Compounds of metals with water. Metals: general characteristics of metals and alloys. IV. Displacement by more active metals of less active metals from solutions of their salts

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 ^.

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 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.

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

The mechanical characteristics include plasticity, elasticity, hardness, strength, viscosity.

The chemical properties of metals include corrosion resistance, solubility, and oxidizability.

Characteristics such as "fluidity", hardenability, weldability, ductility, 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 hue - a color. Among technical metals, only copper and alloys with it have color. The remaining elements are characterized by a shade from silver-white to gray-steel.
2. Fusibility. This characteristic indicates the ability of the element under the influence of temperature to pass into a liquid state from a solid. Fusibility is considered the most important property of metals. In the process of heating, all metals from a solid state pass into a liquid state. When the molten substance is cooled, a reverse transition occurs - from a liquid to a solid state.
3. Electrical conductivity. This characteristic indicates the ability to transfer electricity by free electrons. The electrical conductivity of metallic bodies is thousands of times greater than that of non-metallic ones. As the temperature increases, the conductivity of electricity decreases, and as the temperature decreases, accordingly, it increases. 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. The clearly magnetic (ferromagnetic) elements include only cobalt, nickel, iron, as well as a number of their alloys. However, in the process of heating to a certain temperature, these substances lose their magnetism. Individual iron alloys at room temperature are not ferromagnetic.
5. Thermal conductivity. This characteristic indicates the ability to transfer heat to a less heated one from a more heated body without visible movement of its constituent particles. The high level of thermal conductivity allows even and fast heating and cooling of metals. Among the technical elements, copper has the highest indicator.

Metals occupy a separate 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 an electrochemical or chemical relationship with the 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, platinum are called noble. Has high corrosion resistance Nickel and other non-ferrous metals are subject to destruction faster and more strongly than non-ferrous metals.
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 unlimited solubility in the liquid state can form solid solutions when solidified. In these solutions, atoms from one component are embedded in 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.

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 of these elements, a metallic bond is realized. 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 have characteristic features. These are 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 of oxides. When moving from fluorine to nitrogen (table 1), the oxidizing activity decreases and therefore an increasing number of 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 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

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
Mg2+	Mg 2+ +2 e - \u003d 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
Mn2+	Mn 2+ + 2e - = Mn	-1,180	Mn	Wed activity
Cr2+	Cr 2+ + 2e - = Cr	-0,913	Cr	Wed activity
H2O	2H 2 O+ 2e - \u003d H 2 + 2OH -	-0,826	H 2 , pH=14	Wed activity
Zn2+	Zn 2+ + 2e - = Zn	-0,763	Zn	Wed activity
Cr3+	Cr 3+ +3e - = Cr	-0,744	Cr	Wed activity
Fe2+	Fe 2+ + e - \u003d Fe	-0,440	Fe	Wed activity
H2O	2H 2 O + e - \u003d 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 - \u003d Co	-0,227	co	Wed activity
Ni2+	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
Fe3+	Fe 3+ + 3e - \u003d 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	Small active
Cu2+	Cu 2+ + 2e - = Cu	0,337	Cu	Small active
Cu+	Cu + + e - = Cu	0,521	Cu	Small active
Hg 2 2+	Hg 2 2+ + 2e - = Hg	0,788	Hg 2	Small active
Ag+	Ag + + e - = Ag	0,799	Ag	Small active
Hg2+	Hg 2+ + 2e - \u003d Hg	0,854	hg	Small active
Pt 2+	Pt 2+ + 2e - = Pt	1,2	Pt	Small active
Au 3+	Au 3+ + 3e - = Au	1,498	Au	Small active
Au +	Au++e-=Au	1,691	Au	Small active

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 voltages characterizes its ability to redox interactions in aqueous solutions under 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. The standard electrode potential of a hydrogen ion is 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

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 +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. The highest value of 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 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 dilute acid as an oxidizing agent is 0.96 V. Due to such a large value, 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 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 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.

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

The most active metals are located at the beginning of this row (on the left), the most inactive - at the end (on the right).
Reactions with simple substances. Metals react with non-metals 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 \u003d Na 2 O 2

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

2Mg + O 2 \u003d t 2MgO.

Inactive 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 \u003d 2LiOH + H 2;
Ba + 2H 2 O \u003d Ba (OH) 2 + H 2.

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

Mg + 2H 2 O \u003d 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 water vapor is passed over hot iron shavings, the following reaction occurs:

3Fe + 4H 2 O \u003d 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 \u003d CaCl 2 + H 2;
2Al + 3H 2 SO 4 \u003d Al 2 (SO 4) 3 + 3H 2.

Inactive metals are often practically insoluble in acids. This is due to the formation of an insoluble salt film on their surface. For example, lead, which is in the activity series up to hydrogen, practically does not dissolve in dilute sulfuric and hydrochloric acids due to the formation of a film of insoluble salts (PbSO 4 and PbCl 2) on its surface.

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