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The concept of “compound reaction” is an antonym of the concept of “decomposition reaction”. Try, using the technique of contrast, to define the concept of “compound reaction.” Right! You have the following formulation.

Let's consider this type of reaction using another, new for you, form of recording chemical processes - the so-called chains of transitions, or transformations. For example, circuit

shows the transformation of phosphorus into phosphorus oxide (V) P 2 O 5, which, in turn, is then converted into phosphoric acid H 3 PO 4.

The number of arrows in the diagram of the transformation of substances corresponds to the minimum number of chemical transformations - chemical reactions. In the example under consideration, these are two chemical processes.

1st process. Obtaining phosphorus oxide (V) P 2 O 5 from phosphorus. Obviously, this is a reaction between phosphorus and oxygen.

Let's put some red phosphorus in a burning spoon and set it on fire. Phosphorus burns with a bright flame producing white smoke consisting of small particles of phosphorus (V) oxide:

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

2nd process. Let's add a spoonful of burning phosphorus into the flask. It is filled with thick smoke from phosphorus (V) oxide. Take the spoon out of the flask, pour water into the flask and shake the contents, after closing the neck of the flask with a stopper. The smoke gradually thins out, dissolves in the water and finally disappears completely. If you add a little litmus to the solution obtained in the flask, it will turn red, which is evidence of the formation of phosphoric acid:

R 2 O 5 + ZN 2 O = 2H 3 PO 4.

The reactions that are carried out to carry out the transitions under consideration occur without the participation of a catalyst, which is why they are called non-catalytic. The reactions discussed above proceed only in one direction, that is, they are irreversible.

Let us analyze how many and what substances entered into the reactions discussed above and how many and what substances were formed in them. In the first reaction, one complex substance was formed from two simple substances, and in the second, from two complex substances, each of which consists of two elements, one complex substance was formed, consisting of three elements.

One complex substance can also be formed as a result of the reaction of combining a complex and a simple substance. For example, in the production of sulfuric acid from sulfur oxide (IV) sulfur oxide (VI) is obtained:

This reaction proceeds both in the forward direction, i.e., with the formation of a reaction product, and in the reverse direction, i.e., the decomposition of the reaction product into the starting substances occurs, therefore, instead of the equal sign, they put the reversibility sign.

This reaction involves a catalyst - vanadium (V) oxide V 2 O 5, which is indicated above the reversibility sign:

A complex substance can also be obtained by combining three substances. For example, nitric acid is produced by a reaction whose scheme is:

NO 2 + H 2 O + O 2 → HNO 3.

Let's consider how to select coefficients to equalize the scheme of this chemical reaction.

There is no need to equalize the number of nitrogen atoms: there is one nitrogen atom in both the left and right parts of the diagram. Let's equalize the number of hydrogen atoms - before the acid formula we write the coefficient 2:

NO 2 + H 2 O + O 2 → 2HNO 3.

but in this case the equality of the number of nitrogen atoms will be violated - one nitrogen atom remains on the left side, and there are two on the right side. Let's write the coefficient 2 before the formula of nitric oxide (IV):

2NO 2 + H 2 O + O 2 → 2HNO 3.

Let's count the number of oxygen atoms: there are seven on the left side of the reaction diagram, and six on the right side. To equalize the number of oxygen atoms (six atoms in each part of the equation), remember that before the formulas of simple substances you can write the fractional coefficient 1/2:

2NO 2 + H 2 O + 1/2O 2 → 2HNO 3.

Let's make the coefficients integers. To do this, we rewrite the equation by doubling the coefficients:

4NO 2 + 2H 2 O + O 2 → 4HNO 3.

It should be noted that almost all reactions of the compound are exothermic reactions.

Laboratory experiment No. 15
Calcination of copper in the flame of an alcohol lamp

Examine the copper wire (plate) given to you and describe its appearance. Heat the wire, holding it with crucible tongs, in the upper part of the flame of an alcohol lamp for 1 minute. Describe the reaction conditions. Describe a sign that indicates a chemical reaction has occurred. Write an equation for the reaction that took place. Name the starting materials and products of the reaction.

Explain whether the mass of the copper wire (plate) changed after the end of the experiment. Justify your answer using your knowledge of the law of conservation of mass of substances.

Key words and phrases

1. Combination reactions are antonyms of decomposition reactions.
2. Catalytic (including enzymatic) and non-catalytic reactions.
3. Chains of transitions or transformations.
4. Reversible and irreversible reactions.

Work with computer

1. Refer to the electronic application. Study the lesson material and complete the assigned tasks.
2. Find email addresses on the Internet that can serve as additional sources that reveal the content of keywords and phrases in the paragraph. Offer your help to the teacher in preparing a new lesson - make a report on the key words and phrases of the next paragraph.

Questions and tasks

The chemical properties of substances are revealed in a variety of chemical reactions.

Transformations of substances accompanied by changes in their composition and (or) structure are called chemical reactions. The following definition is often found: chemical reaction is the process of converting starting substances (reagents) into final substances (products).

Chemical reactions are written using chemical equations and diagrams containing the formulas of the starting substances and reaction products. In chemical equations, unlike diagrams, the number of atoms of each element is the same on the left and right sides, which reflects the law of conservation of mass.

On the left side of the equation the formulas of the starting substances (reagents) are written, on the right side - the substances obtained as a result of the chemical reaction (reaction products, final substances). The equal sign connecting the left and right sides indicates that the total number of atoms of the substances involved in the reaction remains constant. This is achieved by placing integer stoichiometric coefficients in front of the formulas, showing the quantitative relationships between the reactants and reaction products.

Chemical equations may contain additional information about the characteristics of the reaction. If a chemical reaction occurs under the influence of external influences (temperature, pressure, radiation, etc.), this is indicated by the appropriate symbol, usually above (or “below”) the equal sign.

A huge number of chemical reactions can be grouped into several types of reactions, which have very specific characteristics.

As classification characteristics the following can be selected:

1. The number and composition of starting substances and reaction products.

2. Physical state of the reagents and reaction products.

3. The number of phases in which the reaction participants are located.

4. The nature of the transferred particles.

5. Possibility of the reaction occurring in forward and reverse directions.

6. The sign of the thermal effect divides all reactions into: exothermic reactions occurring with exo-effect - release of energy in the form of heat (Q>0, ∆H<0):

C + O 2 = CO 2 + Q

And endothermic reactions occurring with the endo effect - the absorption of energy in the form of heat (Q<0, ∆H >0):

N 2 + O 2 = 2NO - Q.

Such reactions are referred to as thermochemical.

Let's take a closer look at each type of reaction.

Classification according to the number and composition of reagents and final substances

1. Compound reactions

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:

	t o
4HNO3	=	2H 2 O + 4NO 2 O + O 2 O.

2AgNO3 = 2Ag + 2NO2 + O2,
(NH 4) 2 Cr 2 O 7 = Cr 2 O 3 + N 2 + 4H 2 O.

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,

2КlO 3 + l 2 = 2KlO 3 + Сl 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 neutralization reactions:

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.

5. Transfer reactions.

In transfer reactions, an atom or group of atoms moves from one structural unit to another:

AB + BC = A + B 2 C,

A 2 B + 2CB 2 = DIA 2 + DIA 3.

For example:

2AgCl + SnCl 2 = 2Ag + SnCl 4,

H 2 O + 2NO 2 = HNO 2 + HNO 3.

Classification of reactions according to phase characteristics

Depending on the state of aggregation of the reacting substances, the following reactions are distinguished:

1. Gas reactions

H2+Cl2		2HCl.

2. Reactions in solutions

NaOH(solution) + HCl(p-p) = NaCl(p-p) + H 2 O(l)

3. Reactions between solids

	t o
CaO(tv) + SiO 2 (tv)	=	CaSiO 3 (sol)

Classification of reactions according to the number of phases.

A phase is understood as a set of homogeneous parts of a system with the same physical and chemical properties and separated from each other by an interface.

From this point of view, the entire variety of reactions can be divided into two classes:

1. Homogeneous (single-phase) reactions. These include reactions occurring in the gas phase and a number of reactions occurring in solutions.

2. Heterogeneous (multiphase) reactions. These include reactions in which the reactants and reaction products are in different phases. For example:

gas-liquid-phase reactions

CO 2 (g) + NaOH(p-p) = NaHCO 3 (p-p).

gas-solid-phase reactions

CO 2 (g) + CaO (tv) = CaCO 3 (tv).

liquid-solid-phase reactions

Na 2 SO 4 (solution) + BaCl 3 (solution) = BaSO 4 (tv) ↓ + 2NaCl (p-p).

liquid-gas-solid-phase reactions

Ca(HCO 3) 2 (solution) + H 2 SO 4 (solution) = CO 2 (r) + H 2 O (l) + CaSO 4 (sol)↓.

Classification of reactions according to the type of particles transferred

1. Protolytic reactions.

TO protolytic reactions include chemical processes, the essence of which is the transfer of a proton from one reacting substance to another.

This classification is based on the protolytic theory of acids and bases, according to which an acid is any substance that donates a proton, and a base is a substance that can accept a proton, for example:

Protolytic reactions include neutralization and hydrolysis reactions.

2. Redox reactions.

These include reactions in which reacting substances exchange electrons, thereby changing the oxidation states of the atoms of the elements that make up the reacting substances. For example:

Zn + 2H + → Zn 2 + + H 2,

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

The vast majority of chemical reactions are redox reactions; they play an extremely important role.

3. Ligand exchange reactions.

These include reactions during which the transfer of an electron pair occurs with the formation of a covalent bond via a donor-acceptor mechanism. For example:

Cu(NO 3) 2 + 4NH 3 = (NO 3) 2,

Fe + 5CO = ,

Al(OH) 3 + NaOH = .

A characteristic feature of ligand exchange reactions is that the formation of new compounds, called complexes, occurs without changing the oxidation state.

4. Reactions of atomic-molecular exchange.

This type of reaction includes many of the substitution reactions studied in organic chemistry that occur via a radical, electrophilic or nucleophilic mechanism.

Reversible and irreversible chemical reactions

Reversible chemical processes are those whose products are capable of reacting with each other under the same conditions in which they were obtained to form the starting substances.

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

Two oppositely directed arrows indicate that, under the same conditions, both forward and reverse reactions occur simultaneously, for example:

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

Irreversible chemical processes are those whose products are not able to react with each other to form the starting substances. Examples of irreversible reactions include the decomposition of Berthollet salt when heated:

2КlО 3 → 2Кl + ЗО 2,

or oxidation of glucose by atmospheric oxygen:

C 6 H 12 O 6 + 6 O 2 → 6 CO 2 + 6 H 2 O.

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.

“Physics Thermonuclear Reactions” - Thermonuclear reaction. Problem: Plasma is difficult to retain. A controlled thermonuclear reaction is an energetically favorable reaction. Details about the reaction. Physics presentation on the topic: Self-sustaining thermonuclear reactions occur in stars. What is a thermonuclear reaction? TOKAMAK (toroidal magnetic chamber with current).

“Types of chemical reactions” - All reactions are accompanied by thermal effects. Reversible reactions are chemical reactions that occur simultaneously in two opposite directions (forward and reverse) For example: 3H2 + N2? 2NH3 Laboratory work. How can we call the process that is taking place? Chemical reactions occur: when mixing or physical contact of reagents, spontaneously when heated with the participation of catalysts, the action of light, electric current, mechanical action, etc.

“Classification of reactions” - Endothermic reactions: P (red)<=>P (white). S (rhombic)<=>S (plastic). Classification of reactions. The overwhelming majority of such reactions are. Decomposition of potassium permanganate when heated: Lithium combustion reaction: Phosphorus allotropy: Calcium combustion reaction in air: Interesting reactions.

“Nuclear reactions” - Radioactive radiation has a detrimental effect on living cells. Nuclear reactions are accompanied by energy transformations. Biological action. Biological effects of radioactive radiation. The effect of radiation on humans. Thermonuclear reactions. Application of nuclear reactions. Nuclear reactor.

“Reactions of acids” - BaCL2 + H2SO4 = BaSO4 + 2HCL Ba2+ + SO42- = BaSO4. Acids. Answers. Classification of acids. Check yourself. Generalization. Typical acid reactions.

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.