds.

In general, this can be expressed by the following diagram:
Me 0 – ne → Me +n,
where Me is a metal - a simple substance, and Me 0+n is a metal, a chemical element in a compound.

Metals are capable of donating their valence electrons to non-metal atoms, hydrogen ions, and ions of other metals, and therefore will react with non-metals - simple substances, water, acids, salts. However, the reducing ability of metals varies. The composition of the reaction products of metals with various substances depends on the oxidizing ability of the substances and the conditions under which the reaction occurs.

At high temperatures, most metals burn in oxygen:

2Mg + O2 = 2MgO

Only gold, silver, platinum and some other metals do not oxidize under these conditions.

Many metals react with halogens without heating. For example, aluminum powder, when mixed with bromine, ignites:

2Al + 3Br 2 = 2AlBr 3

When metals interact with water, hydroxides are formed in some cases. Under normal conditions, alkali metals, as well as calcium, strontium, and barium, interact very actively with water. The general scheme of this reaction looks like this:

Me + HOH → Me(OH) n + H 2

Other metals react with water when heated: magnesium when it boils, iron in water vapor when it boils red. In these cases, metal oxides are obtained.

If a metal reacts with an acid, it is part of the resulting salt. When a metal interacts with acid solutions, it can be oxidized by hydrogen ions present in the solution. The abbreviated ionic equation can be written in general form as follows:

Me + nH + → Me n + + H 2

The anions of oxygen-containing acids, such as concentrated sulfuric and nitric, have stronger oxidizing properties than hydrogen ions. Therefore, those metals that are not able to be oxidized by hydrogen ions, for example, copper and silver, react with these acids.

When metals interact with salts, a substitution reaction occurs: electrons from the atoms of the replacing – more active metal – pass to the ions of the replaced – less active metal. Then the network replaces metal with metal in salts. These reactions are not reversible: if metal A displaces metal B from the salt solution, then metal B will not displace metal A from the salt solution.

In descending order of chemical activity manifested in the reactions of displacement of metals from each other from aqueous solutions of their salts, metals are located in the electrochemical series of voltages (activities) of metals:

Li → Rb → K → Ba → Sr → Ca → Na→ Mg → Al → Mn → Zn → Cr → → Fe → Cd→ Co → Ni → Sn → Pb → H → Sb → Bi → Cu → → Ag → Pd → Pt → Au

Metals located to the left in this row are more active and are able to displace the following metals from salt solutions.

Hydrogen is included in the electrochemical series of voltages of metals, as the only non-metal that shares a common property with metals - to form positively charged ions. Therefore, hydrogen replaces some metals in their salts and can itself be replaced by many metals in acids, for example:

Zn + 2 HCl = ZnCl 2 + H 2 + Q

Metals that come before hydrogen in the electrochemical voltage series displace it from solutions of many acids (hydrochloric, sulfuric, etc.), but all those following it, for example, copper, do not displace it.

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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 give up valence electrons and become oxidized. The more energy levels and fewer electrons a metal atom has, the easier it is for it to give up electrons and undergo reactions. Therefore, metallic properties increase from top to bottom and from right to left in the periodic table.

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

The activity of simple substances is shown in the electrochemical voltage series of metals. To the left of hydrogen are active metals (activity increases towards the left), to the right are inactive metals.

The greatest activity is exhibited by alkali metals that are in group I of the periodic table and are to the left of hydrogen in the electrochemical voltage series. 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 in the electrochemical series from aluminum to hydrogen (medium activity) require additional conditions to enter into reactions.

Rice. 2. Electrochemical series of metal voltages.

Some metals exhibit amphoteric properties or duality. Metals, their oxides and hydroxides react with acids and bases. Most metals react only with certain acids, displacing hydrogen and forming a salt. The most pronounced dual properties are exhibited by:

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

Each metal is capable of displacing another metal standing 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 supporting structures.

The main areas of application are listed 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 a metal is, the more easily it reacts under normal conditions. Active metals react with halogens, nonmetals, water, acids, and salts. Amphoteric metals react with alkalis. Low-active metals do not react with water, halogens, and most non-metals. We briefly reviewed the areas of application. Metals are used in medicine, industry, metallurgy, and electronics.

Test on the topic

Evaluation of the report

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Metal atoms relatively easily give up valence electrons and become positively charged ions. Therefore, metals are reducing agents. Metals react 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 for HN03) with the release of hydrogen: Zn + 2HC1 = ZnCl2 + H2f. Metals react with aqueous solutions of salts of less active metals: Ni + CuS04 = NiS04 + Cu J. Metals react with oxidizing acids: C. Methods for producing metals Modern metallurgy produces more than 75 metals and numerous alloys based on them. Depending on the methods of obtaining metals, pyrohydro- and electrometallurgy are distinguished. GG) Pyrometallurgy covers methods of obtaining metals from ores using reduction reactions carried out at high temperatures. Coal, active metals, carbon monoxide (II), hydrogen, and methane are used as reducing agents. Cu20 + C - 2Cu + CO, t° Cu20 + CO - 2Cu + C02, t° Cr203 + 2A1 - 2Cg + A1203, (aluminothermy) t° TiCl2 + 2Mg - Ti + 2MgCl2, (magnesiumthermy) t° W03 + 3H2 = W+3H20. (hydrogenothermy) |C Hydrometallurgy is the production of metals from solutions of their salts. For example, when copper ore containing copper oxide (I) is treated with dilute sulfuric acid, copper goes into solution in the form of sulfate: CuO + H2S04 = CuS04 + H20. Copper is then removed from the solution either by electrolysis or by displacement using iron powder: CuS04 + Fe = FeS04 + Cu. [h] Electrometallurgy is methods for producing 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 the periodic table of D.I. Mendeleev. 2. Show the physical and chemical properties of metals. 3. Explain the reason for the common 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 table. 5. How do the metallic properties of elements of periods II and III change? Name the most refractory and the most fusible metals. 7. Indicate which metals are found in nature in a 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. Indicate the most important methods for obtaining metals from ores. 10l Name the types 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 advantages of this method over others? 12. Give examples of the production of metals using electrometallurgy. In what case is this method used? 13. What are the modern methods for producing high-purity metals? 14. What is “electrode potential”? Which metal has the highest and which has the lowest electrode potential 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, and sodium? Why? 17. What is the principle of operation of galvanic cells? What metals can be used in them? 18. What processes are classified as corrosion? 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 its contact with other metals affect the corrosion of iron? Which metal will be destroyed first on a damaged surface of tinned, 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 molten sodium chloride, aqueous solutions of sodium chloride, copper sulfate, sodium sulfate, sulfuric acid. 22. What role does the electrode material play during electrolysis processes? Give examples of electrolysis processes occurring with soluble and insoluble electrodes. 23. The alloy used to prepare 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.u.) was released. aluminum. 24. A sample of metal carbide weighing 6 g was burned in oxygen. In this case, 2.24 liters of carbon monoxide (IV) (no.) was formed. Determine what metal was included in 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 sulfate, 2.8 liters of gas (n.e.) were released at the anode. What gas is this? What and in what quantity was released at the cathode? 27. Draw up 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) is released at the anode? u.)? What and in what quantity was released at the cathode?