Redox reactions in nature. Redox reactions Presentation on the topic of redox reactions

Oxidation is the process of donating electrons from an atom, molecule, or ion. An atom turns into a positively charged ion: Zn 0 - 2e Zn 2+ a negatively charged ion becomes a neutral atom: 2Cl - -2e Cl 2 0 S 2- -2e S 0 The value of a positively charged ion (atom) increases according to the number of donated electrons: Fe 2 + -1e Fe 3+ Mn +2 -2e Mn +4

Recovery is the process of adding electrons to an atom, molecule or ion. The atom turns into a negatively charged ion S 0 + 2e S 2 Br 0 + e Br The value of a positively charged ion (atom) decreases according to the number of attached electrons: Mn e Mn +2 S e S +4 or it can go into a neutral atom: H + + e H 0 Cu e Cu 0

Reducing agents are atoms, molecules, or ions that donate electrons. They are oxidized in the process of OVR Typical reducing agents: metal atoms with large atomic radii (I-A, II-A groups), as well as Fe, Al, Zn simple non-metal substances: hydrogen, carbon, boron; negatively charged ions: Cl, Br, I, S 2, N 3. Fluoride ions F are not a reducing agent. metal ions in the lower s.o.: Fe 2+, Cu +, Mn 2+, Cr 3+; complex ions and molecules containing atoms with an intermediate s.o.: SO 3 2, NO 2; CO, MnO 2, etc.

Oxidizing agents are atoms, molecules, or ions that accept electrons. They are reduced in the process of OVR. Typical oxidizing agents: atoms of non-metals VII-A, VI-A, V-A groups in the composition of simple substances, metal ions in the highest s.d.: Cu 2+, Fe 3+, Ag + ... complex ions and molecules containing atoms with higher and high s.d.: SO 4 2, NO 3, MnO 4, СlО 3, Cr 2 O 7 2-, SO 3, MnO 2, etc.

Sulfur oxidation states: -2.0, +4, +6 H 2 S -2 - reducing agent 2H 2 S + 3O 2 \u003d 2H 2 O + 2SO 2 S 0,S +4 O 2 - oxidizing agent and reducing agent S + O 2 \u003d SO 2 2SO 2 + O 2 \u003d 2SO 3 (reductant) S + 2Na \u003d Na 2 S SO 2 + 2H 2 S \u003d 3S + 2H 2 O (oxidizing agent) H 2 S +6 O 4 - oxidizing agent Cu + 2H 2 SO 4 \u003d CuSO 4 + SO 2 + 2H 2 O

Determination of the oxidation states of atoms of chemical elements С.о. atoms h / e in the composition of a simple being = 0 Algebraic sum of s.d. of all elements in the composition of the ion is equal to the charge of the ion Algebraic sum s.d. of all elements in the composition of a complex substance is 0. K +1 Mn +7 O x + 4 (-2) \u003d 0

Classification of redox reactions Reactions of intermolecular oxidation 2Al 0 + 3Cl 2 0 2Al +3 Cl 3 -1 Reactions of intramolecular oxidation 2KCl +5 O KCl O 2 0 Reactions of disproportionation, dismutation (self-oxidation-self-recovery): 3Cl KOH (gor.) KCl + 5 O 3 + 5KCl -1 + 3H 2 O 2N +4 O 2 + H 2 O HN +3 O 2 + HN +5 O 3

This is useful to know. The oxidation states of the elements in the salt anion are the same as in the acid, for example: (NH 4) 2 Cr 2 +6 O 7 and H 2 Cr 2 +6 O 7 The oxidation state of oxygen in peroxides is -1 The oxidation state sulfur in some sulfides is -1, for example: FeS 2 Fluorine is the only non-metal that does not have a positive oxidation state in compounds. In compounds NH 3, CH 4, etc., the sign of the electropositive element hydrogen is in second place

Oxidizing properties of concentrated sulfuric acid Sulfur reduction products: H 2 SO 4 + och.akt. metal (Mg, Li, Na…) H 2 S H 2 SO 4 + act. metal (Mn, Fe, Zn…) S H 2 SO 4 + inactive metal (Cu, Ag, Sb…) SO 2 H 2 SO 4 + HBr SO 2 H 2 SO 4 + non-metals (C, P, S…) SO 2 Note: it is often possible to form a mixture of these products in different proportions

Hydrogen peroxide in redox reactions Solution medium Oxidation (H 2 O 2 -reducing agent) Reduction (H 2 O 2 -oxidizing agent) acidic H 2 O 2 -2eO 2 + 2H + (O - 2eO 2 0) H 2 O 2 + 2H + + 2e2H 2 O (O e2O - 2) alkaline H 2 O 2 + 2OH -O 2 + 2H 2 O (O - 2eO 2 0) H 2 O 2 + 2e2OH - (O e2O - 2) neutral H 2 O 2 - 2eO 2 + 2H + (O - 2eO 2 0) H 2 O 2 + 2e2OH - (O e2O - 2)

Nitric acid in redox reactions Nitrogen reduction products: Concentrated HNO 3: N +5 +1e N +4 (NO 2) (Ni, Cu, Ag, Hg; C, S, P, As, Se); passivates Fe, Al, Cr Diluted HNO 3: N +5 +3e N +2 (NO) (Metals in ECHRNM Al …Cu; non-metals S, P, As, Se) Diluted HNO 3: N +5 +4e N +1 (N 2 O) Ca, Mg, Zn Dilute HNO 3: N +5 +5e N 0 (N 2) Very dilute: N e N -3 (NH 4 NO 3) (active metals in ECHRNM up to Al)

Importance of OVR OVR is extremely common. They are associated with metabolic processes in living organisms, respiration, decay, fermentation, photosynthesis. OVR provide the cycle of substances in nature. They can be observed during fuel combustion, corrosion and metal smelting. With their help, alkalis, acids and other valuable chemicals are obtained. OVR underlie the conversion of the energy of interacting chemicals into eclectic energy in galvanic batteries.

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The concept of redox reactions Chemical reactions that occur with a change in the degree of oxidation of the elements that make up the reactants are called redox reactions.

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Oxidation is the process of donating electrons from an atom, molecule, or ion. An atom turns into a positively charged ion: Zn0 - 2e → Zn2+ a negatively charged ion becomes a neutral atom: 2Cl- -2e →Cl20 S2- -2e →S0 The value of a positively charged ion (atom) increases according to the number of donated electrons: Fe2+ -1e →Fe3+ Mn +2 -2e →Mn+4

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Recovery is the process of adding electrons to an atom, molecule or ion. An atom turns into a negatively charged ion S0 + 2e → S2− Br0 + e → Br − The value of a positively charged ion (atom) decreases according to the number of attached electrons: or it can go into a neutral atom: H+ + e → H0 Cu2+ + 2e → Cu0

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Reducing agents are atoms, molecules, or ions that donate electrons. They are oxidized during the redox process Typical reducing agents: ● metal atoms with large atomic radii (I-A, II-A groups), as well as Fe, Al, Zn ● simple non-metal substances: hydrogen, carbon, boron; ● negatively charged ions: Cl−, Br−, I−, S2−, N−3. Fluoride ions F– are not reducing agents. ● metal ions in the lowest s.d.: Fe2+, Cu+, Mn2+, Cr3+; ● complex ions and molecules containing atoms with intermediate d.s.: SO32−, NO2−; CO, MnO2, etc.

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Oxidizing agents are atoms, molecules, or ions that accept electrons. They are reduced during the redox process. Typical oxidizing agents: ● atoms of non-metals of groups VII-A, VI-A, V-A in the composition of simple substances ● metal ions in the highest s.d.: Cu2+, Fe3+, Ag+ ... ● complex ions and molecules containing atoms with highest and high s.d.: SO42-, NO3-, MnO4-, СlО3-, Cr2O72-, SO3, MnO2, etc.

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The manifestation of redox properties is influenced by such a factor as the stability of a molecule or ion. The stronger the particle, the less it exhibits redox properties.

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For example, nitrogen has a high electronegativity and could be a strong oxidizing agent in the form of a simple substance, but its molecule has a triple bond, the molecule is very stable, nitrogen is chemically passive.

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Or HCLO is a stronger oxidizing agent in solution than HCLO4, since HCLO is a less stable acid.

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If a chemical element is in an intermediate oxidation state, then it exhibits the properties of both an oxidizing agent and a reducing agent.

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Sulfur oxidation states: -2.0, +4, +6 2Na=Na2S SO2+2H2S=3S+2H2O (oxidizing agent) H2S+6O4 - oxidizing agent Cu+2H2SO4=CuSO4+SO2+2H2O

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Classification of redox reactions Intermolecular oxidation reactions 2Al0 + 3Cl20 → 2Al+3 Cl3-1 Intramolecular oxidation reactions 2KCl + 5O3-2 → 2KCl-1 + 3O20 Reactions of disproportionation, dismutation (self-oxidation-self-recovery): 3Cl20 + 6KOH (horizontal) → KCl+5O3 +5KCl-1+3H2O 2N+4O2+ H2O →HN+3O2 + HN+5O3

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This is useful to know. The oxidation states of the elements in the composition of the salt anion are the same as in the acid, for example: (NH4)2Cr2 + 6O7 and H2Cr2 + 6O7 The oxidation state of oxygen in peroxides is -1 The oxidation state of sulfur in some sulfides is -1, for example: FeS2 Fluorine is the only non-metal that does not have a positive oxidation state in compounds. In the compounds NH3, CH4, etc., the sign of the electropositive element hydrogen is in second place

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Oxidizing properties of concentrated sulfuric acid Sulfur reduction products: H2SO4 + pt. metal (Mg, Li, Na…) → H2S H2SO4 + act. metal (Mn, Fe, Zn…) → S H2SO4 + inactive. metal (Cu, Ag, Sb…) → SO2 H2SO4 + HBr → SO2 H2SO4 + non-metals (C, P, S…) → SO2 Note: it is often possible to form a mixture of these products in different proportions

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Completed by: Chemistry teacher Baimukhametova Batila Turginbaevna Redox reactions

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The motto of the lesson is “Someone loses, and someone finds ...” By yourself, by working, you will do everything for your loved ones and for yourself, and if there is no success during work, failure is not a problem, try again. D. I. Mendeleev.

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Lesson topic: “Redox reactions” Purpose: To get acquainted with redox reactions and find out what is the difference between exchange reactions and redox reactions. Learn to identify oxidizing and reducing agents in reactions. Learn to draw diagrams of the processes of giving and receiving electrons. To get acquainted with the most important redox reactions occurring in nature.

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Perhaps these electrons are Worlds where there are five continents, Arts, knowledge, wars, thrones And the memory of forty centuries! Also, perhaps, each atom is the Universe, where there are a hundred planets; There - everything that is here, in a compressed volume, But also what is not here. V. Bryusosov.

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What is an oxidation state? The oxidation state is the conditional charge of an atom of a chemical element in a compound, calculated on the basis of the assumption that all compounds consist only of ions. The oxidation state can be positive, negative or equal to zero, depending on the nature of the respective compounds. Some elements have: constant oxidation states, others - variables. Elements with a constant positive oxidation state include - alkali metals: Li + 1, Na + 1, K + 1, Rb + 1, Cs + 1, Fr + 1, the following elements of group II of the periodic system: Be + 2, Mg + 2, Ca + 2, Sr + 2, Ba + 2, Ra + 2, Zn + 2, as well as an element of III A group - A1 + 3 and some others. Metals in compounds always have a positive oxidation state. Of non-metals, F has a constant negative oxidation state (-1). In simple substances formed by atoms of metals or non-metals, the oxidation states of elements are zero, for example: Na °, Al °, Fe °, H2, O2, F2, Cl2, Br2. Hydrogen is characterized by oxidation states: +1 (H20), -1 (NaH). Oxygen is characterized by oxidation states: -2 (H20), -1 (H2O2), +2 (OF2).

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The most important reducing agents and oxidizing agents Reductants: Oxidizing agents: Metals-simple substances Hydrogen Carbon Carbon monoxide (II) (CO) Hydrogen sulfide (H2S) Sulfur oxide (IV) (SO2) Sulfurous acid H2SO3 and its salts Hydrohalic acids and their salts Metal cations in intermediate degrees oxidation: SnCl2, FeCl2, MnSO4, Cr2(SO4)3 Nitrous acid HNO2 Ammonia NH3 Nitrogen(II) oxide (NO) Halogens Potassium permanganate (KMnO4) Potassium manganate (K2MnO4) Manganese (IV) oxide (MnO2) Potassium dichromate (K2Cr2O7) Nitrogenous acid (HNO3) Sulfuric acid (conc. H2SO4) Copper(II) oxide (CuO) Lead(IV) oxide (PbO2) Hydrogen peroxide (H2O2) Iron(III) chloride (FeCl3) Organic nitro compounds

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The oxidation state of manganese in the potassium permanganate compound KMnO4. 1. The oxidation state of potassium +1, oxygen -2. 2. Calculate the number of negative charges: 4 (-2) \u003d - 8 3. The number of positive charges in manganese is 1. 4. We make the following equation: (+1) + x + (-2) * 4 \u003d 0 1+ x - 8 \u003d 0 X \u003d 8 - 1 \u003d 7 X \u003d +7 +7 is the oxidation state of manganese in potassium permanganate.

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Rules for determining oxidation states 1. The oxidation state of an element in a simple substance is 0. For example: Ca, H2, Cl2, Na. 2. The oxidation state of fluorine in all compounds except F2 is - 1. Example: S + 6F6-1 3. The oxidation state of oxygen in all compounds except O2, O3, F2-1O + 2 and peroxide compounds Na2 + 1 O- 12; H2 + 1O-12 is equal to -2 Examples: Na2O-2, BaO-2, CO2-2. 4. The oxidation state of hydrogen is +1 if there is at least one non-metal in the compounds, -1 in compounds with metals (hydrides) 5. The oxidation state of O in H2 Examples: C-4H4 + 1 Ba + 2H2-1 H2 The oxidation state of metals is always positive (except for simple substances). The oxidation state of metals of the main subgroups is always equal to the group number. The oxidation state of side subgroups can take different meanings. Examples: Na+ Cl-, Al2+3O3-2, Cr2+3 O3-2, Cr+2O-2. 6. The maximum positive oxidation state is equal to the group number (exceptions Cu+2, Au+3). The minimum oxidation state is the group number minus eight. Examples: H+1N+5O-23, N-3H+13. 7. The sum of the oxidation states of atoms in a molecule (ion) is equal to 0 (ion charge).

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Laboratory work Safety regulations. Experience 1. Carry out a chemical reaction between solutions of copper (II) sulfate and sodium hydroxide. Experience 2. 1. Place an iron nail in a solution of copper (II) sulfate. 2. Make the equations of chemical reactions. 3. Determine the type of each chemical reaction. 4. Determine the oxidation state of the atom of each chemical element before and after the reaction. 5. Think about how these reactions differ?

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Answers: Cu + 2S + 6O4-2 + 2Na + 1O-2H + 1Cu + 2 (O -2H + 1) 2 + Na2 + 1S + 6O4-2 - exchange reaction Cu + 2S + 6O4-2 + Fe0 Fe + 2 S + 6O4 -2 + Сu0 - substitution reaction Reaction No. 2 differs from reaction No. 1 in that in this case the oxidation state of the atoms of chemical elements changes before and after the reaction. Note this important difference between the two reactions. The second reaction is OVR. We emphasize in the reaction equation the symbols of chemical elements that have changed their oxidation state. Let's write them down and indicate what the atoms did with their electrons (Given away or received?), i.e. electron transitions. Cu + 2 + 2 e-  Cu0 - oxidizing agent, reduced Fe0 - 2 e-  Fe + 2 - reducing agent, oxidized

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Classification of redox reactions 1. Intermolecular redox reactions The oxidizing and reducing agents are in different substances; the exchange of electrons in these reactions occurs between different atoms or molecules: 2Ca0 + O20 → 2 Ca + 2O-2 Ca is the reducing agent; O2 - oxidizing agent Cu+2O + C+2O → Cu0 + C+4O2 CO - reducing agent; CuO is an oxidizing agent Zn0 + 2HCl → Zn+2Cl2 + H20 Zn is a reducing agent; HСl - oxidizing agent Mn+4O2 + 2KI-1 + 2H2SO4 → I20 + K2SO4 + Mn+2SO4 + 2H2O KI - reducing agent; MnO2 is an oxidizing agent.

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2. Intramolecular redox reactions In intramolecular reactions, the oxidizing agent and reducing agent are in the same molecule. Intramolecular reactions proceed, as a rule, during the thermal decomposition of substances containing an oxidizing agent and a reducing agent. 4Na2Cr2O7 → 4Na2CrO4 + 2Cr2O3 + 3O2 Cr+6- oxidizer; O-2 - reducing agent

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3. Disproportionation reactions Redox reactions in which one element simultaneously raises and lowers the degree of oxidation. 3S + 6NaOH → Na2SO3 + 2Na2S + 3H2O Sulfur in oxidation state 0 is both an oxidizing agent and a reducing agent. 4. Comportionation reactions Redox reactions in which atoms of one element in different oxidation states acquire one oxidation state as a result of the reaction. 5NaBr + NaBrO3 + 3H2SO4 → 3Na2SO4 + 3Br2 + 3H2O Br+5 is an oxidizing agent; Br-1 - reducing agent

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Algorithm for compiling equations of redox reactions using the electron balance method 1. Write down the reaction scheme KMnO4 + KI + H2SO4 → MnSO4 + I2 + K2SO4 + H2O 2. Put down the oxidation state of the atoms of the elements in which it changes KMn + 7O4 + KI- + H2SO4 → Mn + 2SO4 + I20+ K2SO4+ H2O 3. Elements that change oxidation states are isolated and the number of electrons accepted by the oxidizing agent and given away by the reducing agent is determined. Mn + 7 + 5 ē → Mn + 2 2I-1 - 2 ē → I20 4. Equalize the number of received and given electrons, thereby establishing the coefficients for compounds in which there are elements that change the oxidation state. Mn + 7 + 5ē → Mn + 22 2I-1 - 2ē → I205 2Mn + 7 + 10I-1 → 2Mn + 2 + 5I20 5. Select the coefficients for all other participants in the reaction. 2KMnO4+10KI+8H2SO4→2MnSO4+5I2+6K2SO4+ 8H2O

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Electronic balance is a method for finding coefficients in the equations of redox reactions, in which the exchange of electrons between atoms of elements that change their oxidation state is considered. The number of electrons donated by the reducing agent is equal to the number of electrons received by the oxidizing agent.

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Redox reactions are reactions in which oxidation and reduction processes occur simultaneously and, as a rule, the oxidation states of elements change. Consider the process using the example of the interaction of zinc with dilute sulfuric acid:

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Let's remember: 1. Oxidation-reduction reactions are such reactions in which the transfer of electrons from one atom, molecule or ion to another occurs. 2. Oxidation is the process of electron donation, the degree of oxidation increases. 3. Restoration is the process of adding electrons, while the oxidation state decreases. 4.Atoms, molecules or ions that donate electrons are oxidized; are restorers. 5.Atoms, ions or molecules that accept electrons are reduced; are oxidizing agents. 6. Oxidation is always accompanied by reduction, reduction is associated with oxidation. 7. Oxidation - reduction reactions - the unity of two opposite processes: oxidation and reduction.

REDOX REACTIONS

1. OVR. OVR classification.
2.Method of electronic balance.
3. Method of half-reactions.

Targets and goals:

To consolidate the ability of students to apply the concept of "oxidation state" in practice.
Summarize and supplement students' knowledge about the basic concepts of the OVR theory.
To improve the ability of students to apply these concepts to the explanation of facts.

Targets and goals:

Introduce students to the essence of the half-reaction method.
To form the ability to express the essence of redox reactions occurring in solutions using the ion-electronic method.

Oxidizing agent and reducing agent

oxidizing agent A reagent that accepts electrons in a redox reaction is called.
restorer is a reagent that donates electrons in a redox reaction.

OXIDATION AND REDUCTION PROCESS

Oxidation called the process of donating electrons by an atom, molecule or ion, which is accompanied by an increase in the degree of oxidation.
Recovery call the process of adding electrons to an atom, molecule or ion, which is accompanied by a decrease in the degree of oxidation.

Rules for determining the function of a compound in redox reactions.

1. If the element exhibits in the connection highest degree of oxidation then this connection can be oxidizing agent.
2. If the item exhibits in connection lower oxidation state then this connection can be reducing agent.
3. If the item exhibits in connection intermediate oxidation state then this connection can be like reducer, so oxidizing agent.
Exercise:
Predict the functions of substances in redox reactions:

The most important oxidizing and reducing agents Basic concepts of the OVR theory

Questions:
1. What is called the recovery process?
2. How does the oxidation state of an element change during reduction?
3. What is called the oxidation process?
4. How does the oxidation state of an element change during oxidation?
5. Define the concept of "reductant".
6. Define the concept of "oxidant".
7. How to predict the function of a substance by the oxidation state of an element?
8. Name the most important reducing agents and oxidizing agents.
9. What reactions are called redox reactions?

chemical reactions chemical reactions

By changing the oxidation state of atoms of elements

Redox

Without changing the oxidation state of the atoms of the elements
These include all ion exchange reactions, as well as many compound reactions.

REDOX REACTIONS

redox
called reactions that are accompanied by a change in the oxidation states of the chemical elements that make up the reagents.

OVR classification

intermolecular oxidation-reduction reactions

intramolecular oxidation-reduction reactions,

reactions of disproportionation, dismutation or self-oxidation-self-recovery

Intermolecular reactions:

Electron donor particles (reductants) - and electron acceptor particles (oxidizers) - are in different substances.
This type includes the majority of OVR.

Intramolecular reactions

An electron donor - a reducing agent - and an electron acceptor - an oxidizing agent - are in the same substance.

Reactions of dismutation, or disproportionation, or self-oxidation-self-healing

Atoms of the same element in a substance simultaneously perform the functions of both electron donors (reducing agents) and electron acceptors (oxidizing agents).
These reactions are possible for substances containing atoms of chemical elements in an intermediate oxidation state.

Compilation of redox reactions

For the preparation of redox reactions, use:
1) electronic balance method
2) Drawing up equations of redox reactions by the half-reaction method, or by the ion-electron method

Compilation of redox reactions electronic balance method

The method is based on comparing the oxidation states of atoms in the initial substances and reaction products and on balancing the number of electrons shifted from the reducing agent to the oxidizing agent.
The method is applied for compiling equations of reactions occurring in any phases. This is the versatility and convenience of the method.
The disadvantage of the method- when expressing the essence of reactions occurring in solutions, the existence of real particles is not reflected.

Algorithmic prescription for compiling equations of redox reactions by the electron balance method

1. Draw up a reaction scheme.
2. Determine the oxidation states of the elements in the reactants and reaction products.
3. Determine whether the reaction is redox or it proceeds without changing the oxidation states of the elements. In the first case, perform all subsequent operations.
4. Underline the elements whose oxidation states change.

5. Determine which element is oxidized (its oxidation state increases) and which element is reduced (its oxidation state decreases) during the reaction.
6. On the left side of the diagram, use arrows to denote the oxidation process (displacement of electrons from an atom of an element) and the reduction process (displacement of electrons to an atom of an element)
7. Determine the reducing agent (the atom of the element from which the electrons are displaced) and the oxidizing agent (the atom of the element to which the electrons are displaced).

Algorithmic prescription for compiling equations of redox reactions by the electron balance method

8. Balance the number of electrons between the oxidizing agent and reducing agent.
9. Determine the coefficients for the oxidizing agent and reducing agent, oxidation and reduction products.
10. Write down the coefficient in front of the formula of the substance that determines the environment of the solution.
11. Check the reaction equation.

Drawing up equations of redox reactions by the half-reaction method, or by the ion-electron method

The method is based on the compilation of ion-electronic equations for the processes of oxidation and reduction, taking into account real-life particles and their subsequent summation into a general equation.
Method applied to express the essence of redox reactions occurring only in solutions.
Advantages of the method.
1. In the electron-ion equations of half-reactions, ions that actually exist in an aqueous solution are written, and not conditional particles. (For example, ions rather than a nitrogen atom with an oxidation state of +3 and a sulfur atom with an oxidation state of +4.)
2. The concept of "oxidation state" is not used.
3. When using this method, you do not need to know all the substances: they are determined when deriving the reaction equation.
4. The role of the environment as an active participant in the entire process is visible.

The main stages of compiling equations of redox reactions ion-electronic method

(on the example of the interaction of zinc with concentrated nitric acid)
1. We write down the ionic scheme of the process, which includes only the reducing agent and its oxidation product, and the oxidizing agent and its reduction product:

Source

USE. CHEMISTRY: Universal reference book / O.V. Meshkova.- M.: EKSMO, 2010.- 368s.

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