If in periodic table elements of D.I. Mendeleev, draw a diagonal from beryllium to astatine, then at the bottom left along the diagonal there will be metal elements (these also include elements of secondary subgroups, highlighted in blue), and on the top right - nonmetal elements (highlighted in 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.

According to its chemical nature Metals are chemical elements whose atoms donate electrons from an external or pre-external energy level, thus forming positively charged ions.

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

The most typical metals are located at the beginning of the 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 the 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 IA-group from Li to Fr. elements of 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) at the external energy level of typical metal atoms suggests a slight loss of these electrons and the manifestation of strong reducing properties, which reflect low values ​​of electronegativity. Hence, the chemical properties and methods of obtaining typical metals are limited.

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

Metals of A-groups forming a diagonal of amphotericity in Periodic table Be-Al-Ge-Sb-Po, as well as adjacent metals (Ga, In, Tl, Sn, Pb, Bi) do not exhibit typically metallic properties. The general electronic formula of their atoms ns 2 np 0-4 assumes a greater variety of oxidation states, a greater ability to hold their own electrons, a gradual decrease in their reductive ability and the appearance of oxidizing ability, especially in high oxidation states (typical examples are compounds Tl III, Pb IV, Bi v). A similar chemical behavior is typical for most (d-elements, i.e., elements of B-groups of the Periodic Table (typical examples are amphoteric elements Cr and Zn).

This manifestation of the duality (amphotericity) of properties, both metallic (basic) and non-metallic, is due to the nature 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 (in whole or in part). 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 of AlCl 3 it dissociates almost completely, and HgCl 2 - to a very small extent (and then on ions НgСl + and Сl -).

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, be drawn into wire, rolled 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 directional 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 with an increase in temperature, the vibrations of atoms and ions in the nodes of the crystal lattice intensify, which complicates the directional movement of the "electron gas".

4) Thermal conductivity. It is caused by the high mobility of free electrons, due to which there is a rapid equalization of temperature over the mass of the metal. Bismuth and mercury have the highest thermal conductivity.

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 the smaller, the less the atomic mass of the metal and the greater the radius of the atom. The lightest is lithium (ρ = 0.53 g / cm3); the heaviest is osmium (ρ = 22.6 g / cm3). Metals with a density of less than 5 g / cm3 are considered "light metals".

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

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 gray:
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 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 (К) → ZnSO 4 + SO 2 + 2H 2 O

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

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

2H 2 SO 4 (k) + 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) + Cu → Cu (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 of less active metals from solutions of their salts by more active metals:

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 copper-zinc alloys ( brass) are already quite solid and are widely used in mechanical engineering. Aluminum has high ductility and sufficient lightness (low density), but too soft. On its basis, an alloy with magnesium, copper and manganese is prepared - duralumin (duralumin), which, without losing the useful properties of aluminum, acquires high hardness and becomes suitable in aircraft construction. Alloys of iron with carbon (and additives of other metals) are widely known cast iron and steel.

Free metals are reducing agents. However, the reactivity of some metals is low due to the fact that they are coated surface oxide film, to varying degrees of resistance to the action of chemicals such as water, solutions of acids and alkalis.

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

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

When interacting with oxidants in acidic solutions, most metals are converted 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 acidic solution transmitted by a number of voltages. Most of the metals are converted into a solution with hydrochloric and dilute sulfuric acids, but Cu, Ag and Hg - only sulfuric (concentrated) and nitric acids, and Pt and Au - "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, corrosion of iron products in water is widely known, as a result of which rust is formed and the products are crumbled into powder.

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

The place of contact of two dissimilar metals ( contact corrosion). A galvanic pair arises between one metal, such as Fe, and another metal, such as Sn or Cu, placed in water. The flow of electrons goes from the more active metal, which is to the left in the series of voltages (Fe), to the less active metal (Sn, Cu), and the more active metal is destroyed (corroded).

It is because of this that the tinned surface of tin cans (iron coated with tin) rusts when stored in a humid atmosphere and carelessly handling them (iron quickly collapses after the appearance of even a small scratch that allows 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 in the presence of 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; thus, plating iron with chromium or making an iron-chromium alloy eliminates the corrosion of iron. Chromium-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 reduction of metals from ores at high temperatures (for example, the production of iron in a blast furnace);

hydrometallurgy, i.e., the separation of metals from solutions of their salts with more active metals (for example, obtaining 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). By prevalence in earth crust metals are different: from the most common - Al, Na, Ca, Fe, Mg, K, Ti) to the rarest ones - Bi, In, Ag, Au, Pt, Re.

Inorganic substances are simple and complex. Simple substances are divided into metals (K, Na, Li) and non-metals (O, Cl, P). Complex substances are divided into oxides, hydroxides (bases), salts and acids.

Oxides

Oxides- compounds of a chemical element (metal or non-metal) with oxygen (oxidation state -2), while oxygen is associated with a less electronegative element.

Allocate:

1. Acid oxides- oxides showing acidic properties. Formed by non-metals and oxygen. Examples: SO3, SO2, CO2, P2O5, N2O5.

2. Amphoteric oxides- oxides that can exhibit both basic and acidic properties (this property is called amphotericity). Examples: Al2O3, CrO3, ZnO, BeO, PbO.

3. Basic oxides- metal oxides, while the metals exhibit an oxidation state of +1 or +2. Examples: K2O, MgO, CaO, BaO, Li2O, Na2O.

4. Non-salt-forming oxides- practically do not enter into reactions, do not have corresponding acids and hydroxides. Examples: CO, NO.

Chemical properties basic oxides

1. Interaction with water

Only oxides of alkali and alkaline earth metals enter into the reaction, the hydroxides of which form a soluble base

basic oxide + water → alkali

K2O + H2O → 2KOH

CaO + H2O → Ca (OH) 2

2. Interaction with acid

basic oxide + acid → salt + water

MgO + H2SO4 → MgSO4 + H2O

Na2O + H2S (g) → 2NaHS + H2O

MgO (g) + HCl → Mg (OH) Cl

3. Interaction with acidic or amphoteric oxides

basic oxide + acidic / amphoteric oxide → salt

In this case, the metal in the basic oxide becomes a cation, and the acidic / amphoteric oxide becomes an anion (acidic residue). Reactions between solid oxides occur when heated. Water-insoluble basic oxides do not interact with gaseous acidic oxides.

BaO + SiO2 (t) → BaSiO3

K2O + ZnO (t) → K2ZnO2

FeO + CO2 ≠

4. Interaction with amphoteric hydroxides

basic oxide + amphoteric hydroxide → salt + water

Na2O + 2Al (OH) 3 (t) → 2NaAlO2 + 3H2O

5. Decomposition at a temperature of oxides of noble metals and mercury

2Ag2O (t) → 4Ag + O2

2HgO (t) → 2Hg + O2

6. Interaction with carbon (C) or hydrogen (H2) at high temperature.

When the oxides of alkali, alkaline earth metals and aluminum are reduced in this way, not the metal itself is released, but its carbide.

FeO + C (t) → Fe + CO

3Fe2O3 + C (t) → 2Fe3O4 + CO

CaO + 3C (t) → CaC2 + CO

CaO + 2H2 (t) → CaH2 + H2O

7. Active metals reduce less active metals from their oxides at high temperatures

CuO + Zn (t) → ZnO + Cu

8. Oxygen oxidizes lower oxides to higher ones.

Alkali and alkaline earth metal oxides are converted to peroxides

4FeO + O2 (t) → 2Fe2O3

2BaO + O2 (t) → 2BaO2

2NaO + O2 (t) → 2Na2O2

Chemical properties of acidic oxides

1. Interaction with water

acid oxide + water → acid

SO3 + H2O → H2SO4

SiO2 + H2O ≠

Some oxides do not have corresponding acids, in which case a disproportionation reaction occurs

2NO2 + H2O → HNO3 + HNO2

3NO2 + H2O (t) → 2HNO3 + NO

2ClO2 + H2O → HClO3 + HClO2

6ClO2 + 3H2O (t) → 5HClO3 + HCl

Depending on the number of water molecules attached to P2O5, three different acids are formed - metaphosphoric НРО3, pyrophosphoric Н4Р2О7 or orthophosphoric Н3РО4.

P2O5 + H2O → 2HPO3

P2O5 + 2H2O → H4P2O7

P2O5 + 3H2O → 2H3PO4

Chromium oxide corresponds to two acids - chromic H2CrO4 and dichromic H2Cr2O7 (III)

CrO3 + H2O → H2CrO4

2CrO3 + H2O → H2Cr2O7

2. Interaction with bases

acid oxide + base → salt + water

Insoluble acid oxides react only when fusion, and soluble - under normal conditions.

SiO2 + 2NaOH (t) → Na2SiO3 + H2O

With an excess of oxide, an acidic salt is formed.

CO2 (g) + NaOH → NaHCO3

P2O5 (g) + 2Ca (OH) 2 → 2CaHPO4 + H2O

P2O5 (g) + Ca (OH) 2 + H2O → Ca (H2PO4) 2

With an excess of base, a basic salt is formed

CO2 + 2Mg (OH) 2 (g) → (MgOH) 2CO3 + H2O

Oxides that do not have corresponding acids undergo disproportionation reactions and form two salts.

2NO2 + 2NaOH → NaNO3 + NaNO2 + H2O

2ClO2 + 2NaOH → NaClO3 + NaClO2 + H2O

CO2 reacts with some amphoteric hydroxides (Be (OH) 2, Zn (OH) 2, Pb (OH) 2, Cu (OH) 2) to form basic salt and water.

CO2 + 2Be (OH) 2 → (BeOH) 2CO3 ↓ + H2O

CO2 + 2Cu (OH) 2 → (CuOH) 2CO3 ↓ + H2O

3. Interaction with basic or amphoteric oxide

acidic oxide + basic / amphoteric oxide → salt

Reactions between solid oxides take place during fusion. Amphoteric and water-insoluble basic oxides interact only with solid and liquid acidic oxides.

SiO2 + BaO (t) → BaSiO3

3SO3 + Al2O3 (t) → Al2 (SO4) 3

4. Interaction with salt

acidic non-volatile oxide + salt (t) → salt + acidic volatile oxide

SiO2 + CaCO3 (t) → CaSiO3 + CO2

P2O5 + Na2CO3 → 2Na3PO4 + 2CO2

5. Acidic oxides do not interact with acids, but P2O5 reacts with anhydrous oxygen-containing acids.

This forms HPO3 and the corresponding acid anhydride

P2O5 + 2HClO4 (anhydrous) → Cl2O7 + 2HPO3

P2O5 + 2HNO3 (anhydrous) → N2O5 + 2HPO3

6. Enter into redox reactions.

1. Recovery

At high temperatures, some non-metals can reduce oxides.

CO2 + C (t) → 2CO

SO3 + C → SO2 + CO

H2O + C (t) → H2 + CO

Magnesium thermal is often used to reduce non-metals from their oxides.

CO2 + 2Mg → C + 2MgO

SiO2 + 2Mg (t) → Si + 2MgO

N2O + Mg (t) → N2 + MgO

2. Lower oxides are converted into higher ones when interacting with ozone (or oxygen) at high temperatures in the presence of a catalyst

NO + O3 → NO2 + O2

SO2 + O3 → SO3 + O2

2NO2 + O3 → N2O5 + O2

2CO + O2 (t) → 2CO2

2SO2 + O2 (t, kat) → 2SO3

P2O3 + O2 (t) → P2O5

2NO + O2 (t) → 2NO2

2N2O3 + O2 (t) → 2N2O4

3. Oxides also enter into other redox reactions

SO2 + NO2 → NO + SO3 4NO2 + O2 + 2H2O → 4HNO3

2SO2 + 2NO → N2 + 2SO3 2N2O5 → 4NO2 + O2

SO2 + 2H2S → 3S ↓ + 2H2O 2NO2 (t) → 2NO + O2

2SO2 + O2 + 2H2O → 2H2SO4 3N2O + 2NH3 → 4N2 + 3H2O

2CO2 + 2Na2O2 → 2Na2CO3 + O2 10NO2 + 8P → 5N2 + 4P2O5

N2O + 2Cu (t) → N2 + Cu2O

2NO + 4Cu (t) → N2 + 2Cu2O

N2O3 + 3Cu (t) → N2 + 3CuO

2NO2 + 4Cu (t) → N2 + 4CuO

N2O5 + 5Cu (t) → N2 + 5CuO

Chemical properties of amphoteric oxides

1. Do not interact with water

amphoteric oxide + water ≠

2. Interaction with acids

amphoteric oxide + acid → salt + water

Al2O3 + 3H2SO4 → Al2 (SO4) 3 + 3H2O

With an excess of polybasic acid, an acid salt is formed

Al2O3 + 6H3PO4 (g) → 2Al (H2PO4) 3 + 3H2O

With an excess of oxide, a basic salt is formed

ZnO (g) + HCl → Zn (OH) Cl

Double oxides form two salts

Fe3O4 + 8HCl → FeCl2 + 2FeCl3 + 4H2O

3. Interaction with acidic oxide

amphoteric oxide + acidic oxide → salt

Al2O3 + 3SO3 → Al2 (SO4) 3

4. Interaction with alkali

amphoteric oxide + alkali → salt + water

During fusion, medium salt and water are formed, and in solution - complex salt

ZnO + 2NaOH (tv) (t) → Na2ZnO2 + H2O

ZnO + 2NaOH + H2O → Na2

5. Interaction with basic oxide

amphoteric oxide + basic oxide (t) → salt

ZnO + K2O (t) → K2ZnO2

6. Interaction with salts

amphoteric oxide + salt (t) → salt + volatile acidic oxide

Amphoteric oxides displace volatile acid oxides from their salts during fusion

Al2O3 + K2CO3 (t) → KAlO2 + CO2

Fe2O3 + Na2CO3 (t) → 2NaFeO2 + CO2

Chemical properties of bases

Bases are substances that include a metal cation and a hydroxide anion. Bases are soluble (alkalis - NaOH, KOH, Ba (OH) 2) and insoluble (Al2O3, Mg (OH) 2).

1. Soluble base + indicator → color change

When the indicator is added to the base solution, its color changes:

Colorless phenolphthalein - raspberry

Purple litmus - blue

Methyl orange - yellow

2. Interaction with acid (neutralization reaction)

base + acid → salt + water

By the reaction, medium, acidic or basic salts can be obtained. With an excess of a polyacidic acid, an acidic salt is formed, with an excess of a polyacidic base, a basic salt is formed.

Mg (OH) 2 + H2SO4 → MGSO4 + 2H2O

Mg (OH) 2 + 2H2SO4 → MG (HSO4) 2 + 2H2O

2Mg (OH) 2 + H2SO4 → (MgOH) 2SO4 + 2H2O

3. Interaction with acidic oxides

base + acid oxide → salt + water

6NH4OH + P2O5 → 2 (NH4) 3PO4 + 3H2O

4. Interaction of alkali with amphoteric hydroxide

alkali + amphoteric hydroxide → salt + water

In this reaction, amphoteric hydroxide exhibits acidic properties. During the reaction in the melt, average salt and water are obtained, and in solution, a complex salt is obtained. Iron (III) and chromium (III) hydroxides dissolve only in concentrated alkali solutions.

2KOH (tv) + Zn (OH) 2 (t) → K2ZnO2 + 2H2O

KOH + Al (OH) 3 → K

3NaOH (conc) + Fe (OH) 3 → Na3

5. Interaction with amphoteric oxide

alkali + amphoteric oxide → salt + water

2NaOH (s) + Al2O3 (t) → 2NaAlO2 + H2O

6NaOH + Al2O3 + 3H2O → 2Na3

6. Interaction with salt

An ion exchange reaction occurs between the base and the salt. It occurs only during the precipitation of a precipitate or during the evolution of gas (with the formation of NH4OH).

A. Reaction between soluble base and soluble acid salt

soluble base + soluble acid salt → medium salt + water

If the salt and base are formed by different cations, then two middle salts are formed. In the case of acidic ammonium salts, excess alkali leads to the formation of ammonium hydroxide.

Ba (OH) 2 + Ba (HCO3) 2 → 2BaCO3 ↓ + 2H2O

2NaOH (g) + NH4HS → Na2S + NH4OH + H2O

B. Reaction of a soluble base with a soluble medium or basic salt.

Several scenarios are possible

soluble base + soluble medium / basic salt → insoluble salt ↓ + base

→ salt + insoluble base ↓

→ salt + weak electrolyte NH4OH

→ reaction does not go

Reactions occur between soluble bases and a medium salt only if the result is an insoluble salt, or an insoluble base, or a weak electrolyte NH4OH

NaOH + KCl ≠ the reaction does not go

If the initial salt is formed by a multi-acid base, with a lack of alkali, a basic salt is formed

Under the action of alkalis on silver and mercury (II) salts, not their hydroxides are released, which dissolve at 25C, but insoluble oxides Ag2O and HgO.

7. Decomposition at temperature

basic hydroxide (t) → oxide + water

Ca (OH) 2 (t) → CaO + H2O

NaOH (t) ≠

Some bases (AgOH, Hg (OH) 2 and NH4OH) decompose even at room temperature

LiOH (t) → Li2O + H2O

NH4OH (25C) → NH3 + H2O

8. Interaction of alkali and transition metal

alkali + transition metal → salt + H2

2Al + 2KOH + 6H2O → 2K + 3H2

Zn + 2NaOH (s) (t) → Na2ZnO2 + H2

Zn + 2NaOH + 2H2O → Na2 + H2

9. Interaction with non-metals

Alkalis interact with some non-metals - Si, S, P, F2, Cl2, Br2, I2. In this case, two salts are often formed as a result of disproportionation.

Si + 2KOH + H2O → K2SiO3 + 2H2

3S + 6KOH (t) → 2K2S + K2SO3 + 3H2O

Cl2 + 2KOH (conc) → KCl + KClO + H2O (for Br, I)

3Cl2 + 6KOH (conc) (t) → 5KCl + KClO3 + 3H2O (for Br, I)

Cl2 + Ca (OH) 2 → CaOCl2 + H2O

4F2 + 6NaOH (decomp) → 6NaF + OF2 + O2 + 3H2O

4P + 3NaOH + 3H2O → 3NaH2PO2 + PH3

Hydroxides with reducing properties can be oxidized by oxygen

4Fe (OH) 2 + O2 + 2H2O → 4Fe (OH) 3 (= Cr)

Chemical properties of acids

1. Change the color of the indicator

soluble acid + indicator → color change

Violet litmus and methyl orange turn red, phenolphthalein becomes transparent

2. Interaction with bases (neutralization reaction)

acid + base → salt + water

H2SO4 + Mg (OH) 2 → MgSO4 + 2H2O

3. Interaction with basic oxide

acid + basic oxide → salt + water

2HCl + CuO → CuCl2 + H2O

4. Interaction with amphoteric hydroxides with the formation of medium, acidic or basic salts

acid + amphoteric hydroxide → salt + water

2HCl + Be (OH) 2 → BeCl2 + 2H2O

H3PO4 () + Zn (OH) 2 → ZNHPO4 + 2H2O

HCl + Al (OH) 3 () → Al (OH) 2Cl + H2O

5. Interaction with amphoteric oxides

acid + amphoteric oxide → salt + water

H2SO4 + ZnO → ZnSO4 + H2O

6. Interaction with salts

General reaction scheme: acid + salt → salt + acid

An ion exchange reaction takes place, which goes to the end only in the case of gas formation or precipitation.

For example: HCl + AgNO3 → AgCl ↓ + HNO3

2HBr + K2SiO3 → 2KBr + H2SiO3 ↓

A. Reaction with a salt of a more volatile or weaker acid to form a gas

HCl + NaHS → NaCl + H2S

B. Reaction between a strong acid and a salt of a strong or moderate acid to form an insoluble salt

strong acid + strong / medium acid salt → insoluble salt + acid

Non-volatile phosphoric acid displaces strong, but volatile hydrochloric and nitric acids from their salts, subject to the formation of an insoluble salt

B. Interaction of an acid with a basic salt of the same acid

acid1 + basic acid salt1 → medium salt + water

HCl + Mg (OH) Cl → MgCl2 + H2O

D. The interaction of a polybasic acid with a medium or acidic salt of the same acid with the formation of an acidic salt of the same acid containing more hydrogen atoms

polybasic acid1 + medium / acidic acid salt1 → acidic acid salt1

H3PO4 + Ca3 (PO4) 2 → 3CaHPO4

H3PO4 + CaHPO4 → Ca (H2PO4) 2

E. Interaction of hydrogen sulfide acid with salts of Ag, Cu, Pb, Cd, Hg with the formation of insoluble sulfide

acid H2S + salt Ag, Cu, Pb, Cd, Hg → Ag2S / CuS / PbS / CdS / HgS ↓ + acid

H2S + CuSO4 → CuS ↓ + H2SO4

E. Reaction of an acid with a medium or complex salt with an amphoteric metal in the anion

a) in the case of a lack of acid, a medium salt and amphoteric hydroxide are formed

acid + medium / complex salt in amphoteric metal in anion → medium salt + amphoteric hydroxide

b) in the case of an excess of acid, two average salts and water are formed

acid + medium / complex salt with amphoteric metal in the anion → medium salt + medium salt + water

G. In some cases, acids with salts enter into redox reactions or complexation reactions:

H2SO4 (conc) and I‾ / Br‾ (products H2S and I2 / SO2 and Br2)

H2SO4 (conc) and Fe² + (products SO2 and Fe³ +)

HNO3 diluted / conc and Fe² + (products NO / NO2 and Fe³ +)

HNO3 open / conc and SO3²‾ / S²‾ (products NO / NO2 and SO4²‾ / S or SO4²‾)

HClconc and KMnO4 / K2Cr2O7 / KClO3 (products Cl2 and Mn² + / Cr² + / Cl‾)

3. Interaction of concentrated sulfuric acid with solid salt

Non-volatile acids can displace volatile acids from their solid salts

7. Interaction of acid with metal

A. Interaction of acid with metals in a row before or after hydrogen

acid + metal up to Н2 → silt metal in the minimum oxidation state + Н2

Fe + H2SO4 (diluted) → FeSO4 + H2

acid + metal after H2 ≠ the reaction does not go

Cu + H2SO4 (decomp) ≠

B. Interaction of concentrated sulfuric acid with metals

H2SO4 (conc) + Au, Pt, Ir, Rh, Ta ≠ the reaction does not proceed

H2SO4 (conc) + alkali / alkaline earth metal and Mg / Zn → H2S / S / SO2 (depending on conditions) + metal sulfate in the maximum oxidation state + H2O

Zn + 2H2SO4 (conc) (t1) → ZnSO4 + SO2 + 2H2O

3Zn + 4H2SO4 (end) (t2> t1) → 3ZnSO4 + S ↓ + 4H2O

4Zn + 5H2SO4 (end) (t3> t2) → 4ZnSO4 + H2S + 4H2O

H2SO4 (conc) + other metals → SO2 + metal sulfate in the maximum oxidation state + H2O

Cu + 2H2SO4 (conc) (t) → CuSO4 + SO2 + 2H2O

2Al + 6H2SO4 (conc) (t) → Al2 (SO4) 3 + 3SO2 + 6H2O

B. Reaction of concentrated nitric acid with metals

HNO3 (conc) + Au, Pt, Ir, Rh, Ta, Os ≠ the reaction does not proceed

HNO3 (conc) + Pt ≠

HNO3 (conc) + alkali / alkaline earth metal → N2O + metal nitrate in the maximum oxidation state + H2O

4Ba + 10HNO3 (conc) → 4Ba (NO3) 2 + N2O + 5H2O

HNO3 (conc) + other metals at temperature → NO2 + metal nitrate in the maximum oxidation state + H2O

Ag + 2HNO3 (conc) → AgNO3 + NO2 + H2O

It interacts with Fe, Co, Ni, Cr and Al only when heated, since under normal conditions these metals are passivated with nitric acid - they become chemically resistant

D. Reaction of dilute nitric acid with metals

HNO3 (decomposition) + Au, Pt, Ir, Rh, Ta ≠ the reaction does not take place

Very passive metals (Au, Pt) can be dissolved in aqua regia - a mixture of one volume of concentrated nitric acid with three volumes of concentrated hydrochloric acid. The oxidizing agent in it is atomic chlorine, which is split off from nitrosyl chloride, which is formed as a result of the reaction: HNO3 + 3HCl → 2H2O + NOCl + Cl2

HNO3 (decomp) + alkaline / alkaline earth metal → NH3 (NH4NO3) + metal nitrate in maximum oxidation state + H2O

NH3 is converted to NH4NO3 in excess of nitric acid

4Ca + 10HNO3 (diluted) → 4Ca (NO3) 2 + NH4NO3 + 3H2O

HNO3 (broken) + metal in the series of stresses up to Н2 → NO / N2O / N2 / NH3 (depending on conditions) + metal nitrate in the maximum oxidation state + Н2О

With the rest of the metals, which stand in the series of voltages up to hydrogen and non-metals, HNO3 (diluted) forms salt, water and, mainly, NO, but, depending on the conditions, both N2O, and N2, and NH3 / NH4NO3 (the more diluted the acid , the lower the oxidation state of nitrogen in the emitted gaseous product)

3Zn + 8HNO3 (decomp) → 3Zn (NO3) 2 + 2NO + 4H2O

4Zn + 10HNO3 (decomp) → 4Zn (NO3) 2 + N2O + 5H2O

5Zn + 12HNO3 (decomp) → 5Zn (NO3) 2 + N2 + 6H2O

4Zn + 10HNO3 (fine parsed) → 4Zn (NO3) 2 + NH4NO3 + 3H2O

HNO3 (diluted) + metal after Н2 → NO + metal nitrate in the maximum oxidation state + H2O

With low-activity metals, standing after H2, HNO3 dissociates forms salt, water and NO

3Cu + 8HNO3 (decomp) → 3Cu (NO3) 2 + 2NO + 4H2O

8. Decomposition of acids at temperature

acid (t) → oxide + water

H2CO3 (t) → CO2 + H2O

H2SO3 (t) → SO2 + H2O

H2SiO3 (t) → SiO2 + H2O

2H3PO4 (t) → H4P2O7 + H2O

H4P2O7 (t) → 2HPO3 + H2O

4HNO3 (t) → 4NO2 + O2 + 2H2O

3HNO2 (t) → HNO3 + 2NO + H2O

2HNO2 (t) → NO2 + NO + H2O

3HCl (t) → 2HCl + HClO3

4H3PO3 (t) → 3H3PO4 + PH3

9. Interaction of acid with non-metals (redox reaction). In this case, the non-metal is oxidized to the corresponding acid, and the acid is reduced to a gaseous oxide: H2SO4 (conc) - to SO2; HNO3 (conc) - up to NO2; HNO3 (diluted) - to NO.

S + 2HNO3 (decomp) → H2SO4 + 2NO

S + 6HNO3 (conc) → H2SO4 + 6NO2 + 2H2O

S + 2H2SO4 (conc) → 3SO2 + CO2 + 2H2O

C + 2H2SO4 (conc) → 2SO2 + CO2 + 2H2O

C + 4HNO3 (conc) → 4NO2 + CO2 + 2H2O

P + 5HNO3 (decomp) + 2H2O → 3H3PO4 + 5NO

P + 5HNO3 (conc) → HPO3 + 5NO2 + 2H2O

H2S + G2 → 2HG + S ↓ (except F2)

H2SO3 + G2 + H2O → 2HG + H2SO4 (except F2)

2H2S (aq) + O2 → 2H2O + 2S ↓

2H2S + 3O2 → 2H2O + 2SO2 (combustion)

2H2S + O2 (short) → 2H2O + 2S ↓

More active halogens displace less active ones from NG acids (exception: F2 reacts with water, not acid)

2HBr + Cl2 → 2HCl + Br2 ↓

2HI + Cl2 → 2HCl + I2 ↓

2HI + Br2 → 2HBr + I2 ↓

10. Redox reactions between acids

H2SO4 (conc) 2HBr → Br2 ↓ + SO2 + 2H2O

H2SO4 (conc) + 8HI → 4I2 ↓ + H2S + 4H2O

H2SO4 (conc) + HCl ≠

H2SO4 (conc) + H2S → S ↓ + SO2 + 2H2O

3H2SO4 (conc) + H2S → 4SO2 + 4H2O

H2SO3 + 2H2S → 3S ↓ + 3H2O

2HNO3 (conc) + H2S → S ↓ + 2NO2 + 2H2O

2HNO3 (conc) + SO2 → H2SO4 + 2NO2

6HNO3 (conc) + HI → HIO3 + 6NO2 + 3H2O

2HNO3 (conc) + 6HCl → 3Cl2 + 2NO + 4H2O

Chemical properties of amphoteric hydroxides

1. Interaction with basic oxide

amphoteric hydroxide + basic oxide → salt + water

2Al (OH) 3 + Na2O (t) → 2NaAlO2 + 3H2O

2. Interaction with amphoteric or acidic oxide

amphoteric hydroxide + amphoteric / acidic oxide ≠ no reaction

Some amphoteric oxides (Be (OH) 2, Zn (OH) 2, Pb (OH) 2) react with acidic oxide CO2 to form precipitates of basic salts and water

2Be (OH) 2 + CO2 → (BeOH) 2CO3 ↓ + H2O

3. Interaction with alkali

amphoteric hydroxide + alkali → salt + water

Zn (OH) 2 + 2KOH (s) (t) → K2ZnO2 + 2H2O

Zn (OH) 2 + 2KOH → K2

4. Does not interact with insoluble bases or amphoteric hydroxides

amphoteric hydroxide + insoluble base / amphoteric hydroxide ≠ no reaction

5. Interaction with acids

amphoteric hydroxide + acid → salt + water

Al (OH) 3 + 3HCl → AlCl3 + 3H2O

6. Do not react with salts

amphoteric hydroxide + salt ≠ the reaction does not go

7. Do not react with metals / non-metals (simple substances)

amphoteric hydroxide + metal / non-metal ≠ no reaction

8. Thermal decomposition

amphoteric hydroxide (t) → amphoteric oxide + water

2Al (OH) 3 (t) → Al2O3 + 3H2O

Zn (OH) 2 (t) → ZnO + H2O

General information about salts

Let's imagine that we have an acid and an alkali, we carry out a neutralization reaction between them and get an acid and a salt.

NaOH + HCl → NaCl (sodium chloride) + H2O

It turns out that the salt consists of a metal cation and an acid residue anion.

Salts are:

1. Acidic (with one or two hydrogen cations (that is, they have an acidic (or slightly acidic) environment) - KHCO3, NaHSO3).

2. Medium (I have a metal cation and an anion of an acid residue, the medium must be determined using a pH meter - BaSO4, AgNO3).

3. Basic (have a hydroxide ion, that is, an alkaline (or weakly alkaline) medium - Cu (OH) Cl, Ca (OH) Br).

There are also double salts that form cations of two metals (K) upon dissociation.

Salts, with a few exceptions, are crystalline solids with high melting points. Most salts white(KNO3, NaCl, BaSO4, etc.). Some salts are colored (K2Cr2O7 - orange, K2CrO4 - yellow, NiSO4 - green, CoCl3 - pink, CuS - black). According to their solubility, they can be divided into soluble, slightly soluble and practically insoluble. Acidic salts are generally more soluble in water than the corresponding average, and basic ones are worse.

Chemical properties of salts

1. Salt + water

When many salts are dissolved in water, their partial or complete decomposition occurs - hydrolysis... Some salts form crystalline hydrates. When medium salts containing an amphoteric metal in the anion are dissolved in water, complex salts are formed.

NaCl + H2O → NaOH + HCl

Na2ZnO2 + 2H2O = Na2

2. Salt + Basic oxide ≠ the reaction does not go

3. Salt + amphoteric oxide → (t) acid volatile oxide + salt

Amphoteric oxides displace volatile acid oxides from their salts during fusion.

Al2O3 + K2CO3 → KAlO2 + CO2

Fe2O3 + Na2CO3 → 2NaFeO2 + CO2

4. Salt + acidic non-volatile oxide → acidic volatile oxide + salt

Non-volatile acidic oxides displace volatile acidic oxides from their salts upon fusion.

SiO2 + CaCO3 → (t) CaSiO3 + CO2

P2O5 + Na2CO3 → (t) 2Na3PO4 + 3CO2

3SiO2 + Ca3 (PO4) 2 → (t) 3CaSiO3 + P2O5

5. Salt + base → base + salt

Reactions between salts and bases are ion exchange reactions. Therefore, under normal conditions, they proceed only in solutions (both salt and base must be soluble) and only under the condition that a precipitate or a weak electrolyte (H2O / NH4OH) is formed as a result of exchange; gaseous products are not formed in these reactions.

A. Soluble base + soluble acidic salt → medium salt + water

If the salt and the base are formed by different cations, then two average salts are formed; in the case of acidic ammonium salts, an excess of alkali leads to the formation of ammonium hydroxide.

Ba (OH) 2 + Ba (HCO3) → 2BaCO3 + 2H2O

2KOH + 2NaHCO3 → Na2CO3 + K2CO3 + 2H2O

2NaOH + 2NH4HS → Na2S + (NH4) 2S + 2H2O

2NaOH (g) + NH4Hs → Na2S + NH4OH + H2O

B. Soluble base + soluble medium / basic salt → insoluble salt ↓ + base

Soluble base + soluble medium / basic salt → salt + insoluble base ↓

Soluble base + soluble medium / basic salt → salt + weak electrolyte NH4OH

Soluble base + soluble medium / basic salt → no reaction

The reaction between soluble bases and the middle / basic salt occurs only if, as a result of ion exchange, an insoluble salt, or an insoluble base, or a weak electrolyte NH4OH is formed.

Ba (OH) 2 + Na2SO4 → BaSO4 ↓ + 2NaOH

2NH4OH + CuCl2 → 2NH4Cl + Cu (OH) 2 ↓

Ba (OH) 2 + NH4Cl → BaCl2 + NH4OH

NaOH + KCl ≠

If the starting salt is formed by a multi-acid base, a base salt is formed when there is a lack of alkali.

NaOH (short) + AlCl3 → Al (OH) Cl2 + NaCl

Under the action of alkalis on silver and mercury (II) salts, not AgOH and Hg (OH) 2 are released, which decompose at room temperature, but insoluble oxides Ag2O and HgO.

2AgNO3 + 2NaOH → Ag2O ↓ 2NaNO3 + H2O

Hg (NO3) 2 + 2KOH → HgO ↓ + 2KNO3 + H2O

6. Salt + amphoteric hydroxide → reaction does not go

7. Salt + acid → acid + salt

Basically. the reactions of acids with salts are reactions of ion exchange, therefore they occur in solutions and only if, in this case, an acid-insoluble salt or a weaker and more volatile acid is formed.

HCl + AgNO3 → AgCl ↓ + HNO3

2HBr + K2SiO3 → 2KBr + H2SiO3 ↓

2HNO3 + Na2CO3 → 2NaNO3 + H2O + CO2

A. Acid1 + salt of more volatile / weak acid2 → salt of acid1 + more volatile / weak acid2

Acids interact with solutions of salts of weaker or volatile acids... Regardless of the salt composition (medium, acidic, basic), as a rule, a medium salt and a weaker volatile acid are formed.

2CH3COOH + Na2S → 2CH3COONa + H2S

HCl + NaHS → NaCl + H2S

B. Strong acid + strong / medium acid salt → insoluble salt ↓ + acid

Strong acids interact with solutions of salts of other strong acids to form an insoluble salt. Non-volatile Н3РО4 (medium strength acid) displaces strong, but volatile hydrochloric HCl and nitric HNO3 acids from their salts, provided that an insoluble salt is formed.

H2SO4 + Ca (NO3) 2 → CaSO4 ↓ + 2HNO3

2H3PO4 + 3CaCl2 → Ca3 (PO4) 2 ↓ + 6HCl

H3PO4 + 3AgNO3 → Ag3PO4 ↓ + 3HNO3

B. Acid1 + basic acid salt1 → medium salt + water

When an acid acts on a basic salt of the same acid, a medium salt and water are formed.

HCl + Mg (OH) Cl → MgCl2 + H2O

D. Polybasic acid1 + medium / acidic acid salt1 → acidic acid salt1

When a polybasic acid acts on an average salt of the same acid, an acid salt is formed, and when an acid salt is acted upon, an acid salt is formed containing a greater number of hydrogen atoms.

H3PO4 + Ca3 (PO4) → 3CaHPO4

H3PO4 + CaHPO4 → Ca (H2PO4) 2

CO2 + H2O + CaCO3 → Ca (HCO3) 2

E. Acid H2S + salt Ag, Cu, Pb, Cd, Hg → Ag2S / CuS / PbS / CdS / HgS ↓ + acid

Weak and volatile hydrogen sulfide acid H2S displaces even strong acids from solutions of Ag, Cu, Pb, Cd and Hg salts, forming sulfide precipitates with them, insoluble not only in water, but also in the resulting acid.

H2S + CuSO4 → CuS ↓ + H2SO4

E. Acid + medium / complex salt with amphoteric Me in the anion → medium salt + amphoteric hydroxide ↓

→ medium salt + medium salt + H2O

When an acid acts on a medium or complex salt with an amphoteric metal in the anion, the salt is destroyed and forms:

a) in case of lack of acid - medium salt and amphoteric hydroxide

b) in the case of an excess of acid - two average salts and water

2HCl (weeks) + Na2ZnO2 → 2NaCl + Zn (OH) 2 ↓

2HCl (week) + Na2 → 2NaCl + Zn (OH) 2 ↓ + 2H2O

4HCl (g) + Na2ZnO2 → 2NaCl + ZnCl2 + 2H2O

4HCl (g) + Na2 → 2NaCl + ZnCl2 + 4H2O

It should be borne in mind that in some cases, ORP or complexation reactions occur between acids and salts. So, the OVR is joined by:

H2SO4 conc. and I‾ / Br‾ (products H2S and I2 / SO2 and Br2)

H2SO4 conc. and Fe² + (products SO2 and Fe³ + )

HNO3 dil. / Conc. and Fe² + (products NO / NO2 and Fe 3 + )

HNO3 dil. / Conc. and SO3²‾ / S²‾ (NO / NO2 products and sulfate / sulfur or sulfate)

HCl conc. and KMnO4 / K2Cr2O7 / KClO3 (products are chlorine (gas) and Mn²+ / Cr³ + / Cl‾.

G. The reaction takes place without solvent.

Sulfuric acid conc. + salt (solid) → acidic / medium salt + acidic

Non-volatile acids can displace volatile acids from their dry salts. Most often, the interaction of concentrated sulfuric acid with dry salts of strong and weak acids is used, with the formation of an acid and an acidic or medium salt.

H2SO4 (conc) + NaCl (tv) → NaHSO4 + HCl

H2SO4 (conc) + 2NaCl (tv) → Na2SO4 + 2HCl

H2SO4 (conc) + KNO3 (tv) → KHSO4 + HNO3

H2SO4 (conc) + CaCO3 (tv) → CaSO4 + CO2 + H2O

8. Soluble salt + soluble salt → insoluble salt ↓ + salt

Reactions between salts are exchange reactions. Therefore, under normal conditions, they proceed only if:

a) both salts are soluble in water and taken in the form of solutions

b) as a result of the reaction, a precipitate or a weak electrolyte is formed (the latter is very rare).

AgNO3 + NaCl → AgCl ↓ + NaNO3

If one of the starting salts is insoluble, the reaction proceeds only when an even more insoluble salt is formed as a result. The criterion for "insolubility" is the value of PR (solubility product), however, since its study is beyond the scope of the school course, the cases when one of the reagent salts is insoluble are not further considered.

If a salt is formed in the exchange reaction, which completely decomposes as a result of hydrolysis (in the solubility table there are dashes in place of such salts), then the products of the hydrolysis of this salt become the reaction products.

Al2 (SO4) 3 + K2S ≠ Al2S3 ↓ + K2SO4

Al2 (SO4) 3 + K2S + 6H2O → 2Al (OH) 3 ↓ + 3H2S + K2SO4

FeCl3 + 6KCN → K3 + 3KCl

AgI + 2KCN → K + KI

AgBr + 2Na2S2O3 → Na3 + NaBr

Fe2 (SO4) 3 + 2KI → 2FeSO4 + I2 + K2SO4

NaCl + NaHSO4 → (t) Na2SO4 + HCl

Medium salts sometimes interact with each other to form complex salts. OVR is possible between salts. Some salts interact when fusing.

9. Salt of less active metal + metal more active → metal less active ↓ + salt

The more active metal displaces the less active metal (standing to the right in the series of stress) from the solution of its salt, thus forming a new salt, and the less active metal is released in a free form (settles on the plate of the active metal). Exception - alkali and alkaline earth metals in solution interact with water.

Salts with oxidizing properties in solution enter into other redox reactions with metals.

FeSO4 + Zn → Fe ↓ + ZnSO4

ZnSO4 + Fe ≠

Hg (NO3) 2 + Cu → Hg ↓ + Cu (NO3) 2

2FeCl3 + Fe → 3FeCl2

FeCl3 + Cu → FeCl2 + CuCl2

HgCl2 + Hg → Hg2Cl2

2CrCl3 + Zn → 2CrCl2 + ZnCl2

Metals can displace each other from molten salts (the reaction is carried out without air access). It should be remembered that:

a) when melted, many salts decompose

b) a series of metal stress determines the relative activity of metals only in aqueous solutions (for example, Al in aqueous solutions is less active than alkaline earth metals, and in melts it is more active)

K + AlCl3 (melt) → (t) 3KCl + Al

Mg + BeF2 (melt) → (t) MgF2 + Be

2Al + 3CaCl2 (melt) → (t) 2AlCl3 + 3Ca

10. Salt + non-metal

Reactions of salts with non-metals are few. These are redox reactions.

5KClO3 + 6P → (t) 5KCl + 3P2O5

2KClO3 + 3S → (t) 2KCl + 2SO2

2KClO3 + 3C → (t) 2KCl + 3CO2

More active halogens displace less active halogen salts from solutions. An exception is molecular fluorine, which in solutions reacts not with salt, but with water.

2FeCl2 + Cl2 → (t) 2FeCl3

2NaNO2 + O2 → 2NaNO3

Na2SO3 + S → (t) Na2S2O3

BaSO4 + 2C → (t) BaS + 2CO2

2KClO3 + Br2 → (t) 2KBrO3 + Cl2 (the same reaction is typical for iodine)

2KI + Br2 → 2KBr + I2 ↓

2KBr + Cl2 → 2KCl + Br2 ↓

2NaI + Cl2 → 2NaCl + I2 ↓

11. Decomposition of salts.

Salt → (t) thermal decomposition products

1. Salts of nitric acid

The products of thermal decomposition of nitrates depend on the position of the metal cation in the series of metal stresses.

MeNO3 → (t) (for Me to the left of Mg (excluding Li)) MeNO2 + O2

MeNO3 → (t) (for Me from Mg to Cu and also Li) MeO + NO2 + O2

MeNO3 → (t) (for Me to the right of Cu) Me + NO2 + O2

(during the thermal decomposition of iron (II) / chromium (II) nitrate, iron (III) / chromium (III) oxide is formed.

2. Ammonium salts

All ammonium salts decompose on ignition. Most often, this produces ammonia NH3 and acid or its decomposition products.

NH4Cl → (t) NH3 + HCl (= NH4Br, NH4I, (NH4) 2S)

(NH4) 3PO4 → (t) 3NH3 + H3PO4

(NH4) 2HPO4 → (t) 2NH3 + H3PO4

NH4H2PO4 → (t) NH3 + H3PO4

(NH4) 2CO3 → (t) 2NH3 + CO2 + H2O

NH4HCO3 → (t) NH3 + CO2 + H2O

Sometimes ammonium salts containing oxidizing anions decompose on heating with the release of N2, NO or N2O.

(NH4) Cr2O7 → (t) N2 + Cr2O3 + 4H2O

NH4NO3 → (t) N2O + 2H2O

2NH4NO3 → (t) N2 + 2NO + 4H2O

NH4NO2 → (t) N2 + 2H2O

2NH4MnO4 → (t) N2 + 2MnO2 + 4H2O

3. Salts of carbonic acid

Almost all carbonates decompose to metal oxide and CO2. Carbonates alkali metals except for lithium do not decompose when heated. Silver and mercury carbonates decompose to free metal.

MeCO3 → (t) MeO + CO2

2Ag2CO3 → (t) 4Ag + 2CO2 + O2

All bicarbonates are decomposed to the corresponding carbonate.

MeHCO3 → (t) MeCO3 + CO2 + H2O

4. Sulfurous acid salts

When heated, sulfites disproportionate, forming sulfide and sulfate. The sulfide (NH4) 2S formed during the decomposition of (NH4) 2SO3 immediately decomposes into NH3 and H2S.

MeSO3 → (t) MeS + MeSO4

(NH4) 2SO3 → (t) 2NH3 + H2S + 3 (NH4) 2SO4

Hydrosulfites decompose to sulfites, SO2 and H2O.

MeHSO3 → (t) MeSO3 + SO2 + H2O

5. Sulfuric acid salts

Many sulfates decompose at t> 700-800 C to metal oxide and SO3, which decomposes to SO2 and O2 at this temperature. Sulfates of alkali metals are heat-resistant. Silver and mercury sulfates decompose to free metal. Hydrosulfates decompose first to disulfates and then to sulfates.

2CaSO4 → (t) 2CaO + 2SO2 + O2

2Fe2 (SO4) 3 → (t) 2Fe2O3 + 6SO2 + 3O2

2FeSO4 → (t) Fe2O3 + SO3 + SO2

Ag2SO4 → (t) 2Ag + SO2 + O2

MeHSO4 → (t) MeS2O7 + H2O

MeS2O7 → (t) MeSO4 + SO3

6. Complex salts

Hydroxocomplexes of amphoteric metals decompose mainly into medium salt and water.

K → (t) KAlO2 + 2H2O

Na2 → (t) ZnO + 2NaOH + H2O

7. Basic salts

Many basic salts decompose when heated. Basic salts of anoxic acids decompose into water and oxosalts

Al (OH) 2Br → (t) AlOBr + H2O

2AlOHCl2 → (t) Al2OCl4 + H2O

2MgOHCl → (t) Mg2OCl2 + H2O

Basic salts of oxygen-containing acids decompose into metal oxide and thermal decomposition products of the corresponding acid.

2AlOH (NO3) 2 → (t) Al2O3 + NO2 + 3O2 + H2O

(CuOH) 2CO3 → (t) 2CuO + H2O + CO2

8. Examples of thermal decomposition of other salts

4K2Cr2O7 → (t) 4K2CrO4 + 2Cr2O3 + 3O2

2KMnO4 → (t) K2MnO4 + MnO2 + O2

KClO4 → (t) KCl + O2

4KClO3 → (t) KCl + 3KClO4

2KClO3 → (t) 2KCl + 3O2

2NaHS → (t) Na2S + H2S

2CaHPO4 → (t) Ca2P2O7 + H2O

Ca (H2PO4) 2 → (t) Ca (PO3) 2 + 2H2O

2AgBr → (hν) 2Ag + Br2 (= AgI)

Most of the presented material is taken from the manual by N.E.Deryabina. "Chemistry. Main classes inorganic substances". IPO" At Nikitskiye Vorota "Moscow 2011.

1. Being active oxidizing agents, halogens react with metals. The reactions of metals with fluorine are especially vigorous. Alkali metals react explosively with it. When heated, halogens react even with gold and platinum. In an atmosphere of fluorine and chlorine, a number of metals are burned without preheating. Let us recall some of the features of these interactions. Iron and chromium, when reacted with fluorine, chlorine and bromine, are oxidized to a trivalent cation. The reaction with iodine already requires significant heating and leads to the formation of FeJ 2 and CrJ 2. Some metals are passivated in a halogen environment due to the formation of a protective salt film. In particular, copper interacts with fluorine only at high temperatures due to the formation of a CuF 2 film. Nickel behaves similarly. Gaseous fluorine is stored and transported in vessels made of monel metal (an alloy of nickel with iron and manganese). The reaction of chlorine with some metals is inhibited and greatly accelerated by traces of water, which acts as a catalyst in these cases. Well-dried chlorine, for example, does not react with iron, so liquefied chlorine is stored in steel cylinders. The liquid state of aggregation of bromine is the reason that it reacts with some metals more actively than chlorine, since the concentration of the reagent in the liquid phase is higher than the concentration in the gas. For example, compact aluminum and iron react with bromine at room temperature and with chlorine when heated.

2. Fluorine reacts with hydrogen explosively at room temperature, the reaction proceeds at a noticeable rate even at –252 0 С. Chlorine reacts only under ultraviolet or solar irradiation, since the reaction is free radical in nature. The reaction with bromine is less active and already requires heating, and therefore becomes noticeably reversible due to insufficient thermal stability of the H-Br bond. Communication energy H-J else less, the oxidizing ability of iodine is also noticeably lower than that of other halogens, therefore the equilibrium of the reaction H 2 + J 2 = 2HJ at temperatures at which the reaction rate is not very low is significantly shifted towards the starting materials.

3. Sulfur and phosphorus burn when interacting with fluorine, chlorine and bromine. In this case, compounds are formed with fluorine, in which these elements show their maximum oxidation state: SF 6 and PF 5. The products of the remaining reactions depend on the experimental conditions - PCl 3, PCl 5, PBr 3, PBr 5, S 2 Cl 2, S 2 Br 2, SCl 2.

4. With other non-metals, halogens also react with varying activity. The exception is oxygen and nitrogen, with which halogens do not react directly. Halogen oxides of various structures, depending on the conditions, can be obtained by their reaction with ozone.

5. The activity of fluorine is so great that it is capable of interacting even with noble gases (except for He, Ne, Ar).

6. Interacting with each other, halogens form binary compounds of various compositions, in which a more electronegative halogen exhibits a negative oxidation state, and a less negative one - a positive one. For example, ClF 5, BrCl 3, JF 7, JCl.

Reactions with complex substances

1. Water ignites spontaneously in an atmosphere of fluorine, and the reaction proceeds until the fluorine is completely consumed. Depending on the temperature and other conditions, a number of reactions take place: 3F 2 + 3H 2 O = F 2 O + 4HF + H 2 O 2 2F 2 + H 2 O = F 2 O + 2HF; with steam with explosion: 2F 2 + 2H 2 O = 4HF + O 2 3F 2 + 3H 2 O = 6HF + O 3; with ice: F 2 + H 2 O = HOF + HF. Chlorine, limitedly dissolving in water (2 volumes of chlorine (gas!) Per 1 volume of water), reacts reversibly with it: Cl 2 + H 2 O = HCl + HClO. Bromine behaves similarly, but the equilibrium Br 2 + H 2 O = HBr + HBrO is more strongly shifted to the left. A similar equilibrium for iodine is shifted towards the reagents so much that it can be said that the reaction is not proceeding. In accordance with the above, there are chlorine and bromine water, but iodine and fluorine do not exist. At the same time, an iodide anion was found in an aqueous solution of iodine in low concentrations, the appearance of which is explained by the formation of iodine hydrate in the solution, which is capable of dissociating into J +. H 2 O and J -. The dissociation equilibrium of iodine hydrate is also strongly shifted towards the undissociated form.

2. Consider the reactions of halogens with acids. Redox reactions are possible, in which there is an exchange of electrons between the halogen and the element that is part of the acid. In this case, chlorine and bromine often act as oxidizing agents, and iodine as a reducing agent. Here are the most typical reactions: J 2 + 10HNO 3 (conc) = 2HJO 3 + 10NO 2 + 4H 2 O 3J 2 + 10HNO 3 = 6HNO 3 + 10NO + 2H 2 O 2HBr + Cl 2 = 2HCl + Br 2 H 2 SO 3 (SO 2 + H 2 O) + Br 2 + H 2 O = 2HBr + H 2 SO 4 HCOOH + Cl 2 (Br 2) = CO 2 + 2HCl (HBr). Reactions with fluorine lead to destruction.

3. When interacting with alkalis, halogens disproportionate, that is, they simultaneously increase and decrease their oxidation state. Chlorine reacts in the cold: Cl 2 + 2NaOH = NaCl + NaClO, and when heated - 3Cl 2 + 6NaOH = 5NaCl + NaClO 3 + 3H 2 O, because the hypochlorite anion, when heated in solution, disproportionates into chlorate and chloride. Hypobromites and hypoiodites are even less stable; therefore, bromine and iodine at room temperature already give bromates and iodates. For example: 3J 2 + 6KOH = 5KJ + KJO 3. The reaction of chlorine in the cold with calcium hydroxide leads to the formation of a mixed salt of calcium chloride-hypochlorite - bleach: Cl 2 + Ca (OH) 2 = CaOCl 2 + H 2 O.

4. Unlike most substances, fluorine reacts with silicon dioxide at room temperature. The reaction is catalyzed by traces of water. Since SiO 2 is the main constituent of glass, fluorine dissolves glass in accordance with the reaction: 2F 2 + SiO 2 = SiF 4 + O 2.

5. When interacting with salts, oxides and others binary compounds redox reactions are possible, of which the displacement reactions with a more active (more electronegative) halogen less active from the salt composition should be noted, for example: 2KJ + Cl 2 = 2KCl + J 2. An external sign of this reaction is the appearance of a yellow (brown at a significant concentration) color of molecular iodine. With prolonged passage of chlorine through a solution of potassium iodide, the color disappears, since iodine is further oxidized to HJO 3, the solution of which is colorless: J 2 + 5Cl 2 + 6H 2 O = 10HCl + 2HJO 3.

Halogen compounds

1. Hydrogen halides- gaseous substances under normal conditions. The boiling point of hydrogen fluoride is +19 0 C (HCl -85 0 C, HBr -67 0 C, HJ -35 0 C). It is abnormally large due to the formation of very strong hydrogen bonds in liquid hydrogen fluoride. Due to strong hydrogen bonds in liquid hydrogen fluoride, there are no free ions, and it does not conduct electricity, being a non-electrolyte. All molecules of hydrogen halides have single, highly polar bonds. When moving along the group from top to bottom, the polarity of the bond decreases, since the negative end of the dipole of the hydrogen-halogen bond is halogen, and electronegativity decreases significantly from fluorine to iodine. But the bond strength is largely affected by an increase in the bond length, therefore the strongest bond in the series under consideration is in the HF molecule, and the weakest is in the HJ molecule. All hydrogen halides are readily soluble in water. In this case, ionization and dissociation occurs. Dissociation produces a hydronium cation, therefore aqueous solutions hydrogen halides have the properties of acids. Hydrochloric (hydrochloric), hydrobromic and hydroiodic are strong acids. The strongest of them is hydrogen iodide, not only because of the weaker bond in the molecule, but also because of the greater stability of the iodide ion, the concentration of the charge in which is reduced due to its large size. Hydrofluoric (hydrofluoric) acid is weak due to the presence of hydrogen bonds not only between hydrogen fluoride molecules, but also between hydrogen fluoride and water molecules. These bonds are so strong that acid fluorides can form in concentrated solutions, although hydrofluoric acid is monobasic: KOH + 2HF = KHF 2. The acidic difluoride anion has a strong hydrogen bond:. Hydrofluoric acid also reacts with glass, the reaction in general view looks like this: SiO 2 + 6HF = H 2 + 2H 2 O. Hydrohalic acids exhibit all the properties of non-oxidizing acids. But since many metals are prone to the formation of acido complex anions; they sometimes react with metals that are in the series of voltages after hydrogen. For example, 2Cu + 4HI = 2H + H 2. Hydrogen fluoride and hydrogen chloride are not oxidized by concentrated sulfuric acid, so they can be obtained from dry halides, for example, ZnCl 2 (TV) + H 2 SO 4 (conc) = ZnSO 4 + 2HCl. Hydrogen bromide and hydrogen iodide under these conditions are oxidized: 2HBr + H 2 SO 4 (conc) = Br 2 + SO 2 + 2H 2 O; 8HI + H 2 SO 4 (conc) = 4I 2 + H 2 S + 4H 2 O. Absolute phosphoric acid is used to displace them from the composition of salts, which practically does not exhibit oxidizing properties. Concentrated nitric acid oxidizes hydrogen chloride to chlorine, which is a very strong oxidizing agent at the time of isolation. A mixture of concentrated nitric and hydrochloric acids is called aqua regia and is capable of dissolving gold and platinum: Au + HNO 3 + 4HCl = H + NO + 2H 2 O. Hydrogen chloride and concentrated hydrochloric acid are also oxidized by other strong oxidants (MnO 2, KMnO 4, K 2 Cr 2 O 7). These reactions are used as laboratory methods obtaining molecular chlorine. Hydrogen halides can also be obtained by hydrolysis of most non-metal halides. When HI is obtained, a mixture of iodine with red phosphorus is directly affected by water: 2P + 3I 2 + 6H 2 O = 2H 3 PO 3 + 6HI. It should be recalled that direct synthesis from simple substances is possible only for HF and HCl.

2. Salts of hydrohalic acids... Most salts are soluble. Slightly soluble are salts of bivalent lead and insoluble - salts of silver. The interaction of the silver cation and halide ions is a qualitative reaction: AgF is soluble, AgCl is a white cheesy precipitate, AgBr is a pale yellow precipitate, AgI is a bright yellow precipitate. Some metal halides, such as halides (other than fluoride) of aluminum and mercury, are covalent compounds. Aluminum chloride is capable of sublimation, soluble mercury halides dissociate in water stepwise. Tin (IV) chloride - liquid.

3. A qualitative reaction to molecular iodine is the appearance of a blue coloration with a starch solution.

4. Oxygen compounds halogens... Fluorine forms two compounds with oxygen: F 2 O - oxygen fluoride - light yellow gas with bip = -144.8 ° C; is obtained by rapidly passing fluorine through a 2% sodium hydroxide solution. Dioxygen difluoride - F 2 O 2 is a light brown gas, at -57 ° C it turns into a cherry-red liquid, and at -163 ° C it turns into an orange solid. It turns out F 2 O 2 when simple substances interact with cooling and the action of an electric glow discharge. Above the boiling point, it is already unstable and acts as a strong oxidizing and fluorinating agent. Other halogen oxides are endothermic and unstable. At room temperature, some of them, for example, Cl 2 O 7, exist only due to the kinetic inhibition of the decomposition process. Chlorine oxide (VII) is a colorless liquid with a boiling point of 83 ° C, which decomposes explosively when heated to 120 ° C. The only exothermic compound of halogen and oxygen is J 2 O 5. It is a white crystalline substance that decomposes into simple substances without explosion at temperatures above 300 ° C. It is used to detect and quantify carbon monoxide (II) in air: J 2 O 5 + 5CO = J 2 + 5CO 2.

5. Oxygenated halogen acids... Acids of general formula HEO x are known in which halogens exhibit odd positive oxidation states. For chlorine, this is HClO - hypochlorous acid, weak, unstable. Decomposes according to the equation: HClO = HCl + O, and oxygen at the time of release exhibits very strong oxidizing properties. It turns out by the reaction: 2Cl 2 + 2HgO + H 2 O = HgO. HgCl 2 ↓ + 2HClO, salts are called hypochlorites. HClO 2 - chloride acid is also weak and unstable. Salts - chlorites. HClO 3 - chloric acid. It is already a strong acid, but it is stable only in dilute aqueous solutions. In terms of oxidizing ability, it is somewhat inferior to chlorous acid. Salts - chlorates. Chlorine acid - HClO 4 - one of the strongest inorganic acids. Its aqueous solutions are stable and safe during storage, usually a 72% solution is used, which shows almost no oxidizing properties. Perchloric acid exists in free form as a colorless highly fuming liquid that can explode when stored or heated. Salts are called perchlorates. Thus, with an increase in the number of oxygen atoms, the strength of oxygen-containing chlorine acids increases and their oxidizing ability decreases. The corresponding acids of bromine and iodine have similar properties, but they are much less stable. Especially in the oxidation states of halogens +1 and +3. Solutions hypobromous acids are stable for a short time only at 0 ° С. Bromic acid in everything resembles chloric ... Iodic acid - colorless transparent crystals with t pl = 110 ° C. It is obtained by oxidation of iodine with concentrated nitric acid, hydrogen peroxide, ozone, chlorine in water: J 2 + 5H 2 O 2 = 2HJO 3 + 4H 2 O Bromine acid, in contrast to perchloric acid, is a strong oxidizing agent and is not isolated in a free state, which is associated with the phenomenon of secondary periodicity, as a result of which it is disadvantageous for bromine to exhibit the maximum positive oxidation state. There are several iodine acids: HJO 4, H 5 JO 6 (orthoiodic), H 3 JO 5 (metaiodic). The most stable is H 5 JO 6. This is a colorless crystalline substance with t pl = 122 ° C, it is an acid of medium strength and is prone to the formation of acid salts, since the basic equilibria in its solution are as follows: H 5 JO 6 = H + + H 4 JO 6 - K = 10 -3 H 4 JO 6 - = JO 4 - + 2H 2 OK = 29 H 4 JO 6 - = H + + H 3 JO 6 - K = 2. 10 -7. Let's summarize. Strong acids are HClO 4, HClO 3, HBrO 4, HBrO 3, HJO 3. HClO, HClO 2, HBrO, HBrO 4, H 5 JO 6 have strong oxidizing properties.

6. Oxygenated acid salts more stable than acids. Interestingly, perchlorates and periodates are insoluble for the metals of the potassium subgroup, and for rubidium there are also chlorates, bromates and perbromates, although usually all salts of alkali metals are soluble. Most salts decompose when heated: KClO 4 = KCl + 2O 2. Potassium chlorate, which also has the name "Berthollet's salt", disproportionates when heated: 4KClO 3 = KCl + 3KClO 4 Hypochlorite also behaves: 3KClO = 2KCl + KClO 3 If the salt contains impurities, especially metal oxides, decomposition can partly go the other way : 2KClO 3 = 2KCl + 3O 2. When manganese dioxide is used as a catalyst, this path becomes the main one.

7. Redox reactions of oxohalogenate anions. The salts are completely dissociated in solution. In this case, oxohalogenate anions - EO x - are obtained, which, in the presence of a negative charge, are weaker oxidizing agents than acid molecules. For example, hypochlorous acid can oxidize its own salt: 2HClO + NaClO = NaClO 3 + 2HCl. In solution, salts show noticeable oxidizing properties only in an acidic environment. It is worth noting the counterproportionation reactions: KClO 3 + 6HCl = 3Cl 2 + KCl + 3H 2 O KJO 3 + 5KJ + H 2 SO 4 = 3J 2 ↓ + 3K 2 SO 4 + 3H 2 O. When heated, these salts become strong oxidizing agents. The entire match and pyrotechnic industries are based on the reactions of berthollet salt, for example: 2KClO 3 + 3S = 2KCl + 3SO 2 5KClO 3 + 6P = 5KCl + 3P 2 O 5 KClO 3 + 2Al = Al 2 O 3 + KCl. Complex equilibria lead to the fact that oxygen-containing halogen acids and their salts, acting as oxidants, are most often reduced to Hal -1.

8. Methods for producing halogens. Fluorine is produced by electrolysis of potassium hydrofluoride (KHF 2) melt. In industry, chlorine is obtained by electrolysis of a solution of sodium chloride or hydrochloric acid, according to the Deacon method: 4HCl + O 2 = 2H 2 O + 2Cl 2 (when heated and using CuCl 2 as a catalyst), the interaction of bleach with hydrochloric acid. In the laboratory: the interaction of concentrated hydrochloric acid with KMnO 4, K 2 Cr 2 O 7 or MnO 2 when heated. Bromine is obtained by displacing it with chlorine from the composition of potassium or sodium bromide, as well as by oxidizing bromides with concentrated sulfuric acid. All of these reactions have already been discussed. Iodine can also be displaced by chlorine or bromine from the iodide composition. You can oxidize the iodide anion with manganese dioxide in an acidic environment. Since the iodide anion is easily oxidized, a wide variety of reactions are possible here.

COPPER.

Element with serial number 29, relative atomic mass 63.545. Belongs to the d-element family. In the periodic system, it is in the IV period, the I group, a side subgroup. The structure of the outer electron layer: 3d 10 4s 1. In the ground state, the d-sublevel is filled, but it is not sufficiently stable, therefore, in addition to the oxidation state +1, which can be assumed from the electronic structure of the atom, copper exhibits oxidation states +2, even +3 and very rarely +4. The radius of the copper atom is quite small - 0.128 nm. It is even smaller than the radius of the lithium atom - 0.155 nm. Its only 4s electron, when it is closer to the nucleus, falls under the screen from the completed 3d 10 shell, which increases its attraction to the nucleus, and with it the ionization potential. Therefore, copper is an inactive metal, after hydrogen in the series of voltages.

Physical properties. Copper is a soft red metal, ductile, tough, and easily stretches into a wire. It has high thermal and electrical conductivity, which is second only to gold and silver.

Chemical properties of a simple substance. In dry air, copper is quite inert, as it is covered with a thin film of a mixture of CuO and Cu 2 O, which gives the surface a darker color and prevents further interaction with atmospheric oxygen. In the presence of significant amounts of moisture and carbon dioxide, corrosion occurs, the product of which is green hydroxymedium (II) carbonate: 2Cu + H 2 O + CO 2 + O 2 = (CuO) 2 CO 3.

All chemical elements are divided into metals and non-metals depending on the structure and properties of their atoms. Also, simple substances formed by elements are classified into metals and non-metals, based on their physical and chemical properties.

In the Periodic Table of Chemical Elements D.I. Mendeleev's non-metals are located diagonally: boron - astatine and above it in the main subgroups.

For metal atoms, relatively large radii and a small number of electrons at the outer level from 1 to 3 are characteristic (exception: germanium, tin, lead - 4; antimony and bismuth - 5; polonium - 6 electrons).

Nonmetal atoms, on the contrary, are characterized by small atomic radii and the number of electrons at the outer level from 4 to 8 (with the exception of boron, it has three such electrons).

Hence the tendency of metal atoms to give up external electrons, i.e. reducing properties, and for nonmetal atoms - the desire to receive electrons missing to a stable eight-electron level, i.e. oxidizing properties.

Metals

In metals, there is a metallic bond and a metallic crystal lattice. At the lattice sites there are positively charged metal ions bound by means of shared external electrons belonging to the entire crystal.

This determines all the most important physical properties of metals: metallic luster, electrical and thermal conductivity, plasticity (the ability to change shape under external influence) and some others characteristic of this class of simple substances.

Metals of group I of the main subgroup are called alkali metals.

Group II metals: calcium, strontium, barium - alkaline earth.

Chemical properties of metals

In chemical reactions, metals exhibit only reducing properties, i.e. their atoms donate electrons, resulting in positive ions.

1. Interact with non-metals:

a) oxygen (with the formation of oxides)

Alkali and alkaline earth metals oxidize easily under normal conditions, so they are stored under a layer of petroleum jelly or kerosene.

4Li + O 2 = 2Li 2 O

2Ca + O 2 = 2CaO

Please note: when sodium interacts - peroxide is formed, potassium - superoxide

2Na + O 2 = Na 2 O 2, K + O2 = KO2

and the oxides are obtained by calcining the peroxide with the corresponding metal:

2Na + Na 2 O 2 = 2Na 2 O

Iron, zinc, copper and other less active metals are slowly oxidized in air and actively when heated.

3Fe + 2O 2 = Fe 3 O 4 (a mixture of two oxides: FeO and Fe 2 O 3)

2Zn + O 2 = 2ZnO

2Cu + O 2 = 2CuO

Gold and platinum metals are not oxidized by atmospheric oxygen under any conditions.

b) hydrogen (with the formation of hydrides)

2Na + H 2 = 2NaH

Ca + H 2 = CaH 2

c) chlorine (with the formation of chlorides)

2K + Cl 2 = 2KCl

Mg + Cl 2 = MgCl 2

2Al + 3Cl 2 = 2AlCl 3

Please note: when iron interacts, iron (III) chloride is formed:

2Fe + 3Cl 2 = 2FeCl 3

d) sulfur (with the formation of sulfides)

2Na + S = Na 2 S

Hg + S = HgS

2Al + 3S = Al 2 S 3

Please note: when iron interacts, iron (II) sulfide is formed:

Fe + S = FeS

e) nitrogen (with the formation of nitrides)

6K + N 2 = 2K 3 N

3Mg + N 2 = Mg 3 N 2

2Al + N 2 = 2AlN

2. Interact with complex substances:

It must be remembered that according to their reductive ability, metals are arranged in a row, which is called the electrochemical series of voltages or the activity of metals (displacement series of N.N. Beketov):

Li, K, Ba, Ca, Na, Mg, Al, Mn, Zn, Cr, Fe, Co, Ni, Sn, Pb, (H 2), Cu, Hg, Ag, Au, Pt

a) water

Metals located in a row up to magnesium, under normal conditions, displace hydrogen from water, forming soluble bases - alkalis.

2Na + 2H 2 O = 2NaOH + H 2

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

Magnesium interacts with water when boiled.

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

When removing the oxide film, aluminum reacts violently with water.

2Al + 6H 2 O = 2Al (OH) 3 + 3H 2

The rest of the metals in the row up to hydrogen, under certain conditions, can also react with water with the release of hydrogen and the formation of oxides.

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

b) acid solutions

(Except concentrated sulfuric acid and nitric acid of any concentration. See "Redox reactions" section.)

Please note: do not use insoluble silicic acid to carry out the reactions

Metals ranging from magnesium to hydrogen displace hydrogen from acids.

Mg + 2HCl = MgCl 2 + H 2

Please note: ferrous salts are formed.

Fe + H 2 SO 4 (dil.) = FeSO 4 + H 2

The formation of insoluble salt prevents the reaction from proceeding. For example, lead practically does not react with sulfuric acid solution due to the formation of insoluble lead sulfate on the surface.

Metals ranked next to hydrogen DO NOT displace hydrogen.

c) salt solutions

Metals that rank up to magnesium and actively react with water are not used to carry out such reactions.

For the rest of the metals, the following rule is fulfilled:

Each metal displaces from salt solutions other metals located in a row to the right of it, and itself can be displaced by metals located to the left of it.

Cu + HgCl 2 = Hg + CuCl 2

Fe + CuSO 4 = FeSO 4 + Cu

As with acid solutions, the formation of an insoluble salt prevents the reaction from proceeding.

d) alkali solutions

Metals interact, hydroxides of which are amphoteric.

Zn + 2NaOH + 2H 2 O = Na 2 + H 2

2Al + 2KOH + 6H 2 O = 2K + 3H 2

e) with organic substances

Alkali metals with alcohols and phenol.

2C 2 H 5 OH + 2Na = 2C 2 H 5 ONa + H 2

2C 6 H 5 OH + 2Na = 2C 6 H 5 ONa + H 2

Metals participate in reactions with haloalkanes, which are used to obtain lower cycloalkanes and for syntheses, during which the carbon skeleton of the molecule becomes more complex (A. Würz's reaction):

CH 2 Cl-CH 2 -CH 2 Cl + Zn = C 3 H 6 (cyclopropane) + ZnCl 2

2CH 2 Cl + 2Na = C 2 H 6 (ethane) + 2NaCl

Nonmetals

In simple substances, the atoms of non-metals are bound by a covalent non-polar connection... In this case, single (in H 2, F 2, Cl 2, Br 2, I 2 molecules), double (in O 2 molecules), triple (in N 2 molecules) covalent bonds are formed.

The structure of simple substances - non-metals:

1.molecular

Under normal conditions, most of these substances are gases (H 2, N 2, O 2, O 3, F 2, Cl 2) or solids (I 2, P 4, S 8) and only the only bromine (Br 2) is liquid. All these substances have a molecular structure and are therefore volatile. In the solid state, they are fusible due to the weak intermolecular interaction that holds their molecules in the crystal, and are capable of sublimation.

2.atomic

These substances are formed by crystals, in the nodes of which there are atoms: (B n, C n, Si n, Gen, Se n, Te n). Due to the high strength of covalent bonds, they, as a rule, have a high hardness, and any changes associated with the destruction of the covalent bond in their crystals (melting, evaporation) are performed with a large expenditure of energy. Many of these substances have high temperatures melting and boiling, and their volatility is very low.

Many elements - non-metals - form several simple substances - allotropic modifications. Allotropy can be associated with a different composition of molecules: oxygen O 2 and ozone O 3 and with a different crystal structure: graphite, diamond, carbyne, fullerene are allotropic modifications of carbon. Elements - non-metals with allotropic modifications: carbon, silicon, phosphorus, arsenic, oxygen, sulfur, selenium, tellurium.

Chemical properties of non-metals

Atoms of non-metals are dominated by oxidizing properties, that is, the ability to attach electrons. This ability is characterized by the value of electronegativity. Among non-metals

At, B, Te, H, As, I, Si, P, Se, C, S, Br, Cl, N, O, F

electronegativity increases and oxidizing properties increase.

From this it follows that for simple substances - non-metals, both oxidizing and reducing properties will be characteristic, with the exception of fluorine, the strongest oxidizing agent.

1. Oxidizing properties

a) in reactions with metals (metals are always reducing agents)

2Na + S = Na 2 S (sodium sulfide)

3Mg + N 2 = Mg 3 N 2 (magnesium nitride)

b) in reactions with non-metals located to the left of the given one, that is, with a lower value of electronegativity. For example, in the interaction of phosphorus and sulfur, sulfur will be an oxidizing agent, since phosphorus has a lower electronegativity value:

2P + 5S = P 2 S 5 (phosphorus sulfide V)

Most non-metals will oxidize with hydrogen:

H 2 + S = H 2 S

H 2 + Cl 2 = 2HCl

3H 2 + N 2 = 2NH 3

c) in reactions with some complex substances

Oxidizing agent - oxygen, combustion reactions

CH 4 + 2O 2 = CO 2 + 2H 2 O

2SO 2 + O 2 = 2SO 3

Oxidizing agent - chlorine

2FeCl 2 + Cl 2 = 2FeCl 3

2KI + Cl 2 = 2KCl + I 2

CH 4 + Cl 2 = CH 3 Cl + HCl

Ch 2 = CH 2 + Br 2 = CH 2 Br-CH 2 Br

2. Restorative properties

a) in reactions with fluorine

S + 3F 2 = SF 6

H 2 + F 2 = 2HF

Si + 2F 2 = SiF 4

b) in reactions with oxygen (except for fluorine)

S + O 2 = SO 2

N 2 + O 2 = 2NO

4P + 5O 2 = 2P 2 O 5

C + O 2 = CO 2

c) in reactions with complex substances - oxidizing agents

H 2 + CuO = Cu + H 2 O

6P + 5KClO 3 = 5KCl + 3P 2 O 5

C + 4HNO 3 = CO 2 + 4NO 2 + 2H 2 O

H 2 C = O + H 2 = CH 3 OH

3. Disproportionation reactions: the same non-metal is both an oxidizing agent and a reducing agent

Cl 2 + H 2 O = HCl + HClO

3Cl 2 + 6KOH = 5KCl + KClO 3 + 3H 2 O

The world around us is material. Matter is of two types: matter and field. The object of chemistry is a substance (including the effect on matter of various fields - sound, magnetic, electromagnetic, etc.)

Substance - everything that has a rest mass (i.e. characterized by the presence of mass when it does not move)... So, although the rest mass of one electron (the mass of a non-moving electron) is very small - about 10 -27 g, but even one electron is a substance.

The substance comes in three aggregate states- gaseous, liquid and solid. There is one more state of matter - plasma (for example, there is plasma in thunderstorm and ball lightning), but in school course plasma chemistry is hardly considered.

Substances can be pure, very pure (necessary, for example, for the creation of fiber optics), they can contain appreciable amounts of impurities, they can be mixtures.

All substances are made up of tiny particles - atoms. Substances consisting of atoms of the same type(from atoms of one element), called simple(for example, charcoal, oxygen, nitrogen, silver, etc.). Substances that contain interconnected atoms of different elements are called complex.

If a substance (for example, in air) contains two or more simple substances, and their atoms are not connected with each other, then it is called not a complex, but a mixture of simple substances. The number of simple substances is relatively small (about five hundred), and the number of complex substances is enormous. To date, tens of millions of different complex substances are known.

Chemical transformations

Substances are capable of interacting with each other, and new substances appear. Such transformations are called chemical... For example, a simple substance coal interacts (chemists say - reacts) with another simple substance - oxygen, as a result of which a complex substance is formed - carbon dioxide, in which carbon and oxygen atoms are linked together. Such transformations of some substances into others are called chemical. Chemical transformations are chemical reactions. So, when sugar is heated in air, a complex sweet substance - sucrose (of which sugar is composed) - turns into a simple substance - coal and a complex substance - water.

Chemistry studies the transformation of some substances into others. The task of chemistry is to find out with which substances, under given conditions, this or that substance can interact (react), which is formed in this case. In addition, it is important to find out under what conditions one or another transformation can occur and the required substance can be obtained.

Physical properties of substances

Each substance is characterized by a combination of physical and chemical properties. Physical properties are properties that can be characterized using physical instruments... For example, a thermometer can be used to determine the melting and boiling points of water. Physical methods you can characterize the ability of a substance to conduct an electric current, determine the density of a substance, its hardness, etc. At physical processes substances remain unchanged in composition.

The physical properties of substances are subdivided into countable (those that can be characterized with the help of certain physical devices by a number, for example, by indicating the density, melting and boiling points, solubility in water, etc.) and uncountable (those that cannot be characterized by a number or it is very difficult - such as color, smell, taste, etc.).

Chemical properties of substances

The chemical properties of a substance is a collection of information about what other substances and under what conditions a given substance enters into chemical interactions.. The most important task chemistry - identifying the chemical properties of substances.

The smallest particles of substances - atoms - participate in chemical transformations. During chemical transformations, other substances are formed from some substances, and the original substances disappear, and instead of them new substances (reaction products) are formed. BUT atoms at of all chemical transformations are preserved... They are rearranged, during chemical transformations, old bonds between atoms are destroyed and new bonds appear.

Chemical element

Number various substances huge (and each of them has its own set of physical and chemical properties). Atoms, differing from each other in the most important characteristics, are relatively small in the material world around us - about a hundred. Each type of atom has its own chemical element. A chemical element is a collection of atoms with the same or similar characteristics.... About 90 different chemical elements are found in nature. By now, physicists have learned to create new types of atoms that are absent on Earth. Such atoms (and, accordingly, such chemical elements) are called artificial (in English - man-made elements). To date, more than two dozen artificially obtained elements have been synthesized.

Each element has a Latin name and a one- or two-letter symbol. In the Russian-language chemical literature, there are no clear rules for the pronunciation of the symbols of chemical elements. Some pronounce it like this: they call the element in Russian (symbols for sodium, magnesium, etc.), others - in Latin letters (symbols for carbon, phosphorus, sulfur), and still others - as the name of the element sounds in Latin (iron, silver, gold, mercury ). It is customary for us to pronounce the symbol of the element hydrogen H in the same way as this letter is pronounced in French.

A comparison of the most important characteristics of chemical elements and simple substances is shown in the table below. Several simple substances can correspond to one element (the phenomenon of allotropy: carbon, oxygen, etc.), and maybe one (argon and other inert gases).