All possible chemical reactions. G

Compound reactions (formation of one complex substance from several simple or complex substances) A ​​+ B = AB

Decomposition reactions (decomposition of one complex substance into several simple or complex substances) AB = A + B

Substitution reactions (between simple and complex substances, in which atoms of a simple substance replace atoms of one of the elements in a complex substance): AB + C = AC + B

Exchange reactions (between two complex substances in which substances exchange their components) AB + SD = AD + SV

1. Indicate the correct definition of the 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. Reaction between base and 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 the 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.

4. Indicate the signs of the exchange reaction:

• A. Water formation;

• B. Gas formation only;

• B. Precipitation only;

• D. Precipitation, gas formation, or weak electrolyte formation.

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.

6. What type of reactions is the interaction of salts with acids or bases:

• A. Substitution reactions;

• B. Decomposition reactions;

• B. Exchange reactions;

• D. Compound reactions.

• 7. Substances whose formulas are KNO3 FeCl2, Na2SO4 are called:

• A) salts; B) reasons; B) acids; D) oxides.

• 8 . Substances whose formulas are HNO3, HCl, H2SO4 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 NO2, Fe2O3, Na2O are called:

• A) salts; B) acids; B) reasons; D) oxides.

• 11 . Specify the metals that form alkalis:

• Cu, Fe, Na, K, Zn, Li.

Answers:

• Na, K, Li.

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 view 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 due to exposure 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 components. 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

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), metal melts, iodine sublimes, water turns into ice or steam, etc., as well as nuclear reactions, as a result of which atoms of other elements are formed from atoms of some 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. By sign 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. Based on 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. (mechanism of chemical reactions occurring with the participation organic compounds, 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

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 reacting substances are called redox.

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 Arabic numeral above the element symbol with a “+” or “–” sign. For example, Al 3+, S 2–.

To find oxidation states, they are guided by the following rules:

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).

Maximum negative oxidation state of elements in the main subgroups 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 groups 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 such simple ions as, for example, 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(auto-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

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

Located in highest degree oxidation, manganese 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 (chrompic) 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 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.