In reactions, compounds can be formed. Introduction to general chemistry. Give the correct definition of a substitution reaction

1. Indicate the correct definition of a compound reaction: A. The reaction of the formation of several substances from one simple substance; B. A reaction in which one complex substance is formed from several simple or complex substances. B. A reaction in which substances exchange their constituents.

2. Indicate the correct definition of a substitution reaction: A. The reaction between a base and an acid; B. The reaction of interaction of two simple substances; B. A reaction between substances in which atoms of a simple substance replace atoms of one of the elements in a complex substance.

3. Indicate the correct definition of a decomposition reaction: A. A reaction in which several simple or complex substances are formed from one complex substance; B. A reaction in which substances exchange their constituents; B. Reaction with the formation of oxygen and hydrogen molecules.

5. What type of reaction is the interaction of acidic oxides with basic oxides: 5. What type of reaction is the interaction of acidic oxides with basic oxides: A. Exchange reaction; B. Compound reaction; B. Decomposition reaction; D. Substitution reaction.

7. Substances whose formulas are KNO 3 FeCl 2, Na 2 SO 4 are called: 7. Substances whose formulas are KNO 3 FeCl 2, Na 2 SO 4 are called: A) salts; B) reasons; B) acids; D) oxides. A) salts; B) reasons; B) acids; D) oxides. 8. Substances whose formulas are HNO 3, HCl, H 2 SO 4 are called: 8. Substances whose formulas are HNO 3, HCl, H 2 SO 4 are called: A) salts; B) acids; B) reasons; D) oxides. A) salts; B) acids; B) reasons; D) oxides. 9. Substances whose formulas are KOH, Fe(OH) 2, NaOH are called: 9. Substances whose formulas are KOH, Fe(OH) 2, NaOH are called: A) salts; B) acids; B) reasons; D) oxides. 10. Substances whose formulas are NO 2, Fe 2 O 3, Na 2 O are called: A) salts; B) acids; B) reasons; D) oxides. 10. Substances whose formulas are NO 2, Fe 2 O 3, Na 2 O are called: A) salts; B) acids; B) reasons; D) oxides. A) salts; B) acids; B) reasons; D) oxides. 11. Indicate the metals that form alkalis: 11. Indicate the metals that form alkalis: Cu, Fe, Na, K, Zn, Li. Cu, Fe, Na, K, Zn, Li.

7.1. Basic types of chemical reactions

Transformations of substances, accompanied by changes in their composition and properties, are called chemical reactions or chemical interactions. During chemical reactions, there is no change in the composition of the atomic nuclei.

Phenomena in which the shape or physical state of substances changes or the composition of atomic nuclei changes are called physical. An example of physical phenomena is the heat treatment of metals, during which their shape changes (forging), the melting of the metal, the sublimation of iodine, the transformation of water into ice or steam, etc., as well as nuclear reactions, as a result of which atoms are formed from atoms of some elements other elements.

Chemical phenomena can be accompanied by physical transformations. For example, as a result of chemical reactions occurring in a galvanic cell, an electric current arises.

Chemical reactions are classified according to various criteria.

1. According to the sign of the thermal effect, all reactions are divided into endothermic(proceeding with heat absorption) and exothermic(flowing with the release of heat) (see § 6.1).

2. Based on the state of aggregation of the starting substances and reaction products, they are distinguished:

homogeneous reactions, in which all substances are in the same phase:

2 KOH (p-p) + H 2 SO 4 (p-p) = K 2 SO (p-p) + 2 H 2 O (l),

CO (g) + Cl 2 (g) = COCl 2 (g),

SiO 2(k) + 2 Mg (k) = Si (k) + 2 MgO (k).

heterogeneous reactions, substances in which are in different phases:

CaO (k) + CO 2 (g) = CaCO 3 (k),

CuSO 4 (solution) + 2 NaOH (solution) = Cu(OH) 2 (k) + Na 2 SO 4 (solution),

Na 2 SO 3 (solution) + 2HCl (solution) = 2 NaCl (solution) + SO 2 (g) + H 2 O (l).

3. According to the ability to flow only in the forward direction, as well as in the forward and reverse directions, they distinguish irreversible And reversible chemical reactions (see § 6.5).

4. Based on the presence or absence of catalysts, they distinguish catalytic And non-catalytic reactions (see § 6.5).

5. According to the mechanism of their occurrence, chemical reactions are divided into ionic, radical etc. (the mechanism of chemical reactions occurring with the participation of organic compounds is discussed in the course of organic chemistry).

6. According to the state of oxidation states of the atoms included in the composition of the reacting substances, reactions occurring without changing the oxidation state atoms, and with a change in the oxidation state of atoms ( redox reactions) (see § 7.2) .

7. Reactions are distinguished by changes in the composition of the starting substances and reaction products connection, decomposition, substitution and exchange. These reactions can occur both with and without changes in the oxidation states of elements, table . 7.1.

Table 7.1

Types of chemical reactions

	General scheme	Examples of reactions that occur without changing the oxidation state of elements	Examples of redox reactions
Connections (one new substance is formed from two or more substances)		HCl + NH 3 = NH 4 Cl; SO 3 + H 2 O = H 2 SO 4	H 2 + Cl 2 = 2HCl; 2Fe + 3Cl 2 = 2FeCl 3
Decompositions (several new substances are formed from one substance)	A = B + C + D	MgCO 3 MgO + CO 2; H 2 SiO 3 SiO 2 + H 2 O	2AgNO 3 2Ag + 2NO 2 + O 2
Substitutions (when substances interact, atoms of one substance replace atoms of another substance in a molecule)	A + BC = AB + C	CaCO 3 + SiO 2 CaSiO 3 + CO 2	Pb(NO 3) 2 + Zn = Zn(NO 3) 2 + Pb; Mg + 2HCl = MgCl 2 + H 2
(two substances exchange their constituent parts, forming two new substances)	AB + CD = AD + CB	AlCl 3 + 3NaOH = Al(OH) 3 + 3NaCl; Ca(OH) 2 + 2HCl = CaCl 2 + 2H 2 O

7.2. Redox reactions

As mentioned above, all chemical reactions are divided into two groups:

Chemical reactions that occur with a change in the oxidation state of the atoms that make up the reactants are called redox reactions.

Oxidation is the process of giving up electrons by an atom, molecule or ion:

Na o – 1e = Na + ;

Fe 2+ – e = Fe 3+ ;

H 2 o – 2e = 2H + ;

2 Br – – 2e = Br 2 o.

Recovery is the process of adding electrons to an atom, molecule or ion:

S o + 2e = S 2– ;

Cr 3+ + e = Cr 2+ ;

Cl 2 o + 2e = 2Cl – ;

Mn 7+ + 5e = Mn 2+ .

Atoms, molecules or ions that accept electrons are called oxidizing agents. Restorers are atoms, molecules or ions that donate electrons.

By accepting electrons, the oxidizing agent is reduced during the reaction, and the reducing agent is oxidized. Oxidation is always accompanied by reduction and vice versa. Thus, the number of electrons given up by the reducing agent is always equal to the number of electrons accepted by the oxidizing agent.

7.2.1. Oxidation state

The oxidation state is the conditional (formal) charge of an atom in a compound, calculated under the assumption that it consists only of ions. The oxidation state is usually denoted by an Arabic numeral above the element symbol with a “+” or “–” sign. For example, Al 3+, S 2–.

To find oxidation states, the following rules are used:

the oxidation state of atoms in simple substances is zero;

the algebraic sum of the oxidation states of atoms in a molecule is equal to zero, in a complex ion - the charge of the ion;

the oxidation state of alkali metal atoms is always +1;

the hydrogen atom in compounds with non-metals (CH 4, NH 3, etc.) exhibits an oxidation state of +1, and with active metals its oxidation state is –1 (NaH, CaH 2, etc.);

The fluorine atom in compounds always exhibits an oxidation state of –1;

The oxidation state of the oxygen atom in compounds is usually –2, except for peroxides (H 2 O 2, Na 2 O 2), in which the oxidation state of oxygen is –1, and some other substances (superoxides, ozonides, oxygen fluorides).

The maximum positive oxidation state of elements in a group is usually equal to the group number. The exceptions are fluorine and oxygen, since their highest oxidation state is lower than the number of the group in which they are found. Elements of the copper subgroup form compounds in which their oxidation state exceeds the group number (CuO, AgF 5, AuCl 3).

The maximum negative oxidation state of elements located in the main subgroups of the periodic table can be determined by subtracting the group number from eight. For carbon it is 8 – 4 = 4, for phosphorus – 8 – 5 = 3.

In the main subgroups, when moving from elements from top to bottom, the stability of the highest positive oxidation state decreases; in secondary subgroups, on the contrary, from top to bottom the stability of higher oxidation states increases.

The conventionality of the concept of oxidation state can be demonstrated using the example of some inorganic and organic compounds. In particular, in phosphinic (phosphorous) H 3 PO 2, phosphonic (phosphorous) H 3 PO 3 and phosphoric H 3 PO 4 acids, the oxidation states of phosphorus are respectively +1, +3 and +5, while in all these compounds phosphorus is pentavalent. For carbon in methane CH 4, methanol CH 3 OH, formaldehyde CH 2 O, formic acid HCOOH and carbon monoxide (IV) CO 2, the oxidation states of carbon are –4, –2, 0, +2 and +4, respectively, while as the valency of the carbon atom in all these compounds is four.

Despite the fact that the oxidation state is a conventional concept, it is widely used in composing redox reactions.

7.2.2. The most important oxidizing and reducing agents

Typical oxidizing agents are:

1. Simple substances whose atoms have high electronegativity. These are, first of all, elements of the main subgroups VI and VII of groups of the periodic table: oxygen, halogens. Of the simple substances, the most powerful oxidizing agent is fluorine.

2. Compounds containing some metal cations in high oxidation states: Pb 4+, Fe 3+, Au 3+, etc.

3. Compounds containing some complex anions, the elements in which are in high positive oxidation states: 2–, –, etc.

Reducing agents include:

1. Simple substances whose atoms have low electronegativity are active metals. Non-metals, such as hydrogen and carbon, can also exhibit reducing properties.

2. Some metal compounds containing cations (Sn 2+, Fe 2+, Cr 2+), which, by donating electrons, can increase their oxidation state.

3. Some compounds containing simple ions such as I – , S 2– .

4. Compounds containing complex ions (S 4+ O 3) 2–, (НР 3+ O 3) 2–, in which elements can, by donating electrons, increase their positive oxidation state.

In laboratory practice, the following oxidizing agents are most often used:

potassium permanganate (KMnO 4);

potassium dichromate (K 2 Cr 2 O 7);

nitric acid (HNO 3);

concentrated sulfuric acid (H 2 SO 4);

hydrogen peroxide (H 2 O 2);

oxides of manganese (IV) and lead (IV) (MnO 2, PbO 2);

molten potassium nitrate (KNO 3) and melts of some other nitrates.

Reducing agents used in laboratory practice include:

magnesium (Mg), aluminum (Al) and other active metals;
hydrogen (H 2) and carbon (C);
potassium iodide (KI);
sodium sulfide (Na 2 S) and hydrogen sulfide (H 2 S);
sodium sulfite (Na 2 SO 3);
tin chloride (SnCl 2).

7.2.3. Classification of redox reactions

Redox reactions are usually divided into three types: intermolecular, intramolecular, and disproportionation reactions (self-oxidation-self-reduction).

Intermolecular reactions occur with a change in the oxidation state of atoms that are found in different molecules. For example:

2 Al + Fe 2 O 3 Al 2 O 3 + 2 Fe,

C + 4 HNO 3(conc) = CO 2 + 4 NO 2 + 2 H 2 O.

TO intramolecular reactions These are reactions in which the oxidizing agent and the reducing agent are part of the same molecule, for example:

(NH 4) 2 Cr 2 O 7 N 2 + Cr 2 O 3 + 4 H 2 O,

2 KNO 3 2 KNO 2 + O 2 .

IN disproportionation reactions(self-oxidation-self-reduction) an atom (ion) of the same element is both an oxidizing agent and a reducing agent:

Cl 2 + 2 KOH KCl + KClO + H 2 O,

2 NO 2 + 2 NaOH = NaNO 2 + NaNO 3 + H 2 O.

7.2.4. Basic rules for composing redox reactions

The composition of redox reactions is carried out according to the steps presented in table. 7.2.

Table 7.2

Stages of compiling equations for redox reactions

	Action
	Determine the oxidizing agent and reducing agent.
	Identify the products of the redox reaction.
	Create an electron balance and use it to assign coefficients for substances that change their oxidation states.
	Arrange the coefficients for other substances that take part and are formed in the redox reaction.
	Check the correctness of the coefficients by counting the amount of substance of the atoms (usually hydrogen and oxygen) located on the left and right sides of the reaction equation.

Let's consider the rules for composing redox reactions using the example of the interaction of potassium sulfite with potassium permanganate in an acidic environment:

1. Determination of oxidizing agent and reducing agent

Manganese, which is in the highest oxidation state, cannot give up electrons. Mn 7+ will accept electrons, i.e. is an oxidizing agent.

The S 4+ ion can donate two electrons and go into S 6+, i.e. is a reducing agent. Thus, in the reaction under consideration, K 2 SO 3 is a reducing agent, and KMnO 4 is an oxidizing agent.

2. Establishment of reaction products

K2SO3 + KMnO4 + H2SO4?

By donating two electrons to an electron, S 4+ becomes S 6+. Potassium sulfite (K 2 SO 3) thus turns into sulfate (K 2 SO 4). In an acidic environment, Mn 7+ accepts 5 electrons and in a solution of sulfuric acid (medium) forms manganese sulfate (MnSO 4). As a result of this reaction, additional molecules of potassium sulfate are also formed (due to the potassium ions included in the permanganate), as well as water molecules. Thus, the reaction under consideration will be written as:

K 2 SO 3 + KMnO 4 + H 2 SO 4 = K 2 SO 4 + MnSO 4 + H 2 O.

3. Compiling electron balance

To compile an electron balance, it is necessary to indicate those oxidation states that change in the reaction under consideration:

K 2 S 4+ O 3 + KMn 7+ O 4 + H 2 SO 4 = K 2 S 6+ O 4 + Mn 2+ SO 4 + H 2 O.

Mn 7+ + 5 e = Mn 2+ ;

S 4+ – 2 e = S 6+.

The number of electrons given up by the reducing agent must be equal to the number of electrons accepted by the oxidizing agent. Therefore, two Mn 7+ and five S 4+ must participate in the reaction:

Mn 7+ + 5 e = Mn 2+ 2,

S 4+ – 2 e = S 6+ 5.

Thus, the number of electrons given up by the reducing agent (10) will be equal to the number of electrons accepted by the oxidizing agent (10).

4. Arrangement of coefficients in the reaction equation

In accordance with the balance of electrons, it is necessary to put a coefficient of 5 in front of K 2 SO 3, and 2 in front of KMnO 4. On the right side, in front of potassium sulfate we set a coefficient of 6, since one molecule is added to the five molecules of K 2 SO 4 formed during the oxidation of potassium sulfite K 2 SO 4 as a result of the binding of potassium ions included in the permanganate. Since the reaction involves two permanganate molecules, on the right side are also formed two manganese sulfate molecules. To bind the reaction products (potassium and manganese ions included in the permanganate), it is necessary three molecules of sulfuric acid, therefore, as a result of the reaction, three water molecules. Finally we get:

5 K 2 SO 3 + 2 KMnO 4 + 3 H 2 SO 4 = 6 K 2 SO 4 + 2 MnSO 4 + 3 H 2 O.

5. Checking the correctness of the coefficients in the reaction equation

The number of oxygen atoms on the left side of the reaction equation is:

5 3 + 2 4 + 3 4 = 35.

On the right side this number will be:

6 4 + 2 4 + 3 1 = 35.

The number of hydrogen atoms on the left side of the reaction equation is six and corresponds to the number of these atoms on the right side of the reaction equation.

7.2.5. Examples of redox reactions involving typical oxidizing and reducing agents

7.2.5.1. Intermolecular oxidation-reduction reactions

Below, as examples, we consider redox reactions involving potassium permanganate, potassium dichromate, hydrogen peroxide, potassium nitrite, potassium iodide and potassium sulfide. Redox reactions involving other typical oxidizing and reducing agents are discussed in the second part of the manual (“Inorganic chemistry”).

Redox reactions involving potassium permanganate

Depending on the environment (acidic, neutral, alkaline), potassium permanganate, acting as an oxidizing agent, gives various reduction products, Fig. 7.1.

Rice. 7.1. Formation of potassium permanganate reduction products in various media

Below are the reactions of KMnO 4 with potassium sulfide as a reducing agent in various environments, illustrating the scheme, Fig. 7.1. In these reactions, the product of sulfide ion oxidation is free sulfur. In an alkaline environment, KOH molecules do not take part in the reaction, but only determine the product of the reduction of potassium permanganate.

5 K 2 S + 2 KMnO 4 + 8 H 2 SO 4 = 5 S + 2 MnSO 4 + 6 K 2 SO 4 + 8 H 2 O,

3 K 2 S + 2 KMnO 4 + 4 H 2 O 2 MnO 2 + 3 S + 8 KOH,

K 2 S + 2 KMnO 4 – (KOH) 2 K 2 MnO 4 + S.

Redox reactions involving potassium dichromate

In an acidic environment, potassium dichromate is a strong oxidizing agent. A mixture of K 2 Cr 2 O 7 and concentrated H 2 SO 4 (chromium) is widely used in laboratory practice as an oxidizing agent. Interacting with a reducing agent, one molecule of potassium dichromate accepts six electrons, forming trivalent chromium compounds:

6 FeSO 4 +K 2 Cr 2 O 7 +7 H 2 SO 4 = 3 Fe 2 (SO 4) 3 +Cr 2 (SO 4) 3 +K 2 SO 4 +7 H 2 O;

6 KI + K 2 Cr 2 O 7 + 7 H 2 SO 4 = 3 I 2 + Cr 2 (SO 4) 3 + 4 K 2 SO 4 + 7 H 2 O.

Redox reactions involving hydrogen peroxide and potassium nitrite

Hydrogen peroxide and potassium nitrite exhibit predominantly oxidizing properties:

H 2 S + H 2 O 2 = S + 2 H 2 O,

2 KI + 2 KNO 2 + 2 H 2 SO 4 = I 2 + 2 K 2 SO 4 + H 2 O,

However, when interacting with strong oxidizing agents (such as, for example, KMnO 4), hydrogen peroxide and potassium nitrite act as reducing agents:

5 H 2 O 2 + 2 KMnO 4 + 3 H 2 SO 4 = 5 O 2 + 2 MnSO 4 + K 2 SO 4 + 8 H 2 O,

5 KNO 2 + 2 KMnO 4 + 3 H 2 SO 4 = 5 KNO 3 + 2 MnSO 4 + K 2 SO 4 + 3 H 2 O.

It should be noted that hydrogen peroxide, depending on the environment, is reduced according to the scheme, Fig. 7.2.

Rice. 7.2. Possible hydrogen peroxide reduction products

In this case, as a result of the reactions, water or hydroxide ions are formed:

2 FeSO 4 + H 2 O 2 + H 2 SO 4 = Fe 2 (SO 4) 3 + 2 H 2 O,

2 KI + H 2 O 2 = I 2 + 2 KOH.

7.2.5.2. Intramolecular oxidation-reduction reactions

Intramolecular redox reactions usually occur when substances whose molecules contain a reducing agent and an oxidizing agent are heated. Examples of intramolecular reduction-oxidation reactions are the processes of thermal decomposition of nitrates and potassium permanganate:

2 NaNO 3 2 NaNO 2 + O 2,

2 Cu(NO 3) 2 2 CuO + 4 NO 2 + O 2,

Hg(NO 3) 2 Hg + NO 2 + O 2,

2 KMnO 4 K 2 MnO 4 + MnO 2 + O 2.

7.2.5.3. Disproportionation reactions

As noted above, in disproportionation reactions the same atom (ion) is both an oxidizing agent and a reducing agent. Let us consider the process of composing this type of reaction using the example of the interaction of sulfur with alkali.

Characteristic oxidation states of sulfur: – 2, 0, +4 and +6. Acting as a reducing agent, elemental sulfur donates 4 electrons:

S o – 4e = S 4+.

Sulfur – The oxidizing agent accepts two electrons:

S o + 2е = S 2– .

Thus, as a result of the reaction of sulfur disproportionation, compounds are formed whose oxidation states of the element are – 2 and right +4:

3 S + 6 KOH = 2 K 2 S + K 2 SO 3 + 3 H 2 O.

When nitrogen oxide (IV) is disproportioned in alkali, nitrite and nitrate are obtained - compounds in which the oxidation states of nitrogen are +3 and +5, respectively:

2 N 4+ O 2 + 2 KOH = KN 3+ O 2 + KN 5+ O 3 + H 2 O,

Disproportionation of chlorine in a cold alkali solution leads to the formation of hypochlorite, and in a hot alkali solution - chlorate:

Cl 0 2 + 2 KOH = KCl – + KCl + O + H 2 O,

Cl 0 2 + 6 KOH 5 KCl – + KCl 5+ O 3 + 3H 2 O.

7.3. Electrolysis

The redox process that occurs in solutions or melts when a direct electric current is passed through them is called electrolysis. In this case, oxidation of anions occurs at the positive electrode (anode). Cations are reduced at the negative electrode (cathode).

2 Na 2 CO 3 4 Na + O 2 + 2CO 2 .

During the electrolysis of aqueous solutions of electrolytes, along with transformations of the dissolved substance, electrochemical processes can occur with the participation of hydrogen ions and hydroxide ions of water:

cathode (–): 2 Н + + 2е = Н 2,

anode (+): 4 OH – – 4e = O 2 + 2 H 2 O.

In this case, the reduction process at the cathode occurs as follows:

1. Cations of active metals (up to Al 3+ inclusive) are not reduced at the cathode; hydrogen is reduced instead.

2. Metal cations located in the series of standard electrode potentials (in the voltage series) to the right of hydrogen are reduced to free metals at the cathode during electrolysis.

3. Metal cations located between Al 3+ and H + are reduced at the cathode simultaneously with the hydrogen cation.

The processes occurring in aqueous solutions at the anode depend on the substance from which the anode is made. There are insoluble anodes ( inert) and soluble ( active). Graphite or platinum is used as the material of inert anodes. Soluble anodes are made from copper, zinc and other metals.

During the electrolysis of solutions with an inert anode, the following products can be formed:

1. When halide ions are oxidized, free halogens are released.

2. During the electrolysis of solutions containing the anions SO 2 2–, NO 3 –, PO 4 3–, oxygen is released, i.e. It is not these ions that are oxidized at the anode, but water molecules.

Taking into account the above rules, let us consider, as an example, the electrolysis of aqueous solutions of NaCl, CuSO 4 and KOH with inert electrodes.

1). In solution, sodium chloride dissociates into ions.

DEFINITION

Chemical reaction are called transformations of substances in which a change in their composition and (or) structure occurs.

Most often, chemical reactions are understood as the process of converting starting substances (reagents) into final substances (products).

Chemical reactions are written using chemical equations containing the formulas of the starting substances and reaction products. According to the law of conservation of mass, the number of atoms of each element on the left and right sides of a chemical equation is the same. Typically, the formulas of the starting substances are written on the left side of the equation, and the formulas of the products on the right. The equality of the number of atoms of each element on the left and right sides of the equation is achieved by placing integer stoichiometric coefficients in front of the formulas of substances.

Chemical equations may contain additional information about the characteristics of the reaction: temperature, pressure, radiation, etc., which is indicated by the corresponding symbol above (or “below”) the equal sign.

All chemical reactions can be grouped into several classes, which have certain characteristics.

Classification of chemical reactions according to the number and composition of starting and resulting substances

According to this classification, chemical reactions are divided into reactions of connection, decomposition, substitution, and exchange.

As a result compound reactions from two or more (complex or simple) substances one new substance is formed. In general, the equation for such a chemical reaction will look like this:

For example:

CaCO 3 + CO 2 + H 2 O = Ca(HCO 3) 2

SO 3 + H 2 O = H 2 SO 4

2Mg + O 2 = 2MgO.

2FeCl 2 + Cl 2 = 2FeCl 3

The reactions of the compound are in most cases exothermic, i.e. proceed with the release of heat. If simple substances are involved in the reaction, then such reactions are most often redox reactions (ORR), i.e. occur with changes in the oxidation states of elements. It is impossible to say unambiguously whether the reaction of a compound between complex substances will be classified as ORR.

Reactions that result in the formation of several other new substances (complex or simple) from one complex substance are classified as decomposition reactions. In general, the equation for the chemical reaction of decomposition will look like this:

For example:

CaCO 3 CaO + CO 2 (1)

2H 2 O = 2H 2 + O 2 (2)

CuSO 4 × 5H 2 O = CuSO 4 + 5H 2 O (3)

Cu(OH) 2 = CuO + H 2 O (4)

H 2 SiO 3 = SiO 2 + H 2 O (5)

2SO 3 =2SO 2 + O 2 (6)

(NH 4) 2 Cr 2 O 7 = Cr 2 O 3 + N 2 +4H 2 O (7)

Most decomposition reactions occur when heated (1,4,5). Possible decomposition under the influence of electric current (2). The decomposition of crystalline hydrates, acids, bases and salts of oxygen-containing acids (1, 3, 4, 5, 7) occurs without changing the oxidation states of the elements, i.e. these reactions are not related to ODD. ORR decomposition reactions include the decomposition of oxides, acids and salts formed by elements in higher oxidation states (6).

Decomposition reactions are also found in organic chemistry, but under other names - cracking (8), dehydrogenation (9):

C 18 H 38 = C 9 H 18 + C 9 H 20 (8)

C 4 H 10 = C 4 H 6 + 2H 2 (9)

At substitution reactions a simple substance interacts with a complex substance, forming a new simple and a new complex substance. In general, the equation for a chemical substitution reaction will look like this:

For example:

2Al + Fe 2 O 3 = 2Fe + Al 2 O 3 (1)

Zn + 2HCl = ZnСl 2 + H 2 (2)

2KBr + Cl 2 = 2KCl + Br 2 (3)

2КlO 3 + l 2 = 2KlO 3 + Сl 2 (4)

CaCO 3 + SiO 2 = CaSiO 3 + CO 2 (5)

Ca 3 (PO 4) 2 + 3SiO 2 = 3СаSiO 3 + P 2 O 5 (6)

CH 4 + Cl 2 = CH 3 Cl + HCl (7)

Most substitution reactions are redox (1 – 4, 7). Examples of decomposition reactions in which no change in oxidation states occurs are few (5, 6).

Exchange reactions are reactions that occur between complex substances in which they exchange their constituent parts. Typically this term is used for reactions involving ions in aqueous solution. In general, the equation for a chemical exchange reaction will look like this:

AB + CD = AD + CB

For example:

CuO + 2HCl = CuCl 2 + H 2 O (1)

NaOH + HCl = NaCl + H 2 O (2)

NaHCO 3 + HCl = NaCl + H 2 O + CO 2 (3)

AgNO 3 + KBr = AgBr ↓ + KNO 3 (4)

CrCl 3 + ZNaON = Cr(OH) 3 ↓+ ZNaCl (5)

Exchange reactions are not redox. A special case of these exchange reactions is the neutralization reaction (the reaction of acids with alkalis) (2). Exchange reactions proceed in the direction where at least one of the substances is removed from the reaction sphere in the form of a gaseous substance (3), a precipitate (4, 5) or a poorly dissociating compound, most often water (1, 2).

Classification of chemical reactions according to changes in oxidation states

Depending on the change in the oxidation states of the elements that make up the reagents and reaction products, all chemical reactions are divided into redox reactions (1, 2) and those occurring without changing the oxidation state (3, 4).

2Mg + CO 2 = 2MgO + C (1)

Mg 0 – 2e = Mg 2+ (reducing agent)

C 4+ + 4e = C 0 (oxidizing agent)

FeS 2 + 8HNO 3 (conc) = Fe(NO 3) 3 + 5NO + 2H 2 SO 4 + 2H 2 O (2)

Fe 2+ -e = Fe 3+ (reducing agent)

N 5+ +3e = N 2+ (oxidizing agent)

AgNO 3 +HCl = AgCl ↓ + HNO 3 (3)

Ca(OH) 2 + H 2 SO 4 = CaSO 4 ↓ + H 2 O (4)

Classification of chemical reactions by thermal effect

Depending on whether heat (energy) is released or absorbed during the reaction, all chemical reactions are conventionally divided into exothermic (1, 2) and endothermic (3), respectively. The amount of heat (energy) released or absorbed during a reaction is called the thermal effect of the reaction. If the equation indicates the amount of heat released or absorbed, then such equations are called thermochemical.

N 2 + 3H 2 = 2NH 3 +46.2 kJ (1)

2Mg + O 2 = 2MgO + 602.5 kJ (2)

N 2 + O 2 = 2NO – 90.4 kJ (3)

Classification of chemical reactions according to the direction of the reaction

Based on the direction of the reaction, a distinction is made between reversible (chemical processes whose products are capable of reacting with each other under the same conditions in which they were obtained to form the starting substances) and irreversible (chemical processes whose products are not able to react with each other to form the starting substances). ).

For reversible reactions, the equation in general form is usually written as follows:

A + B ↔ AB

For example:

CH 3 COOH + C 2 H 5 OH ↔ H 3 COOC 2 H 5 + H 2 O

Examples of irreversible reactions include the following reactions:

2КlО 3 → 2Кl + ЗО 2

C 6 H 12 O 6 + 6O 2 → 6CO 2 + 6H 2 O

Evidence of the irreversibility of a reaction can be the release of a gaseous substance, a precipitate, or a poorly dissociating compound, most often water, as reaction products.

Classification of chemical reactions according to the presence of a catalyst

From this point of view, catalytic and non-catalytic reactions are distinguished.

A catalyst is a substance that speeds up the progress of a chemical reaction. Reactions that occur with the participation of catalysts are called catalytic. Some reactions cannot take place at all without the presence of a catalyst:

2H 2 O 2 = 2H 2 O + O 2 (MnO 2 catalyst)

Often one of the reaction products serves as a catalyst that accelerates this reaction (autocatalytic reactions):

MeO+ 2HF = MeF 2 + H 2 O, where Me is a metal.

Examples of problem solving

EXAMPLE 1

When a compound reacts from several reacting substances of relatively simple composition, one substance of a more complex composition is obtained:

As a rule, these reactions are accompanied by the release of heat, i.e. lead to the formation of more stable and less energy-rich compounds.

Reactions of compounds of simple substances are always redox in nature. Compound reactions occurring between complex substances can occur without a change in valence:

CaCO 3 + CO 2 + H 2 O = Ca(HCO 3) 2,

and also be classified as redox:

2FeCl 2 + Cl 2 = 2FeCl 3.

2. Decomposition reactions

Decomposition reactions lead to the formation of several compounds from one complex substance:

A = B + C + D.

The decomposition products of a complex substance can be both simple and complex substances.

Of the decomposition reactions that occur without changing the valence states, noteworthy is the decomposition of crystalline hydrates, bases, acids and salts of oxygen-containing acids:


		CuSO 4 + 5H 2 O


		2H 2 O + 4NO 2 O + O 2 O.

2AgNO3 = 2Ag + 2NO2 + O2, (NH4)2Cr2O7 = Cr2O3 + N2 + 4H2O.

Redox decomposition reactions are especially characteristic for nitric acid salts.

Decomposition reactions in organic chemistry are called cracking:

C 18 H 38 = C 9 H 18 + C 9 H 20,

or dehydrogenation

C4H10 = C4H6 + 2H2.

3. Substitution reactions

In substitution reactions, usually a simple substance reacts with a complex one, forming another simple substance and another complex one:

A + BC = AB + C.

These reactions overwhelmingly belong to redox reactions:

2Al + Fe 2 O 3 = 2Fe + Al 2 O 3,

Zn + 2HCl = ZnСl 2 + H 2,

2KBr + Cl 2 = 2KCl + Br 2,

2KlO 3 + l 2 = 2KlO 3 + Cl 2.

Examples of substitution reactions that are not accompanied by a change in the valence states of atoms are extremely few. It should be noted the reaction of silicon dioxide with salts of oxygen-containing acids, which correspond to gaseous or volatile anhydrides:

CaCO 3 + SiO 2 = CaSiO 3 + CO 2,

Ca 3 (PO 4) 2 + 3SiO 2 \u003d 3СаSiO 3 + P 2 O 5,

Sometimes these reactions are considered as exchange reactions:

CH 4 + Cl 2 = CH 3 Cl + HCl.

4. Exchange reactions

Exchange reactions are reactions between two compounds that exchange their constituents with each other:

AB + CD = AD + CB.

If redox processes occur during substitution reactions, then exchange reactions always occur without changing the valence state of the atoms. This is the most common group of reactions between complex substances - oxides, bases, acids and salts:

ZnO + H 2 SO 4 = ZnSO 4 + H 2 O,

AgNO 3 + KBr = AgBr + KNO 3,

CrCl 3 + ZNaON = Cr(OH) 3 + ZNaCl.

A special case of these exchange reactions is the neutralization reaction:

HCl + KOH = KCl + H 2 O.

Typically, these reactions obey the laws of chemical equilibrium and proceed in the direction where at least one of the substances is removed from the reaction sphere in the form of a gaseous, volatile substance, precipitate or low-dissociating (for solutions) compound:

NaHCO 3 + HCl = NaCl + H 2 O + CO 2,

Ca(HCO 3) 2 + Ca(OH) 2 = 2CaCO 3 ↓ + 2H 2 O,

CH 3 COONa + H 3 PO 4 = CH 3 COOH + NaH 2 PO 4.

During chemical reactions, one substance turns into another (not to be confused with nuclear reactions, in which one chemical element is converted into another).

Any chemical reaction is described by a chemical equation:

Reactants → Reaction products

The arrow indicates the direction of the reaction.

For example:

In this reaction, methane (CH 4) reacts with oxygen (O 2), resulting in the formation of carbon dioxide (CO 2) and water (H 2 O), or more precisely, water vapor. This is exactly the reaction that happens in your kitchen when you light a gas burner. The equation should be read like this: One molecule of methane gas reacts with two molecules of oxygen gas to produce one molecule of carbon dioxide and two molecules of water (water vapor).

The numbers placed before the components of a chemical reaction are called reaction coefficients.

Chemical reactions happen endothermic(with energy absorption) and exothermic(with energy release). Methane combustion is a typical example of an exothermic reaction.

There are several types of chemical reactions. The most common:

connection reactions;
decomposition reactions;
single replacement reactions;
double displacement reactions;
oxidation reactions;
redox reactions.

Compound reactions

In compound reactions, at least two elements form one product:

2Na (t) + Cl 2 (g) → 2NaCl (t)- formation of table salt.

Attention should be paid to an essential nuance of compound reactions: depending on the conditions of the reaction or the proportions of the reagents entering the reaction, its result may be different products. For example, under normal combustion conditions of coal, carbon dioxide is produced:
C (t) + O 2 (g) → CO 2 (g)

If the amount of oxygen is insufficient, then deadly carbon monoxide is formed:
2C (t) + O 2 (g) → 2CO (g)

Decomposition reactions

These reactions are, as it were, essentially opposite to the reactions of the compound. As a result of the decomposition reaction, the substance breaks down into two (3, 4...) simpler elements (compounds):

2H 2 O (l) → 2H 2 (g) + O 2 (g)- water decomposition
2H 2 O 2 (l) → 2H 2 (g) O + O 2 (g)- decomposition of hydrogen peroxide

Single displacement reactions

As a result of single substitution reactions, a more active element replaces a less active one in a compound:

Zn (s) + CuSO 4 (solution) → ZnSO 4 (solution) + Cu (s)

Zinc in a copper sulfate solution displaces the less active copper, resulting in the formation of a zinc sulfate solution.

The degree of activity of metals in increasing order of activity:

The most active are alkali and alkaline earth metals

The ionic equation for the above reaction will be:

Zn (t) + Cu 2+ + SO 4 2- → Zn 2+ + SO 4 2- + Cu (t)

The ionic bond CuSO 4, when dissolved in water, breaks down into a copper cation (charge 2+) and a sulfate anion (charge 2-). As a result of the substitution reaction, a zinc cation is formed (which has the same charge as the copper cation: 2-). Please note that the sulfate anion is present on both sides of the equation, i.e., according to all the rules of mathematics, it can be reduced. The result is an ion-molecular equation:

Zn (t) + Cu 2+ → Zn 2+ + Cu (t)

Double displacement reactions

In double substitution reactions, two electrons are already replaced. Such reactions are also called exchange reactions. Such reactions take place in solution with the formation of:

insoluble solid (precipitation reaction);
water (neutralization reaction).

Precipitation reactions

When a solution of silver nitrate (salt) is mixed with a solution of sodium chloride, silver chloride is formed:

Molecular equation: KCl (solution) + AgNO 3 (p-p) → AgCl (s) + KNO 3 (p-p)

Ionic equation: K + + Cl - + Ag + + NO 3 - → AgCl (t) + K + + NO 3 -

Molecular ionic equation: Cl - + Ag + → AgCl (s)

If a compound is soluble, it will be present in solution in ionic form. If the compound is insoluble, it will precipitate to form a solid.

Neutralization reactions

These are reactions between acids and bases that result in the formation of water molecules.

For example, the reaction of mixing a solution of sulfuric acid and a solution of sodium hydroxide (lye):

Molecular equation: H 2 SO 4 (p-p) + 2NaOH (p-p) → Na 2 SO 4 (p-p) + 2H 2 O (l)

Ionic equation: 2H + + SO 4 2- + 2Na + + 2OH - → 2Na + + SO 4 2- + 2H 2 O (l)

Molecular ionic equation: 2H + + 2OH - → 2H 2 O (l) or H + + OH - → H 2 O (l)

Oxidation reactions

These are reactions of interaction of substances with gaseous oxygen in the air, during which, as a rule, a large amount of energy is released in the form of heat and light. A typical oxidation reaction is combustion. At the very beginning of this page is the reaction between methane and oxygen:

CH 4 (g) + 2O 2 (g) → CO 2 (g) + 2H 2 O (g)

Methane belongs to hydrocarbons (compounds of carbon and hydrogen). When a hydrocarbon reacts with oxygen, a lot of thermal energy is released.

Redox reactions

These are reactions in which electrons are exchanged between reactant atoms. The reactions discussed above are also redox reactions:

2Na + Cl 2 → 2NaCl - compound reaction
CH 4 + 2O 2 → CO 2 + 2H 2 O - oxidation reaction
Zn + CuSO 4 → ZnSO 4 + Cu - single substitution reaction

Redox reactions with a large number of examples of solving equations using the electron balance method and the half-reaction method are described in as much detail as possible in the section