Interaction of metals with non-metals. Chemical properties of simple substances of metals and non-metals

Lecture 11. Chemical properties of metals.

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 of the periodic table from the diagonal drawn from boron to astatine. IN simple substances A metallic bond is realized in these elements. Metal atoms have few electrons in the outer electron shell, in the amount of 1, 2, or 3. Metals exhibit electropositive properties and have low electronegativity, less than two.

Metals are inherent characteristics. This solids, heavier than water, with a metallic sheen. 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, during heating, during rupture (exoelectronic emission).

The main feature 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. 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

The 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, initially 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 oxides. In the transition from fluorine to nitrogen (table 1), the oxidative activity decreases and therefore all more metals are not oxidized. For example, only lithium reacts with nitrogen and alkaline earth metals.

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

In aqueous solutions, the reducing activity of a metal is characterized by the value of its standard redox potential. From the entire range of standard redox potentials, a series of metal voltages is distinguished, which is indicated in table 2.

table 2

Row stress metals

In this series of voltages, the values ​​of the electrode potentials of the hydrogen electrode in acidic (рН=0), neutral (рН=7), alkaline (рН=14) media are also given. The position of a particular metal in a series of stresses characterizes its ability to redox interactions in aqueous solutions at standard conditions. Metal ions are oxidizing agents and metals are reducing agents. The further the metal is located in the series of voltages, the stronger the oxidizing agent in an aqueous solution are its ions. The closer the metal is to the beginning of the row, the stronger the reducing agent it is.

Metals are able to displace each other from salt solutions. The direction of the reaction is determined in this case by their mutual position in the series of voltages. It should be borne 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 series of voltages after magnesium.

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

Table 3

Conditional division of metals

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

φ \u003d -0.059 pH.

In a neutral environment (рН=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 much more negative than -0.41 V, displace hydrogen from water. But already magnesium displaces hydrogen only from hot water. Normally, 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 medium

Interaction of metals with hydrochloric acid.

The oxidizing agent in hydrochloric acid is the hydrogen ion. Standard electrode potential of the hydrogen ion zero. Therefore, all active metals and metals of intermediate activity must react with the acid. Only lead exhibits passivation.

Table 5

The interaction of metals with hydrochloric acid

Copper can be dissolved in very concentrated hydrochloric acid, despite the fact that it belongs to low-active metals.

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

Reaction of metals with dilute sulfuric acid. 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

Diluted sulphuric 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 +6 oxidation state 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 electrode potential in electrode processes with the participation of 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 takes place due to the formation of dense oxide films. Tin is oxidized to the tetravalent state with the formation of tin (IV) sulfate:

Sn + 4 H 2 SO 4 (conc.) \u003d Sn (SO 4) 2 + 2SO 2 + 2H 2 O.

Table 7

Interaction of metals with concentrated sulfuric acid

Lead oxidizes to the divalent state with the formation of soluble lead hydrosulfate. 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 \u003d ZnSO 4 + SO 2 + 2H 2 O;

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

4Zn + 5H 2 SO 4 \u003d 4ZnSO 4 \u003d 4ZnSO 4 + H 2 S + 4H 2 O.

Interaction of metals with dilute nitric acid.

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

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

Interaction of concentrated nitric acid with metals

Interaction of metals with alkali solutions.

Metals cannot be oxidized by alkalis. This is because 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, passive clean water metals interact vigorously with alkali solutions.

Table 10

Interaction of metals with alkali solutions

The dissolution process is presented in the form of two stages: the oxidation of the metal with water and the dissolution of the hydroxide:

Zn + 2HOH \u003d Zn (OH) 2 ↓ + H 2;

Zn (OH) 2 ↓ + 2NaOH \u003d Na 2.

Objective: practically get acquainted with the characteristic chemical properties of metals of various activity and their compounds; to study the features of metals with amphoteric properties. equate redox reactions by the method of electron-ion balance.

Theoretical part

Physical properties of metals. Under normal conditions, all metals, except mercury, are solids that differ sharply in 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. at chemical reactions they lose electrons and become positively charged ions. As a result, most metals react with typical oxidizing agents, such as oxygen, to form oxides, which in most cases cover the metal surface in a dense layer.

Mg°+O 2 °=2Mg +2 O- 2

Mg-2=Mg +2

ABOUT 2 +4 =2O -2

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

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

Intermediate activity metals Mg to H;

Inactive metals – from H to the end of the voltage series (to Au).

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

2K°+2N 2 O=2KOH+H 2 ABOUT

K°-=K + | 2

2H + +2 =H 2 0 | 1

The interaction of metals with acids

All anoxic acids (hydrochloric HCl, hydrobromic HBr, etc.), as well as some oxygen-containing acids (dilute sulfuric acid H 2 SO 4, phosphoric H 3 PO 4 , acetic CH 3 COOH, etc.) react with metals 1 and 2 groups standing in a series of voltages 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 ° | one

Concentrated sulfuric acid oxidizes metals of the 1st, 2nd and partially 3rd group (up to Ag inclusive) while being reduced to SO 2 - a colorless gas with a pungent odor, free sulfur that precipitates as a white precipitate or hydrogen sulfide H 2 S - a gas with a rotten smell eggs. The more active the metal, the more sulfur is reduced, for example:

| 1

| 8

Nitric acid of any concentration oxidizes almost all metals, while forming the 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 - odorless gas, NH 4 NO 3 - colorless solution). The more active the metal and the more dilute the acid, the more nitrogen is reduced in nitric acid.

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

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

Pb 0 -2 = Pb 2+ | 1

2H + +2 =H 2 ° | one

or when interacting with strong mortar alkalis:

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 salt and water), for example:

or in ionic form:

or in ionic form:

Practical part

Experience number 1.Interaction of metals with water .

Take a small piece of alkali or alkaline earth metal (sodium, potassium, lithium, calcium) that is stored in a jar of kerosene, dry it thoroughly with filter paper, and place it in 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 interacts with water, heat the reaction tube for some time on an alcohol lamp.

Experience number 2.Reaction of metals with dilute acids .

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

Experience number 3.Interaction of metals with concentrated acids .

Pour 20 - 25 drops of concentrated nitric and sulfuric (carefully!) Acids into two test tubes, drop metal into them, observe what is happening. If necessary, the test tubes can be heated on an alcohol lamp before the start of the reaction. Gently sniff the test tubes to determine the outgassing.

Experience number 4.The interaction of metals with alkalis .

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

An 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 hydrochloric acid solution to one part, and alkali solution to the other. Mark observations, write equations in molecular, full ionic and short ionic forms, draw a conclusion about the nature of the resulting hydroxide.

Formulation of work and conclusions

For redox reactions, write the equations of the electron-ion balance, write the ion-exchange reactions in molecular and ion-molecular forms.

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

Lab #11

Chemical properties of 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 new salt and 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.

If we draw a diagonal from beryllium to astatine in the periodic table of elements of D.I. Mendeleev, then there will be metal elements on the diagonal at the bottom left (they also include elements of secondary subgroups, highlighted in blue), and on the top right - non-metal elements (highlighted yellow). Elements located near the diagonal - semimetals or metalloids (B, Si, Ge, Sb, etc.) have a dual character (highlighted in pink).

As can be seen from the figure, the vast majority of elements are metals.

By their chemical nature, metals are chemical elements whose atoms donate electrons from the outer or pre-outer energy levels, thus forming positively charged ions.

Almost all metals have relatively large radii and a small number of electrons (from 1 to 3) at the external energy level. Metals are characterized by low electronegativity values ​​and reducing properties.

The most typical metals are located at the beginning of periods (starting from the second), further from left to right, the metallic properties weaken. In a group from top to bottom, metallic properties are enhanced, because the radius of atoms increases (due to an increase in the number of energy levels). This leads to a decrease in the electronegativity (the ability to attract electrons) of the elements and an increase in the reducing properties (the ability to donate electrons to other atoms in chemical reactions).

typical metals are s-elements (elements of the IA group from Li to Fr. elements of the PA group from Mg to Ra). General electronic formula their atoms ns 1-2. They are characterized by oxidation states + I and + II, respectively.

The small number of electrons (1-2) in the outer energy level of typical metal atoms suggests easy loss of these electrons and the manifestation of strong reducing properties, which reflect low electronegativity values. This implies the limited chemical properties and methods for obtaining typical metals.

A characteristic feature of typical metals is the tendency of their atoms to form cations and ionic chemical bonds with non-metal atoms. Compounds of typical metals with non-metals are ionic crystals“Metal cation non-metal anion”, for example K + Br -, Ca 2+ O 2-. Typical metal cations are also included in compounds with complex anions - hydroxides and salts, for example, Mg 2+ (OH -) 2, (Li +) 2CO 3 2-.

The A-group metals that form the amphoteric diagonal in the Be-Al-Ge-Sb-Po Periodic System, as well as the metals adjacent to them (Ga, In, Tl, Sn, Pb, Bi) do not exhibit typically metallic properties. The general electronic formula of their atoms ns 2 np 0-4 implies a greater variety of oxidation states, a greater ability to retain their own electrons, a gradual decrease in their reducing ability and the appearance of an oxidizing ability, especially in high oxidation states (typical examples are compounds Tl III, Pb IV, Bi v). Similar chemical behavior is also characteristic of most (d-elements, i.e., elements of B-groups Periodic system (typical examples- amphoteric elements Cr and Zn).

This manifestation of duality (amphoteric) properties, both metallic (basic) and non-metallic, is due to the nature of the chemical bond. In the solid state, compounds of atypical metals with non-metals contain predominantly covalent bonds (but less strong than bonds between non-metals). In solution, these bonds are easily broken, and the compounds dissociate into ions (completely or partially). For example, gallium metal consists of Ga 2 molecules, in the solid state aluminum and mercury (II) chlorides AlCl 3 and HgCl 2 contain strongly covalent bonds, but in a solution AlCl 3 dissociates almost completely, and HgCl 2 - to a very small extent (and even then into HgCl + and Cl - ions).

General physical properties of metals

Due to the presence of free electrons ("electron gas") in the crystal lattice, all metals exhibit the following characteristic general properties:

1) Plastic- the ability to easily change shape, stretch into a wire, roll into thin sheets.

2) metallic luster and opacity. This is due to the interaction of free electrons with light incident on the metal.

3) Electrical conductivity. It is explained by the directed movement of free electrons from the negative to the positive pole under the influence of a small potential difference. When heated, the electrical conductivity decreases, because. as the temperature rises, vibrations of atoms and ions in the nodes of the crystal lattice increase, which makes it difficult for the directed movement of the "electron gas".

4) Thermal conductivity. Due to the high mobility of free electrons, due to which fast leveling temperature by weight of the metal. The highest thermal conductivity is in bismuth and mercury.

5) Hardness. The hardest is chrome (cuts glass); the softest - alkali metals - potassium, sodium, rubidium and cesium - are cut with a knife.

6) Density. It is less the less atomic mass metal and larger atomic radius. The lightest is lithium (ρ=0.53 g/cm3); the heaviest is osmium (ρ=22.6 g/cm3). Metals having a density less than 5 g/cm3 are considered "light metals".

7) Melting and boiling points. The most fusible metal is mercury (m.p. = -39°C), the most refractory metal is tungsten (t°m. = 3390°C). Metals with t°pl. above 1000°C are considered refractory, below - low melting point.

General chemical properties of metals

Strong reducing agents: Me 0 – nē → Me n +

A number of stresses characterize the comparative activity of metals in redox reactions in aqueous solutions.

I. Reactions of metals with non-metals

1) With oxygen:
2Mg + O 2 → 2MgO

2) With sulfur:
Hg + S → HgS

3) With halogens:
Ni + Cl 2 – t° → NiCl 2

4) With nitrogen:
3Ca + N 2 – t° → Ca 3 N 2

5) With phosphorus:
3Ca + 2P – t° → Ca 3 P 2

6) With hydrogen (only alkali and alkaline earth metals react):
2Li + H 2 → 2LiH

Ca + H 2 → CaH 2

II. Reactions of metals with acids

1) Metals standing in the electrochemical series of voltages up to H reduce non-oxidizing acids to hydrogen:

Mg + 2HCl → MgCl 2 + H 2

2Al+ 6HCl → 2AlCl 3 + 3H 2

6Na + 2H 3 PO 4 → 2Na 3 PO 4 + 3H 2

2) With oxidizing acids:

In the interaction of nitric acid of any concentration and concentrated sulfuric acid with metals hydrogen is never released!

Zn + 2H 2 SO 4 (K) → ZnSO 4 + SO 2 + 2H 2 O

4Zn + 5H 2 SO 4(K) → 4ZnSO 4 + H 2 S + 4H 2 O

3Zn + 4H 2 SO 4(K) → 3ZnSO 4 + S + 4H 2 O

2H 2 SO 4 (c) + Cu → Cu SO 4 + SO 2 + 2H 2 O

10HNO 3 + 4Mg → 4Mg(NO 3) 2 + NH 4 NO 3 + 3H 2 O

4HNO 3 (c) + Сu → Сu (NO 3) 2 + 2NO 2 + 2H 2 O

III. Interaction of metals with water

1) Active (alkali and alkaline earth metals) form a soluble base (alkali) and hydrogen:

2Na + 2H 2 O → 2NaOH + H 2

Ca+ 2H 2 O → Ca(OH) 2 + H 2

2) Metals of medium activity are oxidized by water when heated to oxide:

Zn + H 2 O – t° → ZnO + H 2

3) Inactive (Au, Ag, Pt) - do not react.

IV. Displacement by more active metals of less active metals from solutions of their salts:

Cu + HgCl 2 → Hg + CuCl 2

Fe+ CuSO 4 → Cu+ FeSO 4

In industry, not pure metals are often used, but their mixtures - alloys in which the beneficial properties of one metal are complemented by the beneficial properties of another. So, copper has a low hardness and is of little use for the manufacture of machine parts, while alloys of copper with zinc ( brass) are already quite hard and are widely used in mechanical engineering. Aluminum has high ductility and sufficient lightness (low density), but is too soft. On its basis, an alloy with magnesium, copper and manganese is prepared - duralumin (duralumin), which, without losing useful properties aluminum, acquires high hardness and becomes suitable in the aircraft industry. Alloys of iron with carbon (and additions of other metals) are widely known cast iron And steel.

Metals in free form are reducing agents. but reactivity some metals is small due to the fact that they are covered surface oxide film, to varying degrees resistant to the action of such chemical reagents as water, solutions of acids and alkalis.

For example, lead is always covered with an oxide film; its transition into solution requires not only exposure to a reagent (for example, dilute nitric acid), but also heating. The oxide film on aluminum prevents its reaction with water, but is destroyed under the action of acids and alkalis. Loose oxide film (rust), formed on the surface of iron in moist air, does not interfere with the further oxidation of iron.

Under the influence concentrated acids are formed on metals sustainable oxide film. This phenomenon is called passivation. So, in concentrated sulfuric acid passivated (and then do not react with acid) such metals as Be, Bi, Co, Fe, Mg and Nb, and in concentrated nitric acid - metals A1, Be, Bi, Co, Cr, Fe, Nb, Ni, Pb , Th and U.

When interacting with oxidizing agents in acidic solutions, most metals turn into cations, the charge of which is determined by the stable oxidation state of a given element in compounds (Na +, Ca 2+, A1 3+, Fe 2+ and Fe 3+)

The reducing activity of metals in an acidic solution is transmitted by a series of stresses. Most metals are converted into a solution with hydrochloric and dilute sulfuric acids, but Cu, Ag and Hg - only with sulfuric (concentrated) and nitric acids, and Pt and Au - with "aqua regia".

Corrosion of metals

An undesirable chemical property of metals is their, i.e., active destruction (oxidation) 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, as a result of which rust is formed, and the products crumble into powder.

Corrosion of metals proceeds in water also due to the presence of dissolved CO 2 and SO 2 gases; an acidic environment is created, and H + cations are displaced by active metals in the form of hydrogen H 2 ( hydrogen corrosion).

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

It is because of this that the tinned surface rusts. cans(tin-plated iron) when stored in a humid atmosphere and handled carelessly (iron quickly breaks down after even a small scratch is introduced, allowing iron to come into contact 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 is enhanced when it is coated with a more active metal or when they are fused; for example, coating iron with chromium or making an alloy of iron with chromium eliminates the corrosion of iron. Chrome-plated iron and steel containing chromium ( stainless steel ) have high corrosion resistance.

electrometallurgy, i.e., obtaining metals by electrolysis of melts (for the most active metals) or salt solutions;

pyrometallurgy, i.e., the recovery of metals from ores at high temperature (for example, the production of iron in the 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 CuSO 4 solution by the action of zinc, iron or aluminum).

Native metals are sometimes found in nature (typical examples are Ag, Au, Pt, Hg), but more often metals are in the form of compounds ( metal ores). In terms of prevalence in earth's crust metals are different: from the most common - Al, Na, Ca, Fe, Mg, K, Ti) to the rarest - Bi, In, Ag, Au, Pt, Re.

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

With a large variety, all chemical reactions of metals are redox and can only be of two types: compounds and substitutions. Metals are capable of donating electrons during chemical reactions, that is, they can be reducing agents, and show only a positive oxidation state in the compounds formed.

IN general view this can be expressed in a diagram:
Me 0 - ne → Me + n,
where Me - metal - a simple substance, and Me 0 + n - metal chemical element in connection.

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

At high temperatures most metals burn in oxygen:

2Mg + O 2 \u003d 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 sometimes formed. Alkali metals, as well as calcium, strontium, barium, interact very actively with water under normal conditions. 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 the metal reacts with an acid, then it is part of the resulting salt. When a metal interacts with acid solutions, it can be oxidized by the hydrogen ions present in that solution. The abbreviated ionic equation in general form can be written as follows:

Me + nH + → Me n + + H 2

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

When metals interact with salts, a substitution reaction occurs: electrons from the atoms of the substituting - more active metal pass to the ions of the substituting - 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 their salts, metals are located in the electrochemical series of voltages (activity) of metals:

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

The metals located to the left of this row are more active and are able to displace the metals following them from salt solutions.

Hydrogen is included in the electrochemical series of voltages of metals, as the only non-metal that separates from metals common property- 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 \u003d ZnCl 2 + H 2 + Q

Metals standing in the electrochemical series of voltages up to hydrogen displace it from solutions of many acids (hydrochloric, sulfuric, etc.), and all following it, for example, do not displace copper.

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