What to answer to a person who is interested in how to solve redox reactions? They are unsolvable. However, like any others. Chemists generally do not solve reactions or their equations. For an oxidation-reduction reaction (ORR), you can create an equation and place the coefficients in it. Let's look at how to do this.
Oxidizing agent and reducing agent
A redox reaction is a reaction in which the oxidation states of the reactants change. This happens because one of the particles gives up its electrons (it is called a reducing agent), and the other accepts them (an oxidizing agent).
The reducing agent, losing electrons, oxidizes, that is, it increases the value of the oxidation state. For example, the entry: 
means that zinc gave up 2 electrons, that is, it was oxidized. He is a restorer. The degree of oxidation, as can be seen from the above example, has increased. 
– here sulfur accepts electrons, that is, it is reduced. She is an oxidizing agent. Its oxidation level decreased.
Someone may wonder why, when electrons are added, the oxidation state decreases, but when they are lost, on the contrary, it increases? Everything is logical. An electron is a particle with a charge of -1, therefore, from a mathematical point of view, the entry should be read as follows: 0 – (-1) = +1, where (-1) is the electron. Then it means: 0 + (-2) = -2, where (-2) are the two electrons that the sulfur atom accepted.
Now consider a reaction in which both processes occur:
Sodium reacts with sulfur to form sodium sulfide. Sodium atoms are oxidized, giving up one electron at a time, while sulfur atoms are reduced, gaining two. However, this can only be on paper. In fact, the oxidizing agent must add to itself exactly as many electrons as the reducing agent gave them. In nature, balance is maintained in everything, including redox processes. Let us show the electronic balance for this reaction:
The total multiple between the number of electrons given and received is 2. Dividing it by the number of electrons given up by sodium (2:1=1) and sulfur (2:2=1) we get the coefficients in this equation. That is, on the right and left sides of the equation there should be one sulfur atom each (the value obtained by dividing the common multiple by the number of electrons accepted by sulfur), and two sodium atoms. In the written diagram on the left there is still only one sodium atom. Let's double it by putting a factor of 2 in front of the sodium formula. The right side of the sodium atoms already contains 2 (Na2S).
We have compiled an equation for the simplest redox reaction and placed the coefficients in it using the electronic balance method.
Let's look at how to “solve” more complex redox reactions. For example, when concentrated sulfuric acid reacts with the same sodium, hydrogen sulfide, sodium sulfate and water are formed. Let's write down the diagram:
Let us determine the oxidation states of atoms of all elements:
Changed art. only sodium and sulfur. Let us write the half-reactions of oxidation and reduction:
Let's find the least common multiple between 1 (how many electrons sodium gave up) and 8 (the number of negative charges accepted by sulfur), divide it by 1, then by 8. The results are the number of Na and S atoms on both the right and left.
Let's write them into the equation:
We do not yet put the coefficients from the balance sheet in front of the sulfuric acid formula. We count other metals, if any, then acid residues, then H, and last but not least we check for oxygen.
In this equation, there should be 8 sodium atoms on the right and left. The sulfuric acid residues are used twice. Of these, 4 become salt formers (part of Na2SO4) and one turns into H2S, that is, a total of 5 sulfur atoms must be consumed. We put 5 in front of the sulfuric acid formula.
We check H: there are 5×2=10 H atoms on the left side, only 4 on the right side, which means we put a coefficient of 4 in front of water (it cannot be put in front of hydrogen sulfide, since it follows from the balance that there should be 1 H2S molecules on the right and left. We check for oxygen. On the left there are 20 O atoms, on the right there are 4x4 from sulfuric acid and another 4 from water. Everything matches, which means the actions were performed correctly.
This is one type of activity that someone who asked how to solve redox reactions might have in mind. If this question meant “finish the ORR equation” or “add the reaction products,” then to complete such a task it is not enough to be able to draw up an electronic balance. In some cases, you need to know what the oxidation/reduction products are, how they are affected by the acidity of the environment and various factors that will be discussed in other articles.
Redox reactions - video
The whole variety of chemical reactions can be reduced to two types. If, as a result of a reaction, the oxidation states of elements do not change, then such reactions are called exchange, otherwise - redox reactions.
The occurrence of chemical reactions is due to the exchange of particles between reacting substances. For example, in a neutralization reaction, an exchange occurs between cations and anions of an acid and a base, resulting in the formation of a weak electrolyte - water:
Often the exchange is accompanied by the transfer of electrons from one particle to another. Thus, when zinc replaces copper in a solution of copper (II) sulfate
electrons from zinc atoms go to copper ions:
The process of a particle losing electrons is called oxidation, and the process of acquiring electrons is restoration. Oxidation and reduction occur simultaneously, therefore interactions accompanied by the transfer of electrons from one particle to another are called redox reactions.
The transfer of electrons may be incomplete. For example, in the reaction
Instead of low-polar C-H bonds, highly polar H-Cl bonds appear. For the convenience of writing redox reactions, the concept of oxidation degree is used, which characterizes the state of an element in a chemical compound and its behavior in reactions.
Oxidation state- a value numerically equal to the formal charge that can be assigned to an element, based on the assumption that all the electrons of each of its bonds have transferred to a more electronegative atom of the given compound.
Using the concept of oxidation state, we can give a more general definition of the processes of oxidation and reduction. Redox are called chemical reactions that are accompanied by a change in the oxidation states of the elements of the substances involved in the reaction. During reduction, the oxidation state of an element decreases; during oxidation, it increases. A substance that contains an element that reduces its oxidation state is called oxidizing agent; a substance that contains an element that increases the oxidation state is called reducing agent.
The oxidation state of an element in a compound is determined in accordance with the following rules:
· the oxidation state of an element in a simple substance is zero;
· the algebraic sum of all oxidation states of atoms in a molecule is equal to zero;
· the algebraic sum of all oxidation states of atoms in a complex ion, as well as the oxidation state of an element in a simple monatomic ion, is equal to the charge of the ion;
· a negative oxidation state is exhibited in a compound by the atoms of the element having the highest electronegativity;
· the maximum possible (positive) oxidation state of an element corresponds to the number of the group in which the element is located in the D.I. Periodic Table. Mendeleev.
The oxidation state of atoms of elements in a compound is written above the symbol of a given element, indicating first the sign of the oxidation state, and then its numerical value, for example.
A number of elements in compounds exhibit a constant oxidation state, which is used in determining the oxidation states of other elements:
The redox properties of atoms of various elements manifest themselves depending on many factors, the most important of which are the electronic structure of the element, its oxidation state in the substance, and the nature of the properties of other participants in the reaction. Compounds that contain atoms of elements with a maximum (positive) oxidation state, for example, can only be reduced, acting as oxidizing agents. Compounds containing elements with minimal oxidation states, e.g. 
can only oxidize and act as reducing agents.
Substances containing elements with intermediate oxidation states, e.g. 
have redox duality. Depending on the reaction partner, such substances are capable of both accepting (when interacting with stronger reducing agents) and donating (when interacting with stronger oxidizing agents) electrons.
The composition of reduction and oxidation products also depends on many factors, including the environment in which the chemical reaction occurs, the concentration of reagents, and the activity of the partner in the redox process.
To write down the equation for a redox reaction, you need to know how the oxidation states of elements change and what other states the oxidizing agent and reducing agent go to. Let's look at brief characteristics of the most commonly used oxidizing and reducing agents.
The most important oxidizing agents. Among simple substances, oxidizing properties are typical for typical non-metals: fluorine F 2, chlorine Cl 2, bromine Br 2, iodine I 2, oxygen O 2.
Halogens, being reduced, they acquire an oxidation state of -1, and from fluorine to iodine their oxidizing properties weaken (F 2 has limited use due to its high aggressiveness):
Oxygen, being reduced, acquires an oxidation state of -2:
The most important oxidizing agents among oxygen-containing acids and their salts include nitric acid HNO 3 and its salts, concentrated sulfuric acid H 2 SO 4, oxygen-containing halogen acids HHalO x and their salts, potassium permanganate KMnO 4 and potassium dichromate K 2 Cr 2 O 7.
Nitric acid exhibits oxidizing properties due to nitrogen in the oxidation state +5. In this case, the formation of various reduction products is possible:
The depth of nitrogen reduction depends on the acid concentration, as well as on the activity of the reducing agent, determined by its redox potential:
Fig.1. The depth of nitrogen reduction depending on the acid concentration.
For example, the oxidation of zinc (an active metal) with nitric acid is accompanied by the formation of various reduction products; at a concentration of HNO 3 of approximately 2% (wt.), NH 4 NO 3 is predominantly formed:
at a HNO 3 concentration of approximately 5% (wt.) – N 2 O:
at a HNO 3 concentration of about 30% (wt.) – NO:
and at a concentration of HNO 3 of approximately 60% (wt.), NO 2 is predominantly formed:
The oxidative activity of nitric acid increases with increasing concentration, so concentrated HNO 3 oxidizes not only active, but also slightly active metals, such as copper and silver, forming predominantly nitric oxide (IV):
as well as non-metals, such as sulfur and phosphorus, oxidizing them to acids corresponding to higher oxidation states:
Nitric acid salts ( nitrates) can be reduced in acidic, and when interacting with active metals and in alkaline media, as well as in melts:
Aqua regia– a mixture of concentrated and nitric acids, mixed in a ratio of 1:3 by volume. The name of this mixture is due to the fact that it dissolves even such noble metals as gold and platinum:
The occurrence of this reaction is due to the fact that aqua regia releases nitrosyl chloride NOCl and free chlorine Cl2:
under the influence of which metals turn into chlorides.
Sulfuric acid exhibits oxidizing properties in a concentrated solution due to sulfur in the oxidation state +6:
The composition of the reduction products is determined mainly by the activity of the reducing agent and the acid concentration:
Fig.2. Reducing activity of sulfur depending on
acid concentration.
Thus, the interaction of concentrated H 2 SO 4 with low-active metals, some non-metals and their compounds leads to the formation of sulfur oxide (IV):
Active metals reduce concentrated sulfuric acid to sulfur or hydrogen sulfide:
in this case, H 2 S, S and SO 2 are simultaneously formed in different ratios. However, in this case, the main product of the reduction of H 2 SO 4 is SO 2, since the released S and H 2 S can be oxidized by concentrated sulfuric acid:
and their salts (see Table A.1.1) are often used as oxidizing agents, although many of them exhibit a dual character. As a rule, the reduction products of these compounds are chlorides and bromides (oxidation state -1), as well as iodine (oxidation state 0);
However, even in this case, the composition of the reduction products depends on the reaction conditions, the concentration of the oxidizing agent and the activity of the reducing agent:
Potassium permanganate exhibits oxidizing properties due to manganese in the oxidation state +7. Depending on the environment in which the reaction takes place, it is reduced to different products: in an acidic environment - to manganese (II) salts, in a neutral environment - to manganese (IV) oxide in the hydrated form MnO(O) 2, in an alkaline environment - to manganate -and she
acidic environment
neutral environment
alkaline environment
Potassium dichromate, the molecule of which includes chromium in the oxidation state +6, is a strong oxidizing agent during sintering and in an acid solution
exhibits oxidizing properties in a neutral environment
In an alkaline environment, the equilibrium between chromate and dichromate ions
is shifted towards formation, therefore in an alkaline environment the oxidizing agent is potassium chromate K 2 СrO 4:
however, K 2 CrO 4 is a weaker oxidizing agent compared to K 2 Cr 2 O 7 .
Among the ions, the hydrogen ion H + and metal ions in the highest oxidation state exhibit oxidizing properties. Hydrogen ion H + acts as an oxidizing agent when active metals interact with dilute acid solutions (with the exception of HNO 3)
Metal ions in a relatively high oxidation state, such as Fe 3+, Cu 2+, Hg 2+, being reduced, turning into ions of a lower oxidation state
or are isolated from solutions of their salts in the form of metals
The most important reducing agents. Typical reducing agents among simple substances include active metals, such as alkali and alkaline earth metals, zinc, aluminum, iron and others, as well as some non-metals (hydrogen, carbon, phosphorus, silicon).
Metals in an acidic environment they are oxidized to positively charged ions:
In an alkaline environment, metals exhibiting amphoteric properties are oxidized; in this case, negatively charged anions or hydroxocomponents are formed:
Nonmetals, oxidizing, form oxides or corresponding acids:
Reducing functions are possessed by oxygen-free anions, for example Cl -, Br -, I -, S 2-, H - and metal cations in the highest oxidation state.
In a row halide ions, which, when oxidized, usually form halogens:
reducing properties increase from Cl - to I - .
Hydrides metals exhibit reducing properties due to the oxidation of hydrogen bound (oxidation state -1) to free hydrogen:
Metal cations in the lowest oxidation state, such as Sn 2+, Fe 2+, Cu +, Hg 2 2+ and others, when interacting with oxidizing agents, the oxidation degree increases:
Redox duality. Among simple substances, redox duality is characteristic of elements VIIA, VIA and VA subgroups, which can either increase or decrease their oxidation state.
Often used as oxidizing agents halogens under the influence of stronger oxidizing agents they exhibit reducing properties (with the exception of fluorine). Their oxidizing abilities decrease, and their reducing properties increase from Cl 2 to I 2:
Fig.3. Redox ability of halogens.
This feature is illustrated by the reaction of iodine oxidation with chlorine in an aqueous solution:
The composition of oxygen-containing compounds that exhibit dual behavior in redox reactions also includes elements in an intermediate oxidation state. Oxygen-containing acids of halogens and their salts, the molecules of which include a halogen in an intermediate oxidation state, can act as oxidizing agents
and reducing agents
Hydrogen peroxide, containing oxygen in the oxidation state -1, in the presence of typical reducing agents exhibits oxidizing properties, since the oxidation state of oxygen can decrease to -2:
The latter reaction is used in the restoration of paintings by old masters, the paints of which, containing lead white, turn black due to interaction with hydrogen sulfide in the air.
When interacting with strong oxidizing agents, the oxidation state of oxygen included in hydrogen peroxide increases to 0, H 2 O 2 exhibits the properties of a reducing agent:
Nitrous acid And nitrites, which contain nitrogen in the oxidation state +3, and can also act as oxidizing agents
as well as in the role of restorers
Classification. There are four types of redox reactions.
1. If the oxidizing agent and the reducing agent are different substances, then such reactions belong to intermolecular. All the reactions discussed earlier are examples.
2. During the thermal decomposition of complex compounds, which include an oxidizing agent and a reducing agent in the form of atoms of different elements, redox reactions occur, called intramolecular:
3. Reactions disproportionation (dismutation or, according to outdated terminology, self-oxidation - self-healing) can occur if compounds containing elements in intermediate oxidation states are exposed to conditions where they are unstable (for example, at elevated temperatures). The oxidation state of this element both increases and decreases:
4. Reactions counter-proportionation (switching) are processes of interaction between an oxidizing agent and a reducing agent, which include the same element with different oxidation states. As a result, the product of oxidation and reduction is a substance with an intermediate oxidation state of the atoms of a given element:
There are also mixed reactions. For example, the intramolecular counterproportionation reaction includes the decomposition reaction of ammonium nitrate
Drawing up equations.
Equations for redox reactions are compiled based on the principles of equality of the number of the same atoms before and after the reaction, as well as taking into account the equality of the number of electrons given up by the reducing agent and the number of electrons accepted by the oxidizing agent, i.e. electrical neutrality of molecules. The reaction is represented as a system of two half-reactions - oxidation and reduction, the summation of which, taking into account the indicated principles, leads to the compilation of a general equation for the process.
To compile equations for redox reactions, the method of electron-ion half-reactions and the electron balance method are most often used.
Electron-ion half-reaction method used in drawing up equations for reactions occurring in an aqueous solution, as well as reactions involving substances whose oxidation state of elements is difficult to determine (for example, KNCS, CH 3 CH 2 OH).
According to this method, the following main stages are distinguished in composing the reaction equation.
a) write down the general molecular diagram of the process, indicating the reducing agent, oxidizing agent and the medium in which the reaction occurs (acidic, neutral or alkaline). For example
b) taking into account the dissociation of electrolytes in an aqueous solution, this scheme is presented in the form of molecular-ion interaction. Ions whose oxidation states of atoms do not change are not indicated in the diagram, with the exception of environmental ions (H +, OH -):
c) determine the oxidation degrees of the reducing agent and oxidizing agent, as well as the products of their interaction:
f) add ions that did not participate in the oxidation-reduction process, equalize their amounts on the left and right, and write down the molecular equation of the reaction
The greatest difficulties arise when compiling a material balance for half-reactions of oxidation and reduction, when the number of oxygen atoms that make up the particles of the oxidizer and reducer changes. It should be taken into account that in aqueous solutions the binding or addition of oxygen occurs with the participation of water molecules and ions of the medium.
During the oxidation process, for one oxygen atom that attaches to a reducing agent particle, in acidic and neutral environments, one molecule of water is consumed and two H + ions are formed; in an alkaline environment, two hydroxide ions OH - are consumed and one molecule of water is formed (Table 1.1).
To bind one oxygen atom of the oxidizing agent in an acidic environment, two H + ions are consumed during the reduction process and one water molecule is formed; in neutral and alkaline environments, one H 2 O molecule is consumed and two OH - ions are formed (Tables 1, 2).
Table 1
Addition of oxygen atoms to a reducing agent during oxidation
table 2
Bonding of oxygen atoms of the oxidizing agent during the reduction process
The advantages of the method of electron-ionic half-reactions are that when compiling equations for redox reactions, the real states of particles in solution and the role of the environment in the course of processes are taken into account; there is no need to use the formal concept of oxidation state.
Electronic balance method, based on taking into account changes in the oxidation state and the principle of electrical neutrality of the molecule, is universal. It is usually used to construct equations for redox reactions occurring between gases, solids, and in melts.
The sequence of operations, according to the method, is as follows:
1) write down the formulas of the reagents and reaction products in molecular form:
2) determine the oxidation state of atoms that change it during the reaction:
3) based on the change in oxidation states, the number of electrons given up by the reducing agent and the number of electrons accepted by the oxidizing agent are determined, and an electronic balance is drawn up, taking into account the principle of equality of the number of electrons given up and received:
4) the electronic balance factors are written into the equation of the redox reaction as the main stoichiometric coefficients:
5) select the stoichiometric coefficients of the remaining participants in the reaction:
When drawing up equations, it should be taken into account that the oxidizing agent (or reducing agent) can be consumed not only in the main redox reaction, but also when binding the resulting reaction products, that is, it can act as a medium and a salt former.
An example when the role of the medium is played by an oxidizing agent is the oxidation reaction of a metal in nitric acid, composed by the method of electron-ionic half-reactions:
An example when the reducing agent is the medium in which the reaction occurs is the oxidation of hydrochloric acid with potassium dichromate, compiled by the electronic balance method:
When calculating the quantitative, mass and volume ratios of participants in redox reactions, the basic stoichiometric laws of chemistry and, in particular, the law of equivalents are used. To determine the direction and completeness of redox processes, the values of the thermodynamic parameters of these systems are used, and when reactions occur in aqueous solutions, the values of the corresponding electrode potentials are used.
During the lesson we will study the topic “Oxidation-reduction reactions”. You will learn the definition of these reactions, their differences from other types of reactions. Remember what oxidation number, oxidizing agent and reducing agent are. Learn to draw up electronic balance diagrams for redox reactions, get acquainted with the classification of redox reactions.
Topic: Redox reactions
Lesson: Redox Reactions
Reactions that occur with a change in the oxidation states of the atoms that make up the reacting substances are called redox . The change in oxidation states occurs due to the transfer of electrons from the reducing agent to the oxidizing agent. is the formal charge of an atom, assuming that all bonds in the compound are ionic.
Oxidizer - This is a substance whose molecules or ions accept electrons. If an element is an oxidizing agent, its oxidation state decreases.
O 0 2 + 4e - → 2O -2 (Oxidizing agent, reduction process)
Process reception substances of electrons is called restoration. The oxidizing agent is reduced during the process.
Reductant - a substance whose molecules or ions donate electrons. The reducing agent has an increased oxidation state.
S 0 -4e - →S +4 (Reductant, oxidation process)
Process returns electrons is called . The reducing agent is oxidized during the process.
Example No. 1. Obtaining chlorine in the laboratory
In the laboratory, chlorine is obtained from potassium permanganate and concentrated hydrochloric acid. Potassium permanganate crystals are placed in a Wurtz flask. Close the flask with a stopper with a dropping funnel. Hydrochloric acid is poured into the funnel. Hydrochloric acid is poured from a dropping funnel. Vigorous release of chlorine begins immediately. Through the gas outlet tube, chlorine gradually fills the cylinder, displacing air from it. Rice. 1.
Rice. 1
Using this reaction as an example, let's consider how to draw up an electronic balance.
KMnO 4 + HCI = KCI + MnCI 2 + CI 2 + H 2 O
K + Mn +7 O -2 4 + H + CI - = K + CI - + Mn +2 CI - 2 + CI 0 2 + H + 2 O -2
The oxidation states changed manganese and chlorine.
Mn +7 +5e - = Mn +2 oxidizing agent, reduction process
2 CI - -2e - \u003d CI 0 2 reducing agent, oxidation process
4. Equalize the number of given and received electrons. To do this, we find the least common multiple for the numbers 5 and 2. This is 10. As a result of dividing the least common multiple by the number of electrons given and accepted, we find the coefficients of the oxidizing agent and the reducing agent.
Mn +7 +5e - = Mn +2 2
2 CI - -2е - = CI 0 2 5
2KMnO 4 + ? HCI = ?KCI + 2MnCI 2 + 5CI 2 +? H2O
However, there is no coefficient in front of the hydrochloric acid formula, since not all chloride ions participated in the redox process. The electron balance method allows you to balance only the ions involved in the redox process. Therefore, it is necessary to equalize the number of ions not participating in . Namely potassium cations, hydrogen and chloride anions. The result is the following equation:
2KMnO 4 + 16 HCI = 2KCI + 2MnCI 2 + 5CI 2 + 8H 2 O
Example No. 2. Interaction of copper with concentrated nitric acid. Rice. 2.
A “copper” coin was placed in a glass with 10 ml of acid. The release of brown gas quickly began (brown bubbles in the still colorless liquid looked especially impressive). The entire space above the liquid turned brown, and brown vapors poured out of the glass. The solution turned green. The reaction was constantly accelerating. After about half a minute the solution turned blue, and after two minutes the reaction began to slow down. The coin did not completely dissolve, but lost a lot in thickness (it could be bent with your fingers). The green color of the solution in the initial stage of the reaction is due to the reduction products of nitric acid.
Rice. 2
1. Let's write down the scheme of this reaction:
Cu + HNO 3 = Cu (NO 3) 2 + NO 2 + H 2 O
2. Let us arrange the oxidation states of all elements in the substances participating in the reaction:
Cu 0 + H + N +5 O -2 3 = Cu +2 (N +5 O -2 3) 2 + N +4 O -2 2 + H + 2 O -2
The oxidation states changed for copper and nitrogen.
3. We draw up a diagram reflecting the process of electron transition:
N +5 +е - = N +4 oxidizing agent, reduction process
Cu 0 -2е - = Cu +2 reducing agent, oxidation process
4. Let's equalize the number of given and received electrons. To do this, we find the least common multiple for the numbers 1 and 2. This is 2. As a result of dividing the least common multiple by the number of electrons given and received, we find the coefficients of the oxidizing agent and the reducing agent.
N +5 +e - = N +4 2
Cu 0 -2е - = Cu +2 1
5. We transfer the coefficients to the original diagram and transform the reaction equation.
Cu + ?HNO3 = Cu(NO3)2 + 2NO2 + 2H2O
Nitric acid is involved not only in the redox reaction, so the coefficient is not written at first. As a result, the following equation is finally obtained:
Cu + 4HNO 3 = Cu (NO 3) 2 + 2NO 2 + 2H 2 O
Classification of redox reactions
1. Intermolecular redox reactions .
These are reactions in which the oxidizing and reducing agents are different substances.
H 2 S -2 + Cl 0 2 → S 0 + 2HCl -
2. Intramolecular reactions in which oxidizing and stopping atoms are in the molecules of the same substance, for example:
2H + 2 O -2 → 2H 0 2 + O 0 2
3. Disproportionation (self-oxidation-self-recovery) - reactions in which the same element acts both as an oxidizing agent and as a reducing agent, for example:
Cl 0 2 + H 2 O → HCl + O + HCl -
4. Conproportionation (Reproportionation) - reactions in which one oxidation state is obtained from two different oxidation states of the same element
Homework
1. No. 1-3 (p. 162) Gabrielyan O.S. Chemistry. Grade 11. A basic level of. 2nd ed., erased. - M.: Bustard, 2007. - 220 p.
2. Why does ammonia exhibit only reducing properties, and nitric acid only oxidizing properties?
3. Arrange the coefficients in the reaction equation for the production of nitric acid using the electronic balance method: ?NO 2 + ?H 2 O + O 2 = ?HNO 3
Lesson type. Acquiring new knowledge.
Lesson objectives.Educational. Introduce students to a new classification of chemical reactions based on changes in the oxidation states of elements - oxidation-reduction reactions (ORR); teach students to arrange coefficients using the electronic balance method.
Developmental. Continue the development of logical thinking, ability to analyze and compare, and develop interest in the subject.
Educational. To form the scientific worldview of students; improve work skills.
Methods and methodological techniques. Story, conversation, demonstration of visual aids, independent work of students.
Equipment and reagents. Reproduction with the image of the Colossus of Rhodes, algorithm for arranging coefficients using the electronic balance method, table of typical oxidizing and reducing agents, crossword puzzle; Fe (nail), NaOH, CuSO 4 solutions.
DURING THE CLASSES
Introduction
(motivation and goal setting)
Teacher. In the 3rd century. BC. On the island of Rhodes, a monument was built in the form of a huge statue of Helios (the Greek god of the Sun). The grandiose design and perfect execution of the Colossus of Rhodes - one of the wonders of the world - amazed everyone who saw it.
We don't know exactly what the statue looked like, but it is known that it was made of bronze and reached a height of about 33 m. The statue was created by the sculptor Haret and took 12 years to build.
The bronze shell was attached to an iron frame. The hollow statue began to be built from the bottom and, as it grew, it was filled with stones to make it more stable. About 50 years after its completion, the Colossus collapsed. During the earthquake it broke at the level of the knees.
Scientists believe that the real reason for the fragility of this miracle was metal corrosion. And the corrosion process is based on redox reactions.
Today in the lesson you will learn about redox reactions; learn about the concepts of “reducing agent” and “oxidizing agent”, about the processes of reduction and oxidation; learn to place coefficients in equations of redox reactions. Write down the date and topic of the lesson in your workbooks.
Learning new material
The teacher performs two demonstration experiments: the interaction of copper(II) sulfate with alkali and the interaction of the same salt with iron.
Teacher. Write down the molecular equations of the reactions performed. In each equation, arrange the oxidation states of the elements in the formulas of the starting substances and reaction products.
The student writes reaction equations on the board and assigns oxidation states:
Teacher. Did the oxidation states of the elements change in these reactions?
Student. In the first equation, the oxidation states of the elements did not change, but in the second they changed - for copper and iron.
Teacher. The second reaction is redox. Try to define redox reactions.
Student. Reactions that result in changes in the oxidation states of the elements that make up the reactants and reaction products are called redox reactions.
Students write down in their notebooks, under the teacher’s dictation, the definition of redox reactions.
Teacher. What happened as a result of the redox reaction? Before the reaction, iron had an oxidation state of 0, after the reaction it became +2. As we can see, the oxidation state has increased, therefore, iron gives up 2 electrons.
Copper has an oxidation state of +2 before the reaction, and 0 after the reaction. As we can see, the oxidation state has decreased. Therefore, copper accepts 2 electrons.
Iron donates electrons, it is a reducing agent, and the process of transferring electrons is called oxidation.
Copper accepts electrons, it is an oxidizing agent, and the process of adding electrons is called reduction.
We write the schemes of these processes:
So, give the definition of the concepts of "reducing agent" and "oxidizing agent".
Student. Atoms, molecules or ions that donate electrons are called reducing agents.
Atoms, molecules or ions that gain electrons are called oxidizing agents.
Teacher. How can we define the processes of reduction and oxidation?
Student. Reduction is the process by which an atom, molecule, or ion gains electrons.
Oxidation is the process of transfer of electrons by an atom, molecule or ion.
Students write down definitions from dictation in a notebook and draw.
Remember!
Donate electrons and oxidize.
Take electrons - recover.
Teacher. Oxidation is always accompanied by reduction, and vice versa, reduction is always associated with oxidation. The number of electrons given up by the reducing agent is equal to the number of electrons gained by the oxidizing agent.
To select coefficients in the equations of redox reactions, two methods are used - electronic balance and electron-ion balance (half-reaction method).
We will consider only the electronic balance method. To do this, we use an algorithm for arranging coefficients using the electronic balance method (designed on a piece of Whatman paper).
EXAMPLE Arrange the coefficients in this reaction scheme using the electronic balance method, determine the oxidizing agent and the reducing agent, indicate the processes of oxidation and reduction:
Fe 2 O 3 + CO Fe + CO 2.
We will use the algorithm for placing the coefficients using the electronic balance method.
3. Let’s write down the elements that change oxidation states:
4. Compose electronic equations, determining the number of given and received electrons:
5. The number of given and received electrons must be the same, because Neither the starting materials nor the reaction products are charged. We equalize the number of given and received electrons by choosing the least common multiple (LCM) and additional factors:
6. The resulting multipliers are coefficients. Let's transfer the coefficients to the reaction scheme:
Fe 2 O 3 + 3CO = 2Fe + 3CO 2.
Substances that are oxidizing or reducing agents in many reactions are called typical.
A table made on a piece of Whatman paper is hung.
Teacher. Redox reactions are very common. They are associated not only with corrosion processes, but also with fermentation, decay, photosynthesis, and metabolic processes occurring in a living organism. They can be observed during fuel combustion. Redox processes accompany the cycles of substances in nature.
Did you know that about 2 million tons of nitric acid are formed in the atmosphere every day, or
700 million tons per year, and in the form of a weak solution fall to the ground with rain (man produces only 30 million tons of nitric acid per year).
What is happening in the atmosphere?
Air contains 78% by volume nitrogen, 21% oxygen and 1% other gases. Under the action of lightning discharges, and an average of 100 lightning flashes on Earth every second, nitrogen molecules interact with oxygen molecules to form nitric oxide (II):
Nitric oxide(II) is easily oxidized by atmospheric oxygen to nitric oxide(IV):
NO + O 2 NO 2 .
The resulting nitric oxide (IV) interacts with atmospheric moisture in the presence of oxygen, turning into nitric acid:
NO 2 + H 2 O + O 2 HNO 3.
All these reactions are redox.
Exercise . Arrange the coefficients in the given reaction schemes using the electronic balance method, indicate the oxidizing agent, the reducing agent, the processes of oxidation and reduction.
Solution
1. Let's determine the oxidation states of elements:
2. Let us emphasize the symbols of elements whose oxidation states change:
3. Let’s write down the elements that have changed their oxidation states:
4. Let’s create electronic equations (determine the number of given and received electrons):
5. The number of electrons given and received is the same.
6. Let’s transfer the coefficients from the electronic circuits to the reaction diagram:
Next, students are asked to independently arrange the coefficients using the electronic balance method, determine the oxidizing agent, the reducing agent, and indicate the processes of oxidation and reduction in other processes occurring in nature.
The other two reaction equations (with coefficients) have the form:
The correctness of tasks is checked using an overhead projector.
Final part
The teacher asks students to solve a crossword puzzle based on the material they have studied. The result of the work is submitted for verification.
Having solved crossword, you will learn that the substances KMnO 4, K 2 Cr 2 O 7, O 3 are strong ... (vertical (2)).
Horizontally:
1. What process does the diagram reflect:
3. Reaction
N 2 (g.) + 3H 2 (g.) 2NH 3 (g.) + Q
is redox, reversible, homogeneous, ....
4. ... carbon(II) is a typical reducing agent.
5. What process does the diagram reflect:
6. To select coefficients in the equations of redox reactions, use the electronic... method.
7. According to the diagram, aluminum gave up ... an electron.
8. In reaction:
H 2 + Cl 2 = 2HCl
hydrogen H 2 – ... .
9. What type of reactions are always redox only?
10. The oxidation state of simple substances is….
11. In reaction:
reducing agent –….
Homework assignment. According to the textbook by O.S. Gabrielyan “Chemistry-8” § 43, p. 178–179, ex. 1, 7 in writing.
Task (for home). The designers of the first spaceships and submarines were faced with a problem: how to maintain a constant air composition on the ship and space stations? Get rid of excess carbon dioxide and replenish oxygen? A solution has been found.
Potassium superoxide KO 2, as a result of interaction with carbon dioxide, forms oxygen:
As you can see, this is a redox reaction. Oxygen in this reaction is both an oxidizing agent and a reducing agent.
On a space mission, every gram of cargo counts. Calculate the supply of potassium superoxide that must be taken on a space flight if the flight lasts 10 days and if the crew consists of two people. It is known that a person exhales 1 kg of carbon dioxide per day.
(Answer: 64.5 kg KO 2. )
Assignment (increased level of difficulty). Write down the equations of redox reactions that could lead to the destruction of the Colossus of Rhodes. Keep in mind that this giant statue stood in a port city on an island in the Aegean Sea, off the coast of modern-day Turkey, where the humid Mediterranean air is laden with salts. It was made of bronze (an alloy of copper and tin) and mounted on an iron frame.
Literature
Gabrielyan O.S.. Chemistry-8. M.: Bustard, 2002;
Gabrielyan O.S., Voskoboynikova N.P., Yashukova A.V. Teacher's handbook. 8th grade. M.: Bustard, 2002;
Cox R., Morris N. Seven wonders of the world. The ancient world, the Middle Ages, our time. M.: BMM AO, 1997;
Small children's encyclopedia. Chemistry. M.: Russian Encyclopedic Partnership, 2001; Encyclopedia for children "Avanta+". Chemistry. T. 17. M.: Avanta+, 2001;
Khomchenko G.P., Sevastyanova K.I. Redox reactions. M.: Education, 1989.
How do you know where the oxidizing agent is and where the reducing agent is in a chemical reaction? and got the best answer
Reply from ul.[active]
if after a reaction (after the equal sign) a substance acquires a positive charge, it means it is a reducing agent
and if it acquires a negative charge, it means it is an oxidizing agent
For example
H2 + O2 = H2O
Before the reaction, both hydrogen and oxygen have zero charge
after the reaction
hydrogen acquires a charge of +1 and oxygen -2 means hydrogen is a reducing agent
and oxygen is an oxidizing agent!!
Source: =)) if anything is unclear, write)
Answer from 2 answers[guru]
Hello! Here is a selection of topics with answers to your question: How do you know where in a chemical reaction the oxidizing agent is and where the reducing agent is?
Answer from BeardMax[guru]
To do this, you need to know what the degree of oxidation is.
Learn to determine the oxidation state of any atom in a chemical compound.
Next, look at which atoms of CO increases in the reaction, and which decreases. The former are reducing agents, the latter are oxidizing agents.
In general, chemistry was not necessary to skip.
Answer from OOO[newbie]
A reducing agent is a substance that donates electrons. For example, Ca (2+) - 2e = Ca (0)
An oxidizing agent is a substance that accepts electrons.
Answer from Mariska[newbie]
To find out, you need to look at what are reagents and what is added as a medium. For example, if the starting substances contain Mn (+4) and water, then Mn will change the oxidation state to (+6), if I’m not mistaken. In addition, you can see in what degree of oxidation the elements are (suddenly somewhere it is minimal or, on the contrary, maximum).
