Reactions without changing the oxidation states of atoms. Reactions without and with a change in the oxidation state. III. Types of chemical reactions by direction

A chemical reaction is a process by which starting substances are converted into reaction products. The substances obtained after the completion of the reaction are called products. They may differ from the original ones in structure, composition, or both.

Based on changes in composition, the following types are distinguished: chemical reactions:

• with a change in composition (most of them);
• without changing the composition (isomerization and conversion of one allotropic modification to another).

If the composition of a substance does not change as a result of the reaction, then its structure necessarily changes, for example: Cgraphite↔Salmaz

Let us consider in more detail the classification of chemical reactions that occur with a change in composition.

I. According to the number and composition of substances

Compound reactions

As a result of such chemical processes from several substances one is formed: A + B + ... = C

Can connect:

• simple substances: 2Na + S = Na2S;
• simple with complex: 2SO2 + O2 = 2SO3;
• two complex ones: CaO + H2O = Ca(OH)2.
• more than two substances: 4Fe + 3O2 + 6H2O = 4Fe(OH)3

Decomposition reactions

One substance in such reactions decomposes into several others: A=B+C+...

Products in this case can be:

• simple substances: 2NaCl = 2Na + Cl2
• simple and complex: 2KNO3 = 2KNO2 + O2
• two complex ones: CaCO3 = CaO + CO2
• more than two products: 2AgNO3 = 2Ag + O2 + 2NO2

Substitution reactions

Such reactions in which simple and complex substances react with each other, and atoms of a simple substance replace atoms of one of the elements in a complex one, are called substitution reactions. Schematically, the process of substitution of atoms can be shown as follows: A + BC = B + AC.

For example, CuSO4 + Fe = FeSO4 + Cu

Exchange reactions

This group includes reactions in which two complex substances change their parts: AB + CD = AD + CB. According to Berthollet's rule, irreversible occurrence of such reactions is possible if at least one of the products:

• precipitate (insoluble substance): 2NaOH + CuSO4 = Cu(OH)2 + Na2SO4;
• low dissociating substance: NaOH + HCl = NaCl + H2O;
• gas: NaOH + NH4Cl = NaCl + NH3 + H2O (first, ammonia hydrate NH3 H2O is formed, which upon receipt immediately decomposes into ammonia and water).

II. By thermal effect

1. Exothermic — processes occurring with the release of heat:
   C + O2 = CO2 +Q
2. Endothermic - reactions in which heat is absorbed:
   Cu(OH)2 = CuO + H2O – Q

III. Types of chemical reactions by direction

1. Reversible are reactions occurring at the same time in both the forward and reverse directions: N2+O2 ↔ 2NO
2. Irreversible processes proceed to completion, that is, until at least one of the reacting substances is completely consumed. Examples of irreversible exchange reactions were discussed above.

IV. According to the presence of a catalyst

V. According to the state of aggregation of substances

1. If all reactants are in the same states of aggregation, the reaction is called homogeneous. Such processes occur throughout the entire volume. For example: NaOH + HCl = NaCl + H2O
2. Heterogeneous are reactions between substances in different states of aggregation that occur at the interface. For example: Zn + 2HCl = ZnCl2 + H2

VI. Types of chemical reactions based on changes in the oxidation state of reacting substances

1. Redox (ORR) - reactions in which the oxidation states of the reacting substances change.
2. Reactions taking place without changing oxidation states reagents (BISO).

The processes of combustion and substitution are always redox. Exchange reactions occur without changing the oxidation states of substances. All other processes can be either OVR or BISO.

One of the main concepts is not organic chemistry is the concept of oxidation state (CO).

The oxidation state of an element in a compound is the formal charge of an element's atom, calculated from the assumption that valence electrons are transferred to atoms with higher relative electronegativity (REO) and all bonds in the compound molecule are ionic.

The oxidation state of the element E is indicated at the top of the element symbol with a “+” or “-” sign before the number.

The degree of oxidation of ions that actually exist in a solution or crystals coincides with their charge number and is indicated similarly with a “+” or “ ” sign after the number, for example, Ca 2+.

The Stock method is also used to indicate the oxidation state in Roman numerals after the element symbol: Mn (VII), Fe (III).

The question of the sign of the oxidation state of atoms in a molecule is resolved based on a comparison of the electronegativities of the interconnected atoms that form the molecule. In this case, an atom with lower electronegativity has a positive oxidation state, and an atom with higher electronegativity has a negative oxidation state.

It should be noted that the oxidation state cannot be identified with the valence of an element. Valence, defined as a number chemical bonds, by which a given atom is connected to other atoms, cannot be equal to zero and does not have a “+” or “ ” sign. The oxidation state can have both a positive and negative value, and can also take on a zero or even fractional value. Thus, in the CO 2 molecule, the oxidation state of C is +4, and in the CH 4 molecule, the oxidation state of C is 4. The valence of carbon in both compounds is IV.

Despite the above disadvantages, the use of the concept of oxidation state is convenient when classifying chemical compounds and drawing up equations of redox reactions.

In redox reactions, two occur interconnected process: oxidation and reduction.

Oxidation The process of electron loss is called. Recovery process of adding electrons.

Substances whose atoms or ions donate electrons are called restorers. Substances whose atoms or ions attach electrons (or withdraw a common pair of electrons) are called oxidizing agents.

When an element is oxidized, the oxidation state increases, in other words, the reducing agent during the reaction increases the oxidation state.

On the contrary, when an element is reduced, the oxidation state decreases, i.e., during the reaction, the oxidizing agent reduces the oxidation state.

Thus, we can give the following formulation of redox reactions: redox reactions are reactions that occur with a change in the oxidation state of the atoms of the elements that make up the reacting substances.

Oxidizing agents and reducing agents

To predict products and the direction of redox reactions, it is useful to remember that typical oxidizing agents are simple substances whose atoms have a large RER > 3.0 (elements of groups VIA and VIIA). Of these, the most powerful oxidizing agents are fluorine (OEO = 4.0), oxygen (OEO = 3.0), and chlorine (OEO = 3.5). Important oxidizing agents include PbO 2, KMnO 4, Ca(SO 4) 2, K 2 Cr 2 O 7 , HClO, HClO 3, KSIO 4, NaBiO 3, H 2 SO4 (conc), HNO 3 (conc), Na 2 O 2, (NH 4) 2 S 2 O 8, KSIO 3, H 2 O 2 and other substances , which contain atoms with higher or higher CO.

Typical reducing agents include simple substances whose atoms have a small REO< 1,5 (металлы IA и IIAгрупп и некоторые другие металлы). К важным восстановителям относятся H 2 S, NH 3 , HI, KI, SnCl 2 , FeSO 4 , C, H 2 , CO, H 2 SO 3 , Cr 2 (SO 4) 3 , CuCl, Na 2 S 2 O 3 и другие вещества, которые содержат атомы с низкими СО.

When composing equations for redox reactions, two methods can be used: the electron balance method and the ion-electronic method (half-reaction method). A more correct idea of ​​redox processes in solutions is provided by the ion-electronic method. Using this method, changes that actually exist in a solution are predicted by ions and molecules.

In addition to predicting reaction products, ionic half-reaction equations are necessary for understanding the redox processes that occur during electrolysis and in galvanic cells. This method reflects the role of the environment as a participant in the process. And finally, when using this method, it is not necessary to know in advance all the substances formed, since many of them are obtained by drawing up the equation of redox reactions.

It should be borne in mind that although half-reactions reflect the real processes occurring during redox reactions, they cannot be identified with the real stages (mechanism) of redox reactions.

The nature and direction of redox reactions are influenced by many factors: the nature of the reactants, the reaction of the medium, concentration, temperature, catalysts.

Biological significance of redox processes

Important processes in animal organisms are reactions of enzymatic oxidation of substrate substances: carbohydrates, fats, amino acids. As a result of these processes, organisms receive large number energy. Approximately 90% of the entire energy requirement of an adult male is met by the energy produced in tissues by the oxidation of carbohydrates and fats. The rest of the energy ~10% comes from oxidative cleavage amino acids.

Biological oxidation occurs through complex mechanisms with the participation large number enzymes. In mitochondria, oxidation occurs as a result of the transfer of electrons from organic substrates. As electron carriers, the mitochondrial respiratory chain includes various proteins containing various functional groups that are designed to transfer electrons. As they move along the chain from one intermediate to another, electrons lose free energy. For every pair of electrons transferred through the respiratory chain to oxygen, 3 ATP molecules are synthesized. The free energy released when 2 electrons are transferred to oxygen is 220 kJ/mol.

The synthesis of 1 ATP molecule under standard conditions requires 30.5 kJ. It is clear from this that a fairly significant part of the free energy released during the transfer of one pair of electrons is stored in ATP molecules. From these data, the role of multistage electron transfer from the initial reducing agent to oxygen becomes clear. The large energy (220 kJ) released during the transfer of one pair of electrons to oxygen is divided into a number of portions corresponding to individual stages of oxidation. At three such stages, the amount of energy released approximately corresponds to the energy required for the synthesis of 1 ATP molecule.

Based on changes in the oxidation states of the atoms that make up the reacting substances, chemical reactions are divided into two types.

1) Reactions that occur without changing the oxidation states of atoms.

For example:

2+4-2 t +2 -2 +4 -2
CaCO 3 = CaO + CO 2

In this reaction, the oxidation state of each atom remained unchanged.

2) Reactions that occur with a change in the oxidation states of atoms.

For example:

0 +2 -1 0 +2 -1
Zn + CuCl 2 = Cu + ZnCl 2

In this reaction, the oxidation states of the zinc and copper atoms changed.

Redox reactions are the most common chemical reactions.

In practice, a redox reaction is the gain or loss of electrons. Some atoms (ions, molecules) give or receive electrons from others.

Oxidation.

The process of giving up electrons by an atom, ion or molecule is called oxidation.

When electrons are lost, the oxidation state of an atom increases.

A substance whose atoms, ions or molecules give up electrons is called reducing agent.

In our example, atoms in the oxidation state 0 went into atoms with the oxidation state +2. That is, an oxidation process has occurred. In this case, the zinc atom, which donated two electrons, is a reducing agent (it increased the oxidation state from 0 to +2).

The oxidation process is recorded by an electronic equation, which indicates the change in the oxidation state of atoms and the number of electrons donated by the reducing agent.

For example:

0 +2 0
Zn – 2e – = Zn (oxidation, Zn – reducing agent).

Recovery.

The process of adding electrons is called restoration.

When electrons are added, the oxidation state of the atom decreases.

A substance whose atoms, ions or molecules gain electrons is called oxidizing agent.

In our example, the transition of copper atoms with oxidation state +2 to atoms with oxidation state 0 is a reduction process. In this case, a copper atom with an oxidation state of +2, accepting two electrons, lowers the oxidation state from +2 to 0 and is an oxidizing agent.

The oxidation process is also written using an electronic equation:

2 0 0
Cu + 2e – = Cu (reduction, Cu is an oxidizing agent).

The reduction process and the oxidation process are inseparable and occur simultaneously.

0 +2 0 +2
Zn + CuCl 2 = Cu + ZnCl 2
reducing agent oxidizing agent
oxidized reduced

There are two types of chemical reactions:

A Reactions in which the oxidation state of elements does not change:

Addition reactions

SO 2 + Na 2 O = Na 2 SO 3

Decomposition reactions

Cu(OH) 2 = CuO + H 2 O

Exchange reactions

AgNO 3 + KCl = AgCl + KNO 3

NaOH + HNO 3 = NaNO 3 + H 2 O

B Reactions in which there is a change in the oxidation states of the atoms of the elements that make up the reacting compounds and the transfer of electrons from one compound to another:

2Mg 0 + O 2 0 = 2Mg +2 O -2

2KI -1 + Cl 2 0 = 2KCl -1 + I 2 0

Mn +4 O 2 + 4HCl -1 = Mn +2 Cl 2 + Cl 2 0 + 2H 2 O

Such reactions are called redox reactions.

The oxidation state is the nominal charge of an atom in a molecule, calculated under the assumption that the molecule consists of ions and is generally electrically neutral.

The most electronegative elements in a compound have negative oxidation states, and the atoms of elements with less electronegativity have positive oxidation states.

Oxidation state is a formal concept; in some cases, the oxidation state does not coincide with the valency.

For example:

N 2 H 4 (hydrazine)

nitrogen oxidation degree – -2; nitrogen valence – 3.

Calculation of oxidation state

To calculate the oxidation state of an element, the following points should be taken into account:

1. Oxidation states of atoms in simple substances are equal to zero (Na 0; H 2 0).

2. The algebraic sum of the oxidation states of all atoms that make up a molecule is always equal to zero, and in a complex ion this sum is equal to the charge of the ion.

3. The following atoms have a constant oxidation state in compounds with atoms of other elements: alkali metals(+1), alkaline earth metals (+2), fluorine

(-1), hydrogen (+1) (except for metal hydrides Na + H -, Ca 2+ H 2 - and others, where the oxidation state of hydrogen is -1), oxygen (-2) (except F 2 -1 O + 2 and peroxides containing the –O–O– group, in which the oxidation state of oxygen is -1).

4. For elements, the positive oxidation state cannot exceed a value equal to the group number of the periodic system.

Examples:

V 2 +5 O 5 -2; Na 2 +1 B 4 +3 O 7 -2; K +1 Cl +7 O 4 -2 ; N -3 H 3 +1 ; K 2 +1 H +1 P +5 O 4 -2 ; Na 2 +1 Cr 2 +6 O 7 -2

Oxidation, reduction

In redox reactions, electrons are transferred from one atom, molecule, or ion to another. The process of losing electrons is oxidation. During oxidation, the oxidation state increases:

H 2 0 - 2ē = 2H + + 1/2О 2

S -2 - 2ē = S 0

Al 0 - 3ē = Al +3

Fe +2 - ē = Fe +3

2Br - - 2ē = Br 2 0

The process of adding electrons is reduction: During reduction, the oxidation state decreases.

Mn +4 + 2ē = Mn +2

S 0 + 2ē = S -2

Cr +6 +3ē = Cr +3

Cl 2 0 +2ē = 2Cl -

O 2 0 + 4ē = 2O -2

Atoms, molecules or ions that gain electrons in a given reaction are oxidizing agents, and those that donate electrons are reducing agents.

The oxidizing agent is reduced during the reaction, the reducing agent is oxidized.

Redox properties of a substance and the oxidation state of its constituent atoms

Compounds containing atoms of elements with the maximum oxidation state can only be oxidizing agents due to these atoms, because they have already given up all their valence electrons and are only able to accept electrons. The maximum oxidation state of an element's atom is equal to the number of the group in the periodic table to which the element belongs. Compounds containing atoms of elements with a minimum oxidation state can only serve as reducing agents, since they are only capable of donating electrons, because the external energy level in such atoms it is completed by eight electrons. The minimum oxidation state of metal atoms is 0, for non-metals - (n–8) (where n is the number of the group in periodic table). Compounds containing atoms of elements with intermediate oxidation states can be both oxidizing and reducing agents, depending on the partner with which they interact and the reaction conditions.

The most important reducing and oxidizing agents

Restorers

Carbon(II) monoxide (CO).

Hydrogen sulfide (H 2 S);

sulfur oxide (IV) (SO 2);

sulfurous acid H 2 SO 3 and its salts.

Hydrohalic acids and their salts.

Metal cations in lower oxidation states: SnCl 2, FeCl 2, MnSO 4, Cr 2 (SO4) 3.

Nitrous acid HNO2;

ammonia NH 3;

hydrazine NH 2 NH 2 ;

nitric oxide (II) (NO).

Cathode during electrolysis.

Oxidizing agents

Halogens.

Potassium permanganate (KMnO 4);

potassium manganate (K 2 MnO 4);

manganese (IV) oxide (MnO 2).

Potassium dichromate (K 2 Cr 2 O 7);

potassium chromate (K 2 CrO 4).

Nitric acid(HNO 3).

Sulfuric acid(H 2 SO 4) conc.

Copper(II) oxide (CuO);

lead(IV) oxide (PbO 2);

silver oxide (Ag 2 O);

hydrogen peroxide (H 2 O 2).

Iron(III) chloride (FeCl 3).

Berthollet's salt (KClO 3).

Anode during electrolysis.

On this basis, a distinction is made between redox reactions and reactions that occur without changing the oxidation states of chemical elements.

These include many reactions, including all substitution reactions, as well as those reactions of combination and decomposition in which at least one simple substance is involved, for example:

As you remember, coefficients in complex redox reactions are calculated using the electron balance method:

In organic chemistry, a striking example of redox reactions is the properties of aldehydes.

1. They are reduced to the corresponding alcohols:

2. Aldehydes are oxidized into the corresponding acids:

The essence of all the above examples of redox reactions was presented using the well-known electron balance method. It is based on comparing the oxidation states of atoms in the reactants and products of a reaction and on balancing the number of electrons in the processes of oxidation and reduction. This method is used to compile equations for reactions occurring in any phases. This makes it versatile and convenient. But at the same time, it has a serious drawback - when expressing the essence of redox reactions occurring in solutions, particles are indicated that do not really exist.

In this case, it is more convenient to use another method - the half-reaction method. It is based on the compilation of ion-electronic equations for the processes of oxidation and reduction, taking into account actually existing particles and their subsequent summation in general equation. In this method, the concept of “oxidation state” is not used, and the products are determined by deriving the reaction equation.

Let's demonstrate this method with an example: we will create an equation for the redox reaction of zinc with concentrated nitric acid.

1. We write down the ionic scheme of the process, which includes only a reducing agent and its oxidation product, an oxidizing agent and its reduction product:

2. We compose the ion-electronic equation of the oxidation process (this is the 1st half-reaction):

3. We compose the ion-electronic equation of the reduction process (this is the 2nd half-reaction):

Please note: electron-ion equations are written in accordance with the law of conservation of mass and charge.

4. We write the half-reaction equations so that the number of electrons between the reducing agent and the oxidizing agent is balanced:

5. Let us sum up the half-reaction equations term by term. Making up the general ionic equation reactions:

We check the correctness of the reaction equation in ionic form:

• Maintaining equality in the number of atoms of elements and in the number of charges
  1. The number of atoms of the elements must be equal on the left and right parts ionic reaction equation.
  2. The total charge of particles on the left and right sides of the ionic equation must be the same.

6. Write the equation in molecular form. To do this, we add to the ions included in the ionic equation the required number of ions of opposite charge.