Redox reactions in qualitative analysis. Application of redox reactions in chemical analysis Redox reactions in qualitative analysis

The main subject of study in chemistry is the chemical reaction. A deep knowledge of the nature and patterns of chemical reactions makes it possible to control them and use them for the synthesis of new substances. The assimilation of the general patterns of the course of chemical reactions is necessary for the subsequent study of the properties of inorganic and organic substances. Most of the chemical processes occurring in nature and carried out by man in his practical activities are redox reactions. The following processes can be cited as an example: photosynthesis, natural gas combustion, obtaining metals from ores, ammonia synthesis. The equation of any chemical reaction can be written on paper, but this does not mean that such a reaction is really possible. In some cases, it is enough to change the conditions for the reaction to go on, in others, the reaction cannot be forced to go under any conditions. The same can be attributed to redox reactions, the course of which is largely determined by factors that affect the course of a chemical reaction.

The applications of redox reactions are enormous. They are used in the qualitative and quantitative analysis of substances, biochemical processes, to obtain new important substances. The equations of these reactions are interesting for their complexity and peculiarity.

I set myself the following tasks: to study the relevant scientific literature on this topic, to select for the practical part a methodology that is possible for school conditions, to conduct a series of analyzes according to the selected methodology, to draw appropriate conclusions on this topic.

The work is carried out by order of the administration high school №16.

1. 1. BASIC CONCEPTS AND FACTORS AFFECTING REDOX REACTIONS

1. 1. 1. Basic concepts and definitions.

Along with the acid-base interaction, which is based on the exchange of a proton (H +) between the reagents, redox interaction is widespread in nature, which is characterized by the redistribution of electrons between the reagents.

Redox reactions are called chemical reactions that occur with a change in the oxidation state of atoms due to the redistribution of electrons between them.

In redox reactions, electrons are always given off and added.

Oxidation is the process of giving off electrons by an atom of a substance, accompanied by an increase in the degree of its oxidation.

For example:

– = – = – =

Recovery is the process of adding electrons to an atom of a substance, accompanied by a decrease in the degree of its oxidation.

For example:

In the course of a redox reaction, both processes proceed simultaneously, and the total number of electrons given up during oxidation is equal to the total number of electrons received during reduction.

An oxidizing agent is a substance that includes atoms that attach electrons, that is, an oxidizing agent is an electron acceptor.

A reducing agent is a substance that includes atoms that donate electrons, that is, a reducing agent is an electron donor.

The oxidizing agent, accepting electrons, acquires reducing properties, turning into a conjugated reducing agent: oxidizing agent + e conjugated reducing agent

The reducing agent, donating electrons, acquires oxidizing properties, turning into a conjugated oxidizing agent: reducing agent - e conjugated oxidizing agent

Any redox reaction is a unity of two opposite transformations - oxidation and reduction: conjugated redox pair I conjugated conjugated oxidizing agent I + reducing agent II reducing agent I oxidizing agent II conjugated redox pair II

For example: conjugated redox couple I

Conjugated redox couple II

The combination of an oxidizing agent (reducing agent) with the product of its transformation constitutes a conjugated redox pair, and its interconversion is a reduction (oxidation) half-reaction. In any redox reaction, two conjugated redox pairs take part, between which there is competition for electrons, as a result of which two half-reactions occur: one is associated with the addition of electrons, i.e. reduction, the other with the release of electrons, i.e. e. oxidation.

So, in the reaction:

Pairs are involved, and, and the indicated half-reactions proceed:

And in the reaction:

+ + + ; pairs are involved, and, and half-reactions proceed:

For correct notation of the conjugated redox pair, the oxidized form should first be written, and then the reduced form of the substance.

The interaction of substances in redox reactions, as in other chemical reactions, obeys the law of equivalents.

The equivalent of an oxidizing or reducing agent is its particle (real or conditional), which, respectively, attaches or gives up one electron.

The molar mass equivalent 1/z M(X) of the oxidizing or reducing agent is equal to their molar mass M(X) multiplied by their equivalence factor 1/z in the given reaction.

The equivalence factor of the oxidizing agent or reducing agent is equal to 1/z, where z is the number of electrons accepted or given away by one particle (molecule, atom, ion) of the oxidizing or reducing agent. Therefore, the molar mass equivalent of an oxidizing agent or reducing agent is calculated by the equation:

M(1/z X) = M(X)/z

1. 1. 2. Factors affecting the course of redox reactions. The nature of the course of redox reactions depends on the chemical nature of the interacting substances and on the reaction conditions:

Reagent concentrations:

+ (dilute) + oxidizing agent;

+ (conc) + + oxidizing agent;

Reaction temperatures: cold + + + ; when heated + + + ;

Presence of catalyst: without catalyst + + ; with catalyst + + ;

Influence of the nature of the environment: acidic environment

+ - + ; neutral environment

+ - +; alkaline environment

As can be seen from the analysis of the above half-reactions, redox reactions occurring in aqueous solutions often involve water or its ions, which not only facilitate the transfer of electrons from the reducing agent to the oxidizing agent, but also bind the components of these transformations. This is especially important for biochemical redox reactions that always occur in aqueous solutions.

The basis for determining the direction of the spontaneous occurrence of redox reactions is the following rule:

Redox reactions spontaneously proceed always towards the transformation of a strong oxidizing agent into a weak conjugated reducing agent or a strong reducing agent into a weak conjugated oxidizing agent.

This rule is similar to the rule that determines the direction of acid-base transformations.

A quantitative measure of the redox ability of a given conjugated redox pair is the value of its reduction potential, which depends on:

The nature of the oxidized and reduced form of a given conjugate pair;

The concentration ratios of the oxidized and reduced forms of a given conjugated pair;

Temperatures.

In cases where ions or are involved in the conversion of an oxidizing agent or reducing agent, it also depends on the pH of the solution. The value that takes under standard conditions: the concentration of all components involved in the reaction, including water ions (in an acidic environment) and (in an alkaline environment), is 1 mol / l, a temperature of 298 K, is called the standard reduction potential and is indicated. The value is a quantitative characteristic of the redox properties of a given conjugated redox pair under standard conditions.

There is no way to determine the absolute value of potentials for conjugated redox pairs. Therefore, relative values ​​are used that characterize the potentials of conjugated pairs relative to the reference pair + , the potential of which, under standard conditions, is assumed to be conditionally equal to zero ((,) =).

Redox pairs have a positive value, in which the oxidized form attaches electrons more easily than the hydrogen cation in the reference pair. Redox pairs have a negative value, in which the oxidized form attaches electrons more difficult than in the reference pair. Consequently, the greater (i.e., more positive) the value of a given conjugated redox pair, the more pronounced its oxidizing properties, and the reducing properties, respectively, are weaker.

The essence of redox reactions is the competition for the addition of an electron between the participating oxidizing agents. In this case, the electron is attached to that conjugated pair, the oxidized form of which holds it stronger. This is reflected in the following diagram:

restore I ok-l I + + ok-l II restore II

Redox shift

Comparing the potentials of conjugated pairs involved in the redox reaction, it is possible to determine in advance the direction in which this or that reaction will spontaneously proceed.

When two conjugated redox pairs interact, the oxidizing agent will always be the oxidized form of the pair whose potential has a more positive value.

Example. The reaction mixture contains two conjugated redox pairs:

Since the first pair contains a stronger oxidizing agent () than the second pair (), then under standard conditions a reaction will spontaneously proceed in which it will be an oxidizing agent and a reducing agent: + = +.

To determine the direction of the redox reaction, you can also use the value of its EMF.

The EMF of a redox reaction under standard conditions (E°) is numerically equal to the difference in the standard potentials of the conjugated redox pairs involved in the reaction: E° = –.

The condition for the spontaneous occurrence of a redox reaction is the positive value of its EMF, i.e. E ° \u003d -\u003e 0.

Taking into account this condition, for a spontaneous redox reaction, the value of the redox pair acting as an oxidizing agent should be more than a second redox pair, which plays the role of a reducing agent in this reaction. So, in the example above:

E° = - = 0.54 - 0.17 = 0.37B

If E°=0, then the redox reaction is equally likely to occur both in the forward and reverse directions, and this is the condition for the occurrence of chemical equilibrium for the redox process. The quantitative characteristic of the course of any reversible processes is the equilibrium constant K, which is related to the change in the standard Gibbs energy by the following relation:

∆G = -2.3RT log K

On the other hand, the change in the standard Gibbs energy is related to the EMF of the redox reaction by the relation:

ΔG = –zFE where F = 96 500 C/mol; z is the number of electrons involved in the elementary process.

From these two equations follows:

E ° \u003d (2.3RT log K) / (zF) or log K \u003d (zFE) / (2.3RT)

Using these expressions, it is possible to calculate the equilibrium constant of any redox reaction, but it will have a real value only for those reactions whose EMF is less than 0.35 V, since at high EMF the reactions are considered practically irreversible. Since the EMF of individual stages of redox reactions occurring in living systems usually does not exceed 0.35 V (E °

Redox reactions underlie the metabolism of any organisms. In the case of aerobic metabolism, the main oxidizing agent is molecular oxygen supplied during respiration, and the reducing agent is organic compounds supplied with food. In anaerobic metabolism, it is based mainly on redox reactions, in which organic compounds are both oxidizing and reducing agents.

1. 3. FEATURES OF BIOCHEMICAL REDOX PROCESSES IN ORGANISMS

All biochemical redox processes, the speed and depth of which are controlled by the body, are carried out in the presence of enzymes with the general name of oxidoreductase. Oxidoreductases always contain cofactors or coenzymes. Cofactors are transition metal cations (usually iron and copper, and sometimes manganese and molybdenum), which form a complex compound with the enzyme protein. Coenzymes are complex organic compounds that are quite tightly bound to the enzyme protein. The main feature of cofactors and coenzymes is their ability to be both an oxidizing agent and a reducing agent, since each of them can be in two conjugated forms: oxidized and reduced. Thus, oxidoreductases exhibit redox properties at the expense of their cofactors or coenzymes.

The composition of the final products of the breakdown of molecules of complex organic compounds under aerobic and anaerobic conditions, as can be seen from the above reactions, differs sharply in the case of carbon-containing products. During aerobic oxidation, these products contain only, and under anaerobic conditions, carbon-containing products are formed, in which carbon atoms can have a wide range of oxidation states:

products of aerobic digestion (excess of oxygen) initial organic substances ++++++ products of anaerobic digestion (lack of oxygen)

It can be seen from the diagram that during the biological oxidation of organic compounds, only the oxidation states of their constituent carbon atoms change, since the oxidation states of all other atoms (hydrogen, nitrogen, and sulfur) remain constant. Therefore, when describing the transformations of organic compounds, it is necessary to take into account the degree of oxidation of each carbon atom.

1. 3. 1. Stepwise biochemical redox reactions. A feature of biochemical oxidation-reduction reactions is their multistage nature: the formation of many different intermediate products. Moreover, all biochemical redox processes: glycolysis, β-oxidation of fatty acids, the Krebs cycle, oxidative phosphorylation and others - include many different stages, each of which is carried out under the action of certain enzymes. All the necessary enzymes for each stage of this process are combined into ensembles with a clear spatial organization due to intermolecular bonds. Ensembles of enzymes, as a rule, are fixed on various cell membranes. As a result of the action of all enzymes of the ensemble coordinated in time and space, the chemical transformations of the substrate are carried out gradually, as if on a conveyor belt. In this case, the reaction product of one stage is the starting compound for the next stage.

The stepwise mechanism of the occurrence of biochemical redox reactions creates the possibility of functional control of the formation of substances at each stage.

Thus, the body prevents changes in the metabolic process that are undesirable for it, which, if they occur, are due to the multi-stage process not abruptly, but smoothly, gradually. Due to the gradation for biochemical redox reactions, the reversibility of individual stages becomes more likely. This provides a reaction directed towards equilibrium, spontaneous flow and helps to maintain redox homeostasis in the body.

1. 3. 2. Normal recovery potential. One of the features of biochemical redox processes in living organisms is that most of them proceed in neutral aquatic environment. Therefore, to characterize the redox properties of natural conjugated pairs, instead of standard potential values ​​that correspond to pH = 0 or pH = 14 (pOH = 0), normal values ​​of reduction potentials measured at 1 M concentration of components and at pH = 7.0 are used. Under these conditions, the value of the potential of the hydrogen electrode (,) \u003d - 0.42 V, and the ratio between the values ​​​​of the normal and standard reduction potentials of the compound is expressed by the equation: \u003d - 0.42.

All natural conjugated redox pairs have potentials in the range of -0.42 + 0.82 V, which characterize the electrochemical stability of water. At a potential below -0.42 V, the reduction of water begins with the formation of molecular hydrogen, and at a potential above + 0.82 V, water is oxidized with the formation of molecular oxygen.

1. 3. 3. Exergonicity of biological oxidation reactions. Biological oxidation reactions are exergonic and serve as sources of energy necessary for various life processes. Therefore, many reactions of biological oxidation are associated with the phosphorylation of ADP with the formation of ATP, as a result, the accumulation of energy of redox reactions in the body occurs. At the same time, in the body, some oxidation reactions are associated with reduction reactions, which are endergonic.

1. 3. 4. Classification of biochemical redox processes. In the biochemical literature, it is widely believed that biological oxidation is associated only with the interaction of the substrate with oxygen and its derivatives or with the dehydrogenation of the substrate. In fact, in addition to these processes, very big number biochemical reactions occurring intramolecularly in the presence of enzymes, has a redox character. Therefore, the redox processes occurring in the body are proposed to be divided into the following types:

1. Reactions of intra- and intermolecular redox dismutation due to carbon atoms.

2. Reactions occurring with the participation of oxidoreductases:

Dehydrogenase oxidation-reduction;

Oxygenase oxidation-reduction.

3. Free radical oxidation-reduction.

1. 4. USE OF OXIDIZERS AND REDUCERS IN HEALTH PRACTICE

Many strong oxidizing agents: potassium permanganate KMnO4, hydrogen peroxide H2O2, iodine solution, bleach CaClOCl, as well as chlorine and ozone (for chlorination and ozonation of water) are widely used as bactericidal agents in health care practice, because due to strong oxidizing properties they effectively destroy microorganisms.

The toxic effect of nitrogen oxides, ozone, chlorine, bromine, nitrates and nitrites, chromates and dichromates is associated with their oxidizing properties. In case of poisoning by oxidizing or reducing agents, redox reactions are used to neutralize them. So, in case of poisoning with hydrogen sulfide (a strong reducing agent), the victim is allowed to breathe slightly moistened bleach, from which small amounts of chlorine are released, while the reaction proceeds:

In case of poisoning with bromine vapor (a strong oxidizing agent), ammonia vapor is allowed to be inhaled:

The use of various reducing agents as antioxidants in medicinal and prophylactic products was considered in the previous section.

Redox reactions underlie the methods of oxidimetry: permanganatometry, iodometry, chromatometry, which are widely used in clinical analysis to determine calcium ions, uric acid, cholesterol, sugar, catalase and peroxidase enzymes in the blood. In sanitary and hygienic practice, these methods are used to determine the oxidizability of water, the content of "residual" chlorine in household and drinking water, as well as "active" chlorine in disinfectants (bleach and chloramines).

Thus, redox processes in the body play an extremely important role, supplying it with energy and necessary metabolites, as well as participating in the regulatory mechanisms of life. Due to neurohumoral regulation, an amazing balance is achieved between the content of oxidizing agents, reducing agents and products of their interaction in living organisms, which provides them with a state of redox homeostasis.

1. 5. REDOX

TITRATION (REDOXMETRY)

Theoretical basis

1. 5. 1. Methods of redoxmetry. Redox reactions formed the basis of a group of titrimetric analysis methods, which are called redox titration methods, or redoxmetry. Depending on the type of reactions and working solutions, the following methods are distinguished: permanganometric titration - titrant solution; iodometric titration - based on interaction with iodine; bromatometric titration - titrant solution; dichromatometric titration - titrant solution; nitritometric titration - titrant solution. As titrants in redox titrations, solutions of oxidizing and reducing agents are used. Oxidizing and reducing agents, depending on the conditions, can enter into reactions in which a different number of electrons is transferred, so often oxidizing and reducing agents do not have a constant equivalent. For example, in an acidic environment - it attaches 5 electrons and is reduced to (E = M / 5 = 31.68), in an alkaline environment - attaches 1 electron and is reduced to - (E = M = 158.04), in a neutral one - attaches 3 electrons and is restored to (E = M / 3 = 52.68).

Permanganatometric titration

main reaction. In permanganometric titration, a titrant is used - a solution of potassium permanganate. Titration with a solution of potassium permanganate is most often carried out in an acidic environment, in the presence of H2SO4, and a half-reaction occurs:

E = 1.52 V; = 31.608.

Titrant preparation. Solutions of KMnO4 are usually prepared in 0.1 N. To prepare 1 dm3 of a solution, a sample of a little more than 3.161 g is taken. KMnO4 is a strong oxidizing agent and is capable of changing its concentration in solutions in the presence of a wide variety of reducing substances - organic ammonia impurities. When preparing solutions of KMnO4, they are kept for several days in a dark place in order to undergo all redox processes with impurities contained in water and trapped in the solution with dust. Only after settling does the concentration of the KMnO4 solution become constant, and the solution is standardized.

Titrant standardization. The KMnO4 solution is standardized for oxalic acid. In a sulfuric acid medium, H2C204 reacts with KMn04 according to the equation:

In this reaction, the equivalents of oxalic acid and potassium permanganate are: =126.04/2=63.04; = 31.608. A portion of oxalic acid for the preparation of 100 cm3 0.1 N. solution is 0.6304 g, for the preparation of 250 cm3 0.1 n. solution of 1.6760 g. The titrant serves as an indicator in the method, with the slightest excess of which the titrated solution turns into pink color, therefore, titration with a KMnO4 solution is carried out until pink color analyzed solution. When titrating with a solution of KMnO4 (and other oxidizing agents), burettes with a glass tap are used. The use of burettes with rubber tubes is unacceptable, since the rubber oxidizes, and the KMnO4 solutions eventually change their concentration. It is not recommended to leave the KMnO4 solution in burettes for a long time, because a brown coating of manganese dioxide appears on the walls of the burettes, which is easily washed off with HCl and H2C2O4 solutions. The concentration of the KMnO4 solution is checked regularly.

Application. Permanganometric titration is used to analyze reducing agents - hydrogen peroxide H2O2, magnesium peroxide MgO2 + MgO, sodium nitrite NaNO2, reduced iron.

For example, the determination of hydrogen peroxide is carried out in an acidic environment: ++ ++(g)+

The exact volume of the drug (10 cm3) is diluted in a volumetric flask to 100 cm3. To 10 cm3 of the resulting solution, add 5 cm3 of diluted sulfuric acid and titrate with 0.1 N. KMnO4 until a faint pink color appears. The titer of the KMnO4 solution for the analyte is TKMnO/n2O2 = 0.001701 g/cm3.

Sometimes a solution of KMnO4 is used to determine substances by back titration. For example, sodium nitrite is analyzed by adding an excess of 0.1 N to its solution. KMnO4 and sulfuric acid. NaNO2 is oxidized by KMnO4 to NaNO3. The excess of KMnO4 is determined iodometrically. KI is added to the solution, KMnO4 oxidizes KI to I2, which is titrated with O.1 N. Na2S2O3 solution:

Iodometric titration

Fundamentals of the method. Iodometry is based on the reduction reaction of free iodine to iodidiones and the oxidation of iodidiones to free iodine (reverse reaction):

+ ; EI/I= +0.535V.

Iodometric titration determines both oxidizing and reducing agents. The reactions are carried out in a neutral environment. In an alkaline or acidic environment, there may be adverse reactions. In an acidic environment, KI forms HI, which is unstable and releases free iodine under the influence of light. In a strongly alkaline environment, I2 forms hypoiodites:

Titrants. Apply solutions of sodium thiosulfate and iodine 0.1 and 0.01 N. concentration. Sodium thiosulfate solutions are used to titrate free iodine.

During this reaction, iodine is reduced to iodide ions, thiosulfate ions are oxidized to tetrathionate ions SO. = 126.9; == 248.2. iodine is slightly soluble in water, its solutions are prepared in the presence of KI, which forms a water-soluble complex:

A portion of iodine is taken slightly more than the calculated one, taking into account the possibility of the presence of impurities, iodine is volatile and can volatilize during weighing, therefore, all operations of weighing and preparing an iodine solution are carried out using weighing bottles and as quickly as possible. A weighed portion of iodine (13 g) is dissolved in a concentrated solution of a triple (39 g of KI per 50 cm3 of water) amount of KI, and after complete dissolution of iodine, the solution is adjusted in a volumetric flask with water to 1000 cm3. The concentration of the prepared iodine solution is set according to the sodium thiosulfate solution, for which 15-20 cm3 of the iodine solution is titrated with O.1 N. thiosulfate solution. A starch solution is used as an indicator, the blue color of which fades when the equivalence point is reached. 0.1 N sodium thiosulfate solution is prepared by weighing a sample of Na2S2O3∙5H2O, approximately equal to 0.1 mol. The sample is dissolved in 500 cm3 of water, 0.1 g of Na2CO3 is added and the mixture is adjusted to 1000 cm3 with water. Na2CO3 is added to stabilize Na2S2O3, which interacts with O2 and CO2 of the air:

In the presence of Na2CO3 these reactions slow down. After preparation, the sodium thiosulfate solution is left for 7-10 days in a dark place so that the concentration of the solution stabilizes. The titer and normality of the Na2S2O3 solution is determined using potassium dichromate. Potassium dichromate quantitatively releases iodine from potassium iodide, the released iodine is titrated with a solution of sodium thiosulfate:

About 0.15 g (accurately weighed) of K2Cr2O7 dried to a constant mass is dissolved in a flask with a ground stopper in 50 cm3 of water, 10 cm3 of a 20% KI solution, 5 cm3 of an HCl solution are added. After 10 minutes of standing in a dark place, 200 cm3 of water is added to the solution and titrated with a Na2S2O3 solution until a green-yellow color is obtained. Then add 2-3 cm3 of starch solution and titrate until the blue color of the solution changes to light green. The volume of the Na2S2O3 solution is used to calculate the titer, normality, or correction factor. Solutions of iodine and sodium thiosulfate are stored in a dark place in closed bottles and their concentration is periodically checked. The indicator of the method is a 1% starch solution, which is prepared by grinding 1 g of starch in a mortar with 5 cm3 of cold water, the mixture is poured into 100 cm3 of boiling water with stirring. After 2–3 minutes of boiling, the solution is cooled. The starch solution deteriorates quickly and can be used for no more than 2-3 days. For preservation, 1 g of salicylic acid or ZnCl2 is added to it. 5 cm3 starch solution should be given with 2 drops of 0.1 N. iodine solution turns blue. If the color turns red-brown, then the solution is not suitable for use. Starch solution cannot be added to concentrated iodine solutions. The blue coloration does not disappear for a long time. The starch solution is added at the end of the titration, when most of the iodine has been titrated and the solution turns yellow.

The use of iodometry. From inorganic compounds, iodine analysis analyzes iodine, potassium permanganate, sodium arsenate, calomel, copper sulfate and many organic drugs (forialin, antshirin, etc.). Iodine and tincture of iodine are analyzed by direct titration with Na2S2O3 solution. The analysis of iodine is carried out as follows. About 0.2 g of iodine, ground in a mortar, is added to a weighed flask with 10 cm3 of a 20% potassium iodide solution, and the flask is weighed again. This method of taking a sample of iodine is adopted in order to avoid volatilization of iodine. The solution is diluted to 20 cm3 and titrated with 0.1 N. sodium thiosulfate solution until yellow. Add 1 cm3 of starch solution, the solution turns blue, titration is carried out until the solution becomes colorless. =0.01269 g/cm3.

Back titration is used in the analysis of Hg2CI2 calomel, formaldehyde, and antipyrine. Calomel is analyzed by adding 0.1 N. iodine solution and potassium iodide. In this case, Hg2 is oxidized to Hg. The excess of unreacted iodine is titrated with 0.1 N. solution of Na2S2O3 *=0.01261g/cm.

By titration of the substituent, potassium permanganate, sodium arsenate, and copper sulfate are determined. For example, KMnO4 is analyzed by dissolving its exact weight in water in a volumetric flask, a 20% solution of KI is added to part of the solution, and after standing for 10 minutes, the released iodine is titrated with 0.1 N. Na2S2O3 solution * = 0.003161 g/cm3.

Bromatometric and bromometric titration

Bromatometric titration. Based on the reduction reaction of bromide ions in an acidic medium to bromide ions:

1.44B; = =27.836.

Titrant. A 0.1 H solution of potassium bromate KBrO is used, which is one of the strong oxidizing agents in an acidic environment. 2.8 g of potassium bromate (accurately weighed) are dissolved in ~500 cm3 of H2O and the volume of the solution in the volumetric flask is adjusted to 1 dm3. The titer and normality of the solution are established iodometrically, by substituent titration. KBrO3 in the presence of HCI quantitatively interacts with KI, releasing elemental iodine I:

The released iodine is titrated with 0.1 N. Na2S2O solution until discoloration in the presence of starch.

The method uses methyl orange and methyl red indicators, which are decolorized at the equivalence point by bromine, which is released during the interaction of potassium bromate with potassium bromide, which is added to the analyzed solution:

Application. Preparations containing arsenic - arsenic anhydride, novarsenol, miarsenol and osarsol are analyzed by bromatometric titration. For example, arsenic anhydride (0.1 g) is dissolved in a solution of sodium hydroxide (2–3 cm3), 50 cm3 of water, concentrated sulfuric acid (10 cm3) and potassium bromide (0.6 g) are added to the solution. The solution is heated to boiling and titrated with 0.1 N. potassium bromate solution. The indicator is methyl red. Titration is carried out until the solution becomes colorless. Arsenic anhydride is oxidized by potassium bromate to HAsO4:

At the equivalence point, excess potassium bromate reacts with potassium bromide, releasing free bromine, which discolors the indicator. = 0.004946 g/cm3.

bromometric titration. Based on the interaction of potassium bromate with potassium bromide. The free bromine released during this reaction has the properties of an oxidizing agent and is able to be reduced to bromide ions:

Elements and bromine quantitatively brominates many organic compounds, especially the group of phenols - phenol, resorcinol, thymol, salicylic acid. For example, the thymol bromination reaction is:

HC-CH-CH HC-CH-CH

During bromatometric titration, a sample of a substance is dissolved in water, a KBr solution is added. The resulting mixture was titrated with 0.1 N. KBrO3 solution. In Russia, substituent titration is used to determine phenol and resorcinol. A solution of a substance is mixed with a certain volume of 0.1 N. KBrO3, add KBr and HCI. The mixture is left for 10 minutes, during which time the substance is brominated by the released bromine. The excess of unreacted bromine is determined by adding a solution of KI to the mixture. Bromine quantitatively displaces iodine from CI, which is titrated with a solution of Na2S2O3.

Nitritometric titration

Fundamentals of nitritemetry. NaNO with organic amines forms diazo compounds:

Diazotization of organic amines with sodium nitrite is quantitative and widely used in analysis. medicines having an amino group: - streptocide, sulfacyl, norsulfazol, etazol, sulfadimezin, anesthesin, novocaine and others.

Titrant. 0.5 M and 0.1 M NaNO solutions are used, which are prepared by dissolving 36.5 g of NaNO2 (for a 0.5 M solution) or 7.3 g (for a 0.1 M solution) in 100 cm3 of water. The NaNO2 solution is standardized for sulfanilic acid:

NaNO + 2HCI + NaCI + 2HO

An exact weight of sulfanilic acid, recrystallized and dried to constant weight, and a small amount of NaHCO3 (or NH3 solution) are dissolved in water, HCl, KBr are added, and ODM is titrated with stirring; NaNO2. The titration speed should be slow - at first 2 cm3 / min, at the end of the titration - 0.05 cm3 / min. The diazotization reaction is slow; potassium bromide is added to speed it up.

Indicators. In nitrimetry, an external indicator is used - starch iodide paper, which is prepared by impregnating filter paper with solutions of KI and starch, then the paper is dried. At the equivalence point, an excess of NaNO2 appears in the solution in an acidic environment, interacting with K1 of starch iodide paper. In this case, elemental iodine is released, coloring the starch blue:

Starch iodide paper is an external indicator; it cannot be lowered into the titrated solution. At the end of the titration, after adding each portion of the titrant, drops of the titrated mixture are also applied to starch iodide paper. At the equivalence point, the paper turns blue. In parallel, conduct a control experiment.

As internal indicators in nitritometry, redox indicators are used - tropeolin 00, a mixture of tropeolin 00 with methylene blue. When using tropeolin 00, the color of the solution at the equivalence point changes from red to yellow, a mixture of tropeolin 00 with methylene blue - from raspberry to blue. Analysis by the nitritometric method is carried out similarly to the determination of sulfanilic acid.

1. 5. 2. Applicability of redoximetry. In the process of redoximetry, the redox potential E of the titratable redox pair is constantly changing, which is determined by the Nernst equation by the ratio of the oxidized (approx.) and reduced (reductive) forms:

E= E° ± 0.059/n*lg, where E° is the standard redox potential of a given pair; n is the number of electrons involved in the reaction,

The value of the potential of the reacting redox pairs makes it possible to determine the possibility of using redox titration. If you set the completeness of binding of the analyte to 99.99%, then its residue at the time of equivalence in the system is 1*10−4 (0.01% of 100%) of the initial amount. Given the equivalence of the interaction, the equilibrium constant of the redox reaction should have a value of at least 1 * 108, and the EMF of the reaction is determined by the values: ok1 + vos2 ok2 + vos1; K =

at n=1 EMF= at n=2

1. 5. 3. Titration curves. Using the Nernst equation, one can calculate the potential value for any moment of titration and construct a titration curve that will express the dependence of the potential of the titrated redox pair E on the volume of the added titrant. An example of a titration curve for a solution of FeSO4 with a solution of KMnO4 in an acidic medium is shown in fig. 20. On the titration curve, there is a jump in the potential of the Fe3 + / Fe2 + system in the zone of the equivalence point. At the equivalence point, the oxidation potential of the system is 1.38 V. Titration curves allow you to analyze the course of the titration, highlight the boundaries of the titration jump potential and select an indicator of redox titration.

1. 5. 4. Indicators. Several types of indicators are used in redox methods: a) indicators that change color with a certain change in the redox potential of the system - redox indicators; b) indicators that change their color either when an excess of the reagent appears in the solution, or when the analytes disappear - specific indicators.

Redox indicators. They exist in two forms - oxidized and reduced, and the color of one form differs from the other. The transition of the indicator from one form to another and the change in its color occurs at a certain color transition potential of the indicator, which, in order to reduce errors, must be within the titration jump.

Diphenylamine. Used in titration with oxidizing agents:

1. Restored form

2. Oxidized formula

Diphenylamine 1 is oxidized to diphenylbenzidine violet 2, which has a blue color. Color transition potential, E0 = 0.76 V. An indicator solution in concentrated H2SO4 is used.

Phenylanthranilic acid is similar in structure to diphenylamine. In the reduced form, the indicator is colorless, in the oxidized form it is red-violet. Color transition potential E°=1.08V. Applied as a solution in H2SO4.

specific indicators. Starch is an indicator for the presence of free iodine. In the presence of free iodine, starch turns blue at room temperature, the blue color of starch is due to the adsorption of iodine on its molecules.

Ammonium thiocyanate. It is used in the titration of Fe3+ salts. At the equivalence point, when it is used, the titrated solution turns from red to colorless.

2. PRACTICAL PART

I decided in the laboratory to conduct a quantitative analysis of a hydrogen peroxide solution using permanganometric titration, iodometric titration to determine the amount of iodine in iodine tincture, bromatometric titration to determine resorcinol, dichromate oxidizability to determine the amount of oxygen in tap water.

2. 1. Permanganometric titration

2. 1. 1. Preparation of 250 cm3 0.1 N. titrated solution of potassium permanganate. 0.8 g of KMnO4 is weighed on a technical balance, dissolved in hot water, the solution is cooled and brought to 250 cm3 with water in the cylinder. Leave the solution in a dark place for 7 days, then filter it through a glass filter or glass wool. Store the solution in a bottle with a ground stopper in a dark place. Standardization of the solution is carried out at least once a week.

Standardization 0.1 n. potassium permanganate solution. The KMnO4 solution is standardized with respect to oxalic acid or sodium oxalate. 0.63 g of oxalic acid is placed on a watch glass, the watch glass with a weighed portion is weighed, oxalic acid is poured into a funnel inserted into a 100 cm3 volumetric flask, and the watch glass is weighed again. Add 50 cm3 of distilled water to a volumetric flask, while washing off the powder from the funnel, dissolve oxalic acid and bring the volume of the solution to 100 cm3. Calculate the normality of the solution. In a conical flask of 150-200 cm3, 20 cm3 of the prepared oxalic acid solution is measured, 20 cm3 of 2 N. H2SO4 and titrate with KMnO4 solution until a pale pink color appears, which does not disappear within 30 s. An approximate and three exact titrations are carried out. Calculate the arithmetic mean, titer, normality and a correction of 0.1 n. KMnO4 solution.

2. 1. 2. Determination of the content of hydrogen peroxide in solution. For work, a 5% solution of hydrogen peroxide is used. Measure 10 cm3 of hydrogen peroxide solution and dilute in a volumetric flask to 100 cm3. Pour 10 cm3 of the resulting solution into a titration flask, add 5 cm3 of diluted H2SO4 and titrate with 0.1 N. KMnO4 solution until a pink color appears, which does not disappear within 30 s. At least three titrations are carried out, the arithmetic mean and the percentage of hydrogen peroxide are calculated. = 0.001701 g/cm3.

2. 2. Iodometric titration

2. 2. 1. Preparation of 0.1 N. titrated sodium thiosulfate solution and starch solution. 6.25 g of sodium thiosulfate are weighed on a technical scale and 250 cm3 of freshly prepared and cooled distilled water are dissolved in a flask. The solution is stored in a dark glass bottle, closed with a stopper, equipped with a tube with soda lime (CO2 access is harmful).

A starch solution is prepared by mixing 0.5 g of soluble starch with 5 cm3 of cold water and pouring the resulting suspension into 100 cm3 of boiling water while stirring. Boil the solution for 2 minutes, add 1 g of salicylic acid or ZnC12. The solution is cooled and stored in a bottle with a cork stopper. The prepared solution should give a bright blue color with 2-3 drops of 0.01 N. iodine solution. If a brown color is formed, the starch solution is unsuitable.

2. 2. 2. Standardization 0.1 n. sodium thiosulfate solution. The sodium thiosulfate solution is standardized with potassium dichromate. Potassium dichromate is preliminarily subjected to double recrystallization and dried at 150C to constant weight. Prepare 100 cm3 O.1 n. potassium dichromate solution. About 0.4903 g (accurately weighed) K2Cr2O7 is dissolved in 100 cm3 of water in a volumetric flask. Calculate the titer and correction of the solution. 20 cm3 of the prepared solution are placed in a 500 cm3 titration flask; H2SO4 solution, stand for 3 minutes and add 100 cm3 of water. The released iodine is titrated with sodium thiosulfate solution. The titration is first carried out without an indicator until the solution turns yellow, then 3 cm3 of starch solution is added and titrated until the blue color disappears. At least three titrations are carried out. =0.004904 g/cm3.

2. 2. 3. Preparation of 0.1 N. iodine solution. 4 g of CI are dissolved in 6 cm3 of water, 1.27 g of iodine is added to the resulting solution (the sample is taken on a technical scale). After dissolution of iodine, the volume of the solution is adjusted with water to 100 cm3 (in a measuring cylinder). The iodine solution is stored in a dark bottle with a ground stopper.

2. 2. 4. Standardization 0.1 n. iodine solution. 15 ml of iodine solution are placed in a conical flask and titrated with 0.1 N. Na2S2O3 solution until yellow. Then add 3 ml of starch solution and titrate with sodium thiosulfate solution until the blue color disappears. At least three titrations are carried out. =0.01269 g/cm3.

2. 2. 5. Determination of iodine content in iodine tincture. For analysis, a commercial 5%, tincture of iodine is used. 2 cm3 of tincture of iodine are placed in a 200 cm3 conical flask and titrated with 0.1 N. N2S2O3 solution until the solution turns yellow. Add 3 ml of starch solution and continue titration until the blue color disappears. At least three titrations are carried out. =0.01269 g/cm3.

2. 2. 6. Quantitative determination of sodium thiosulfate. Dissolve about 0.5 g of sodium thiosulfate (accurately weighed) in 25 cm3 of water, add 2 cm3 of starch solution and titrate with 0.1 N iodine solution until a blue color appears. Three titrations are carried out. =0.02482 g/cm3.

2. 3. Bromatometric titration

2. 3. 1. Preparation of O.1 n. potassium bromate solution. About 0.2783 g (accurately weighed) of potassium bromate, previously recrystallized and dried at 150°, is dissolved in distilled water and the volume of the solution is adjusted to 100 cm3 with water in a volumetric flask. Stopper the flask and stir the solution. The solution is stored in a bottle with a closed stopper. A solution of potassium bromate is prepared from a weighed sample of exact concentration and does not require standardization. Calculate (T, N, K of KBrO3 solution) = 27.836 g/mol.

2. 3. 2. Quantitative determination of resorcinol. About 0.1 g (accurately weighed) of resorcinol is dissolved in 20 cm3 of water and the volume of the solution is adjusted to 100 cm3. 20 cm3 of the solution are placed in a flask with a ground stopper with a capacity of 250 cm3, 30 cm3 of 0.1 N. potassium bromate solution, 5 cm3 10% KBr, 5 cm3 50% H23O4, mix, stopper the flask and leave for 15 minutes. Then add 10 cm3 of 10% CI, mix the solution and leave it again in a closed flask for 10 minutes in a dark place. Add 2 cm3 of chloroform to the mixture and titrate with 0.1 N. Na2S2O3 released iodine. =0.001835 g/cm3. Calculate the content of resorcinol in percent.

2. 4. Bichromate oxidizability.

Oxidability is a value that characterizes the amount of oxygen required for the oxidation of organic substances contained in 1 liter.

Let us select a portion of water in such a way that about 50% is spent on its oxidation. We dilute the selected water sample with distilled water up to 20 ml (for river waters, you can usually take 20 ml). Transfer to a 300 ml round bottom flask, add 10 ml of 0.1 N solution, and carefully, in small portions, thoroughly mixing the mixture after each addition, 30 ml cf. =1.84 g/cm. Then add 0.3. 0.4g. It is more convenient to dissolve in advance in conc. used for analysis to the desired 0.3. 0.4 g of the reagent were contained in 30 ml. After all, we add several glass capillaries (beads or pieces of pumice stone) to the flask, attach it to a reflux condenser, heat to a slight boil, and boil for 2 hours. Then we cool the mixture, wash the walls of the refrigerator with 25 ml of distilled water and transfer the contents of the flask to a 500 ml conical flask. We wash the first flask several times with distilled water, collecting the washings in the same conical flask so that the volume of the solution is about 350 ml. We introduce into the mixture 4.5 drops of a solution of feroin or 10.15 drops of a solution of N-phenylanthronilic acid and titrate the excess of potassium bichromate with Mohr's salt.

In parallel, we will conduct a blank experiment, for which we take 20 ml. dist. water and repeat all the operations described above.

COD \u003d ((A-B) * N * K * 8 * 1000) / V,

Where A is the volume of Mohr's solution used for titration in a blank experiment, ml. ;

B is the volume of the same solution used for sample titration, ml;

H is the normality of Mohr's solution;

K - correction factor for the type of Mohr's salt solution;

8 - oxygen equivalent;

V is the volume of analyzed water, ml.

If in water there are chlorides and only easily oxidized substances, then the definition is without.

From the results of COD, it is necessary to calculate the correction: 0.23 mg is consumed per 1 mg. (Content should be determined in advance).

If the water also contains hardly oxidizable substances, then we process: to 20 ml of the sample, add 1 g, 5 ml with p = 1.84 g / cm, 10 ml of a 0.1 N solution. Very carefully pour in 30 ml. , add 0.75 g and reflux for 2 hours and continue the analysis as described above.

Simultaneously, a blank experiment is carried out.

2. 5. Research results

one). Hydrogen peroxide was used for the analysis in the form of a 5% solution, which was provided by the pharmaceutical company Antey. The solution was prepared by OAO Samaramedprom. The composition is fully consistent with that indicated on the label. During storage, hydrogen peroxide is destroyed. It can be concluded that the solution was fresh, which corresponds to the expiration date on the label.

2). Iodine tincture 5% of CJSC "Yaroslavl Pharmaceutical Factory", purchased from the pharmacy company "Antey", also corresponds to the indicated composition.

3). A sample of resorcinol weighing 0.1 g corresponds to its composition by only 99.93%. Resorcinol was provided by the Children's City Hospital in Vologda.

four). Bichromatic COD analysis was carried out during November 2005 with tap water. Experiments were carried out every 3 days. The average amount of oxygen required for the oxidation of organic substances contained in 1 liter of water is 35.00 mg, i.e. COD = 35.00 mgO2/l. this suggests that there is an increased content of organic matter in the water. Based on my results, I made a proposal to the school administration to purchase tap water filters in the school cafeteria.

Conclusion

In this scientific work I was able to study the relevant literature and learn a lot about this topic. I myself was convinced of the complexity and depth of the material under study, in particular, that redox reactions have a huge practical value in industry and medicine, that the products obtained depend on the conditions for the occurrence of OVR, and when these conditions change, a person can receive certain substances.

With the help of redox reactions, a huge amount of quantitative and qualitative analysis is carried out. Redox reactions are the basis of almost all biochemical processes. In my work, I was able to touch on a part of the whole huge theory of OVR. I will continue my work. My next task is to try to use different types titration using OVR in relation to medicines and tap water.

As you know, when a particle (atom, ion, molecule) gives up an electron, it is oxidized and is a reducing agent, if it accepts, it is reduced and is an oxidizing agent. Oxidation and reduction reactions are inextricably linked, they are called redox (redox) reactions. Most of the o-in reactions used in analytical chemistry proceed in solutions.

O - in the reaction are very widely used in analytical chemistry.

1) To transfer ions and compounds from lower to higher oxidation states and vice versa to prevent their harmful effects on the analysis (for example, the oxidation of Fe 2+ to Fe 3+ ;

2) To detect ions that give characteristic reactions with oxidizing or reducing agents, for example Mn 2+ (colorless) ® MnO 2 (black-brown); Mn 2+ (colorless) ® MnO 4 - (raspberry).

3) For the separation of ions that are oxidized or reduced to form sparingly soluble ions.

use various methods:

¾ permanonatometry (based on the reaction Mn 2+ ® MnO 4 - , while changing color);

¾ iodometry (I - ®I 2), etc.

Redox reactions are characterized by the presence in solution of at least 2 oxidized and 2 reduced forms, i.e. 2 oxidizing-reducing pairs.

To determine the direction of the redox reaction, the concept of the redox potential is used, which is different for each redox pair.

It happens standard island potential and unconventional. O-B dependency the potential of a given pair on the concentration of an oxidizing agent and a reducing agent is expressed by means of the Nernst equation.

As is known, when a metal plate is immersed in a salt solution of this metal, a potential arises at the solution/electrode interface, which cannot be measured directly. It is evaluated relative to the zero point - a standard hydrogen electrode, which consists of an inert conductor - a platinum plate coated with black, immersed in an acid solution with a H + = 1, and passed through the solution under a pressure of 1 atm. gaseous hydrogen.

The reaction proceeds H + + e "½ H 2

(H + + e ® H ads

N ads + N ads ®½ H 2)

That. an equilibrium is established on platinum, which, depending on the conditions, can be shifted in any direction. The value of the potential of this electrode is conventionally taken as 0.

For equilibrium conditions, the Nernst equation is valid:

The standard electrode potential of the hydrogen electrode, i.e. potential of a hydrogen electrode immersed in a solution with an activity of ions, with respect to which it is reversible, equal to 1 under standard conditions, R is the universal gas constant, T is absolute temperature, F is the Faraday number, equal to 96500 C. You can write:

If an electrode made of an inert metal, such as platinum, is immersed in a solution in which a redox process occurs, then a potential jump occurs at the electrode/solution interface due to this process. Such systems (and their potentials) are called redox or redox systems (potentials).

where is the standard potential about-in electrode, and are the activities of the oxidized and reduced forms, respectively, n is the number of electrons involved in the redox process.

Let's take a tin electrode, it is a platinum plate immersed in an aqueous solution with a salt of Sn 2+ and Sn 4+, for example. Sn(NO 3) 2 and Sn(NO 3) 4:

Sn 4+ + 2e « Sn 2+ . This reaction is potential-determining on platinum and determines the redox potential of the medium.

or:

Let's take an electrode representing a platinum plate immersed in an aqueous solution with soluble salts Fe 2+ and Fe 3+:

Fe 2+ + e « Fe 3+ This reaction is potential-determining on platinum and determines the redox potential of the medium.

If you connect 2 different electrodes, we get an electrochemical (galvanic) cell, while the electrode with a more positive potential is called the anode, and the negative potential is called the cathode.

Potentials measured in pairs with a hydrogen electrode at a concentration (activity) of ions equal to 1 and a temperature of 25 0 C are called normal.

If we neglect the contact potential difference, then the EMF of the element (E) is equal to the difference in electrode potentials:

,

where are the potentials of the cathode (more positive) and anode (more negative) of the electrodes, respectively. At the cathode there is a reduction process, and at the anode - oxidation.

In addition to galvanic cells, redox reactions in solutions also proceed in the direction when the oxidizing agent is the oxidized form, the redox potential (ORP) of which is higher, and the reducing agent is the reduced form of the pair, the redox potential of which is lower. Therefore, when writing and selecting coefficients in analytical chemistry great value have half-reactions of OB reactions, their equations and Nernst dependence. Therefore, when recording and selecting coefficients in analytical chemistry, only the half-reaction method is used (the electron balance method is not used).

In this case, the law of conservation of the mass of matter and the principle of electrical neutrality are taken into account, i.e. the number of charges must be the same.

Consider the O-B reaction:

KMnO 4 + Na 2 SO 3 + ... ® Mn 2+ + Na 2 SO 4

I half-reaction:

II half-reaction:

According to the Nernst equation:

That. under ln is [H + ], changing it we change the ORP.

It should be borne in mind that after compiling the ion-molecular equation for each of the half-reactions, the number of atoms of all elements of the left and right parts is equalized:

If the original ion or molecule contains more oxygen atoms than the reaction product, then:

1. In an acidic environment, an excess of oxygen atoms is bound by H + ions into water molecules;

2. In a neutral and alkaline environment, an excess of oxygen atoms is bound by water molecules into OH - groups;

If the original ion or molecule contains fewer oxygen atoms than the reaction product, then:

1. The lack of oxygen atoms in acidic and neutral solutions is compensated by water molecules;

2. In alkaline solutions - due to OH ions -

According to the Nernst half-reaction equation, it is possible to change the ORP not only due to the concentration of the oxidized and reduced form, but also due to the pH value of the solution. That. the same oxidizing agent (MnO 4 -) at different pH of the solution gives different products.

a) in an acidic environment, it is oxidized to Mn 2+

MnO 4 - + 8H + + 5e ® Mn 2+ + 4H 2 O 2

SO 3 2– + 2OH - -2e ® SO 4 2– + H 2 O 5

2MnO 4 – + 16H++ 5SO 3 2– + 10OH-® 2Mn 2+ + 5SO 4 2– + 13H 2 O

2MnO 4 – + 6Н + + 5SO 3 2– + ® 2Mn 2+ + 5SO 4 2– + 3Н 2 О

2KMnO 4 + 5Na 2 SO 3 + 3H 2 SO 4 ® 2MnSO 4 + 5Na 2 SO 4 + K 2 SO 4 + 3H 2 O

b) in a neutral environment - up to MnO 2:

KMnO 4 + Na 2 SO 3 + H 2 O ® 2MnO 2 + 5Na 2 SO 4 + KOH

(end-and solid phases are not written).

Federal Agency for Education

____________________________________________________________

State educational institution higher professional education St. Petersburg State Technological Institute (Technical University)

______________________________________________________________

Department of Physical Chemistry

Ilyin A.A., Naraev V.N., Smirnova E.N., Fomicheva T.I.

Fundamentals of chemistry of redox processes

Methodical instructions for practical exercises

St. Petersburg

UDC 541(076.5)

Fundamentals of the chemistry of redox processes: Guidelines for practical exercises / A.A. Ilyin, V.N. Naraev, E.N. Smirnova, T.I. Fomicheva - St. With.

The guidelines briefly review the basic concepts of the theory of redox processes. A technique for predicting reaction products depending on the conditions of its implementation, drawing up equations of redox reactions and selecting stoichiometric coefficients is given. The redox transformations and properties of the most important oxidizing and reducing agents are described. The guide also includes sections

are recommended for students of other non-chemical specialties of universities for use in practical and laboratory classes, in preparation for exams in general and special chemistry, and also as additional material when studying the course "Corrosion of chemical equipment", since the basis of corrosion processes are redox reactions.

Fig.3, table. 2, bibliography. 8 titles

Reviewer:

A.N. Belyaev, Doctor of Chemical Sciences, Professor Department of Chemical Technology of Catalysts, St. Petersburg State Technical University (TU)

Approved at a meeting of the educational and methodological commission of the chemical department on September 11, 2006.

INTRODUCTION

Modern specialist in the field information technologies A person working at chemical industry enterprises must ensure the efficient operation of electrical equipment, communication networks, automated control systems and process control systems, for which he needs to have an understanding of such chemical processes as oxidation, reduction, corrosion, and electrolysis.

Redox processes underlie many electrochemical industries (electrolysis, pyroelectrometallurgy, metallothermy, hydrometallurgy, blast furnace process, etc.), the production of many valuable products, as well as the conversion of chemical energy into electrical energy in galvanic and fuel cells.

Redox reactions are important for the functioning and vital activity of biological systems (photosynthesis, respiration, digestion, fermentation and putrefaction), accompanying the cycle of substances in nature.

In analytical chemistry, methods based on redox reactions (potentiometry, redoximetry, coulometry, polarography, etc.) are widely used.

1. REDOX REACTIONS

Characteristic features of redox reactions

Chemical reactions can be conditionally divided into two significantly different types. The first includes reactions in which the oxidation states of the atoms of the chemical elements that make up the reactants do not change. These reactions are:

exchange reactions such as

BaCl2 + K2 SO4 = BaSO4 + 2 KCl;

compound reactions, for example,

CaO + 2 H2 O \u003d Ca (OH) 2;

decomposition reactions, for example,

CaCO3 = CaO + CO2.

Another type are chemical reactions in which

oxidation states of atoms change . For example, in the reaction

Zn + 2HCl = ZnCl2 + H2

zinc and hydrogen atoms change their oxidation states. Such reactions are called redox. The change in oxidation states occurs as a result of the transfer of electrons from one atom or ion to another.

Oxidation is a process in which a particle (atom, molecule, ion) donates one or more electrons. Substances whose molecules, atoms or ions are capable of donating electrons, increasing the oxidation state of the corresponding atoms, are called reducing agents. In the process of donating electrons the reducing agent is oxidized.

Recovery is a process in which a particle (atom, molecule, ion) attaches one or more electrons. Substances whose molecules, atoms or ions are capable of attaching electrons, lowering the oxidation state of the corresponding atoms, are called oxidizing agents. In the process of electron addition oxidizer is reduced.

As an example, consider the reaction of the formation of iron sulfide from simple substances (iron and sulfur):

During this reaction, the iron atom, losing two electrons, is oxidized, increasing the oxidation state from zero to plus two (oxidation process):

Fe - 2 ē → Fe.

At the same time, the sulfur atom, accepting two electrons, is reduced, lowering the oxidation state from zero to minus two (reduction process):

S + 2 ē → S.

The oxidation of iron occurred due to sulfur, to which, as a more electronegative element, its electrons passed. By accepting electrons and acting as an oxidizing agent, sulfur is reduced, and iron, a reducing agent, is oxidized.

Oxidation states of atoms in compounds

Oxidation state1 of a chemical element (hereinafter we will denote it by the Latin letter n) is the conditional charge of its atom in the compound, calculated on the assumption that all bonds in the oxidizing or reducing agent molecule are ionic. In other words, oxidation state - this is the electric charge that would arise on an atom if all the electron pairs with which it forms chemical bonds with other atoms in compounds were completely displaced to the atoms of the most electronegative element. This value is of a formal nature and, in most cases, is far from the true values ​​of the electric charges arising on atoms due to the displacement of electron clouds.

The oxidation state can take on negative, positive (integer and fractional), zero values ​​and is placed above the element symbol in the compound formula from above with a preceding plus or minus sign, for example:

KMnO4	H2SO4

The negative value of the oxidation state in the compound has the atoms of the most electronegative elements, towards which the binding electron clouds are shifted. Atoms that donate their electrons to other atoms acquire positive oxidation states.

The oxidation state of elements in monatomic ions numerically coincides with their charge. When compiling redox reactions, the charges of real-life ions are usually written with a sign after the number. For example,

Na 1+ Ca 2+ Fe 3+ NO 3 - SO 4 2- PO 4 3-

1 In the literature, it is also called the oxidation number - from English - oxidation number

The unit when recording the oxidation state of an element in a compound (the charge of an ion) is often omitted.

The oxidation state of an element in a compound should be distinguished from its valence, with which the former may not coincide. According to the theory of valence bonds, valency is quantitatively characterized by the number of covalent bonds that an atom forms with other atoms in a compound, i.e. it is expressed as a non-zero unsigned integer. So, for example, the oxidation states of carbon in the compounds CH4, CH3 Cl, CHCl3, CCl4 are, respectively, -4, -2, +2, +4. The valency of carbon, i.e. the number of covalent bonds formed by it, in all these compounds is 4.

To determine the degree of oxidation of atoms of elements in compounds, one should be guided by the following provisions:

1. Hydrogen in most compounds exhibits an oxidation state of +1. Exception: hydrides of alkali and alkaline earth metals

(NaH, KH, CaH2, etc.), in which the oxidation state of hydrogen is n = -1.

2. Oxygen in the vast majority of compounds has an oxidation state equal to n = -2.

Exception:

a) peroxides (H2 O2, Na2 O2, K2 O2, BaO2, etc.), in which the degree of oxidation of oxygen atoms is n = -1;

b) superoxides (KO2, RbO2, CsO2, etc.), in which the oxidation state -1 has a complex superoxide ion -1 and, therefore, formally, the oxidation state of each oxygen atom is n = -1/2;

c) ozonides (KO3, RbO3, CsO3, etc.), in which the ozonide ion -1 has a single negative charge and, therefore, formally, the oxidation state of each oxygen atom is n = -1/3;

d) oxygen fluorides OF2 and O2 F2, where the oxidation state of oxygen is n = +2 and n = +1, respectively.

3. The oxidation states of atoms of elements in simple substances (N2, Cl2, O2, Pb, Cu, etc.) are assumed to be zero.

4. Alkali metals (n = +1), metals of the main subgroup of the second group of the periodic system, zinc and cadmium (n = +2) show a constant degree of oxidation in compounds.

5. The highest positive oxidation state of the atoms of elements is determined by the group number in the Periodic system of elements of D.I. Mendeleev. The exceptions are the elements of the copper subgroup (Cu, Ag, Au), oxygen, fluorine, as well as the metals of the eighth group.

6. The algebraic sum of the oxidation states of all the atoms of the elements that make up the molecule of a chemical compound is always zero.

The algebraic sum of the oxidation states of all the atoms of the elements that make up a complex ion is equal to its charge.

Based on these rules, to determine the oxidation state of an element atom in a compound, for example, nitrogen in ammonium hydroxide NH4 OH, an equation of the form

NH4OH	5 (+1) + 1 (-2) + 1 n = 0

Solving the equation with one unknown value regarding the oxidation state of nitrogen in the compound, we get n = -3.

The main oxidation states of the elements of the main subgroups are shown in Table 1. It can be seen from it that the highest oxidation state of the elements of the main subgroups in compounds, as a rule, coincides with the group number of the Periodic Table.

The rules for using the Periodic Table of Elements, which allow determining stable oxidation states, are given in Table 1.

Table 1. Oxidation states in compounds depending on the position of the chemical element in the Periodic system

Group number in the Periodic system of elements N

					highest			possible
oxidation				group number
element in
connections					Intermediate			oxidation state
equal to number				element		in compounds is
				groups minus two 2 :
					Least			possible
				oxidation of an element in compounds is
				group number minus eight:

For example, for sulfur, which is in the sixth group, stable oxidation states will be +6, +4, -2. For elements of the main subgroup of group VII (except F), as shown in the table, stable oxidation states are: +7, +5, +3, +1, -1.

2 Nitrogen exhibits several intermediate oxidation states: +4, +3, +2, +1.

For transition elements must be remembered the following stable oxidation states: Cu +2, (+1); Ag+1; Au +3, +1; Zn+2; cd+2; Hg +2 (+1); Cr +6, +3; Mn +7, +6, +4, +2; Fe +3, +2; Co+2; Ni+2.

NB3! As you can see, in contrast to valency, the oxidation state

atoms of chemical elements in compounds can take integer and fractional, positive and negative values, and in simple substances it is taken equal to zero.

Types of redox reactions

Among the many redox reactions, three types are distinguished:

-- intermolecular redox;

-- intramolecular redox

-- disproportionation reactions according to oxidation states (also called dismutation reactions or self-oxidation reactions - self-recovery).

AT intermolecular reactions redox the degree of oxidation is changed by the atoms of the elements that make up various molecules, for example:

2 KClO3 = 2 KCl + 3 O2

Disproportionation reactions according to oxidation states

(dismutations) are reactions in which atoms of the same element that make up the same compound act as both a reducing agent and an oxidizing agent, for example:

NH4 NO3 = N2 + 2 H2 O

3NB! = Nota Bene! You should pay attention!

In the above reaction, the oxidation state of nitrogen atoms is aligned: in ammonium nitrate, nitrogen exists in two different oxidation states -3 and +3, and in molecular nitrogen formed as a result of the reaction, nitrogen atoms have a zero oxidation state.

Redox properties of elements depending on the structure of their atoms

The course of redox reactions implies the presence in the reaction system of atoms, molecules or ions that are mutually opposite in their ability to donate or gain electrons.

In periods with an increase in the charge of the nucleus of atoms of elements, the reducing properties of simple substances decrease, and the oxidizing ones increase, reaching a maximum for halogens.

In the main subgroups of the periodic system, with an increase in the charge of the nucleus, the reducing properties increase, and the oxidizing properties weaken.

Metal atoms have only reducing properties, since they have 1, 2 or 3 electrons at the external energy level. The ability of metal atoms to donate electrons and turn into cations is characterized by the ionization potential. The lower the value of the ionization energy, the stronger the reduction ability of the metal is manifested. The strongest reducing agents are alkali and alkaline earth metals.

According to the strength of the reducing properties, the metals are arranged in a row in descending order (see Fig. 1). The arrangement of chemical elements shown in Figure 2 according to their reducing ability is called near voltage or metal activity.

Li Cs K Ca Na Mg Al Zn Fe Co Ni Sn Pb H2 Cu Ag Hg Au

Strengthening the ability (property) of atoms to donate electrons (strengthening the reducing properties of atoms)

Li + Cs + K + Ca 2+ Na + Mg 2+ Al 3+ Zn 2+ Fe 2+ Co 2+ Ni 2+ Sn 2+ Pb 2+ H + Cu 2+ Ag + Hg 2+ Au 3+

Strengthening the property of ions to attach electrons (strengthening the oxidizing properties of ions)

Fig.1 Electrochemical series of voltages of metals (in an acidic environment).

In a similar row, you can arrange the most common oxidizing agents in descending order (from left to right) of their oxidizing ability (see Fig. 2).

F2 > MnO 4 - > PbO2 > HClO > ClO4 - > BrO4 - > Cr 2 O 7 2- > Cl 2 > MnO2 >

> O2 > Br2 > NO 2 - > NO3 - > Fe3+ > I2 > O 4 2-

Fig.2 Oxidizing properties of the most important oxidizing agents in descending order of the strength of the oxidizing ability (in an acidic environment).

As can be seen from Fig. 2, fluorine is the strongest oxidizing agent. The oxidizing power of halogens in the group from top to bottom weakens. At the same time, the reducing ability of halide anions from F- to I- increases.

The redox properties of complex substances depend on the degree of oxidation of the atom of the element that changes it in the reaction.

Oxidizing properties are characteristic of those complex substances in which the atoms of elements capable of changing the degree of oxidation are in the highest degrees oxidation. For example, active oxidizing agents (and only oxidizing agents!) are potassium permanganate KMnO4, potassium dichromate K2 Cr2 O7, sulfuric acid H2 SO4 (concentrated!), since the atoms of manganese, chromium and sulfur are in the highest oxidation states possible for them, and, therefore , they can only accept electrons.

Substances containing atoms in the lowest oxidation states can only act as reducing agents.

Redox duality is manifested by substances in which the degree of oxidation of the atoms that change it has an intermediate value between the maximum and minimum possible values, i.e. in relation to some reagents, such substances can act as an oxidizing agent, and in relation to others - a reducing agent.

Redox transformations of the most important oxidizing and reducing agents

Starting to compile the equations of redox reactions, it is necessary to find out which of the reacting substances will act as an oxidizing agent and reducing agent, as well as possible reaction products, depending on the acidity of the medium.

Redox transformations of the most commonly used oxidizing and reducing agents are given in the Appendix. From the data given in the Appendix, it can be seen that the environment and conditions in which

1. Essence of oxidation-reduction reactions. Standard and real redox potential. Nernst equation

2. Basic provisions of the theory of redox reactions

3. Disproportionation reactions

Fig. 4. Dependence of the values ​​of the redox potentials of the systems on various factors (EMF of the reaction, pH)

5. Application of redox reactions in analytical chemistry

6. Cations of the V analytical group. Reactions of cations of the V analytical group.

KEYWORDS AND TERMS

Redox reactions

Oxidation process

Recovery process

Oxidizer

Reducing agent

Oxidation reactions intramolecular, intermolecular, disproportionation

Galvanic cell Platinum black

Standard redox potentials Real redox potentials Nernst equation

Catalysts

Inhibitors

Autocatalytic reactions Conjugated reactions Induced reactions

V analytical group of cations

Group reagent for V analytical group of cations

Reactions of iron (II) ions

Reactions of iron (III) ions

Reactions of manganese (II) ions

Reactions of bismuth ions

Reactions of antimony (III, V) ions

Systematic course of the analysis of cations of the V analytical group

Redox reactions in analytical chemistry

Redox reactions are reactions that are accompanied by the transfer of electrons from one particle (atoms, molecules and ions) to others, which leads to a change in the oxidation states of elements.

Oxidation is the process by which recoil occurs

No° - d - H2O2 - 2d - 2H+ + O2 BO32- + 2OH- - 2d - BO42- + H2O

Restoration is the process by which attachment occurs

electrons an atom, molecule or ion:

B° + 2d - B2-H2O2 + 2H+ + 2d - 2H2O SG2O72- + 14H+ + 6d - 2Cr3+ + 7H2O MnO4- + 8H+ + 5d - Mn2+ + 4H2O

The most commonly used reducing agents in analytical chemistry are: H2O2, SnCl2, H2B, H2BO3, Na2S2O3; as oxidizing agents -С12; Br2; H2O2; K2Cr2O7; KMpo4; NSh3 and others.

Two substances always participate in oxidation-reduction reactions, one of which is a reducing agent and the other an oxidizing agent. In this case, the number of electrons accepted by the oxidizing agent must always be equal to the number of electrons donated by the reducing agent.

Types of redox reactions

The main types of redox reactions include the following:

1 - intermolecular redox reactions (elements that change the degree of oxidation are part of different molecules):

PbS + 4H2O2 ^ ^PbBO4 + 4H2O

reducing agent oxidizing agent

NaC1 + NaC1O + NaC04 ^ C12 + Na2BO4 + H2O;

2 - intramolecular redox reactions (elements that change the degree of oxidation are part of one molecule):

NH4NO3^N2O + 2H2O

(NnH4)2 C+2O7 - N2° + Cr2O3 + 4H2O;

3 - disproportionation reactions (self-oxidation-self-recovery) (atoms of the same element change their oxidation state in different ways):

C°l2 + 2NaOH - NaCl + NaClO + H2O 2H2O2 - 2H2O + O2

AT general view The redox process is expressed as follows:

where: Red is a reducing agent-particle donating electrons;

Ox is an oxidizing particle that is capable of accepting electrons.

Thus, we are dealing with a conjugated redox or redox couple.

The following can be cited as the simplest redox pairs: Fe3+/Fe2+, Ce4+/Ce3+, etc. Both forms of these pairs are monatomic ions in which there is a direct transition of electrons from one particle to another.

More often, more complex processes take place in which electrons and atoms pass simultaneously, for example:

Mn2+ + 2H2O - 2e -- MnO2^ + 4H+.

A redox reaction can occur only in the direct interaction of an oxidizing agent (oxidized form) of one conjugated redox pair with reducing agents (reduced form) of another conjugated redox pair. In this case, the general redox reaction consists of two private ones:

aoh! + ne aRedl

BRed2-ne bOx2

aOxi + bRed2 ^ aRedl + bOx2

Thus, each redox reaction can be represented as the sum of two half-reactions. One of them reflects the transformation of the oxidizing agent, and the other - the reducing agent.

Redox potential

The division of redox reactions into half-reactions is not only a formal technique that facilitates understanding of the process of electron transfer, but has a well-defined physical meaning.

If the components of each half-reaction are placed in different vessels, the solutions of which are connected by a salt bridge (a tube filled with agar-agar gel impregnated with a saturated KCl solution), inert electrodes (platinum wire) are lowered into each of the vessels and connected to a galvanometer (Fig.), the latter will show the presence of electric current. Such a device is called a galvanic cell, and each of the vessels with a solution and a platinum wire is called a half-cell.

Fe2+ ​​-- Fe3+ + e

Ce4+ + e - Ce3+

Rice. Scheme of a galvanic cell: 1 - platinum electrodes; 2 - electrolytic key (tube filled with KC1 solution); 3 - potentiometer; 4 - glass.

When a reaction takes place in a galvanic cell, chemical energy is converted into electrical energy. The electromotive force (EMF) of a galvanic cell is the potential difference between two electrodes (EMF = E1-E2) and characterizes the ability of the reducing agent electrons to pass to the oxidizing agent.

The potential of an individual redox pair (in our example, Fe/Fe and Ce4+/Ce3+) cannot be measured. In this regard, this problem is solved by combining a half-cell, the potential of which must be determined, with a half-cell, the potential of which is taken as a standard. A standard hydrogen electrode is used as such a half-element (electrode). This electrode is a platinum plate coated with a layer of platinum black, which is immersed in an acid solution with a hydrogen ion activity equal to one. The electrode is washed with hydrogen at a pressure of 1.013-10 Pa (1 atm). In this case, the layer of platinum black is saturated with hydrogen, which behaves like a hydrogen electrode, which is in solution in equilibrium with hydrogen ions:

H2 (g) \u003d 2H + + 2e.

The potential of this electrode is assumed to be zero at all temperatures.

Given that the EMF of the galvanic cell is equal to the potential difference of the electrodes (E1 - E2) and E2 = 0 (standard hydrogen electrode): for the specified galvanic cell we will have EMF = E1. Thus, the potential of this electrode is equal to the EMF of the galvanic cell, which consists of this electrode and a standard hydrogen electrode.

In the case when all components of the half-reaction are in the standard state (a hypothetical one-molar solution with activity coefficients of the components equal to unity, at normal atmospheric pressure and temperature).

The potential of such an electrode is called the standard electrode potential and denoted by E0.

The relationship between the real equilibrium redox potential of the electrode (electrode potential under non-standard conditions - E) and the standard electrode potential (E0) is conveyed by the Nernst equation:

E \u003d E0 + -1n oh

where: R is the universal gas constant equal to 8.312 J/(mol K); T is the absolute temperature; F - Faraday's constant, equal to 96,500 C;

n is the number of electrons involved in the electrode process;

aox, ared are the activities of the oxidized and reduced forms of the substance, respectively, raised to powers equal to the corresponding stoichiometric coefficients. If we take into account that activity can be represented as the product of concentration and activity coefficient, then the Nernst equation can be written in the following form:

E = E0 + ^ log Y°x

nF "Y red"

where: y0x - activity coefficient of the oxidized form;

Yred - activity coefficient of the restored form. When other ions, in addition to the oxidized and reduced forms, are included in the half-reaction equation, their concentrations (activities) must be included in the Nernst equation.

So the potential of the redox pair EMn04-/ Mn2+ , for which the half-reaction equation has the form:

Mn04- + 8H+ + 5e ^ Mn2+ + 4H20,

calculated according to the equation:

*Mn°47 Mn2+ - E Mn°47 Mn2+ + ln "

If we substitute the numerical values ​​of the constants into the Nernst equation and go to decimal logarithms, then for a temperature of 25 ° C it takes the form:

E \u003d E "+ 0.059 18 "Us

0.059 lg |0x] + 0.059 lg

the value

in the last equation is called

formal potential and denoted as E0(1).

If a formal potential is introduced into the Nernst equation, then it takes the form:

E0(D + 0059 log |°x]

It follows from the last equation that E - E0(1) when [°x] - 1.0.

Equality E 0(1)

testifies

the formal potential depends on the ionic strength of the solution. If the ionic strength of the solution is neglected, then we obtain the equality E0(1) = E0. For many analytical calculations, the accuracy of this approximation is quite sufficient.

The standard potential of the redox system is used as an objective characteristic of the redox properties of compounds. The greater the positive value of the standard potential, the stronger the oxidizing agent. At the same time, the reduced forms of strong oxidizing agents have weakly expressed reducing properties, and, conversely, the oxidized forms of strong reducing agents have weak oxidizing properties. Thus, redox reactions proceed in the direction of the formation of weaker oxidizing agents and reducing agents from stronger ones.

Comparison of standard potential values ​​can be used to predict the direction of a redox reaction.

However, it should be taken into account that the standard potentials may differ significantly from the real ones and the direction of the reaction may change.

The value of the real potential is influenced by such factors as pH of the medium, concentrations of reagents, complexation, precipitation, etc.

It should be noted that the pH of the medium affects the real potential not only in cases where the concentrations of H+ and OH- ions are included in the Nernst equation, but sometimes also in cases where they are absent in this equation. This may be due to a change in the form of existence of ions in solution (influence on hydrolytic and other equilibrium processes).

As noted on page 96, the potential of the E°mpO4tmp2+ redox pair depends on the concentration of hydrogen ions (the indicated value can vary from 1.51 V to 1.9 V), and therefore this property is used for the fractionated oxidation of halide anions to free halogens. So at pH 5 to 6 permanganate oxidizes only iodides (E°y 21- = 0.53 V), at pH 3 bromides are oxidized (E°Br2 / 2Br- = 1.06 V) and only at a much higher acidity are oxidized chlorides (E ° su 2sg \u003d 1.395 V).

A change in the pH value can affect not only the value of the redox potential, but, sometimes, the direction of the reaction. For example, the reaction:

LvO43" + 2H+ + 21" - LbOz3" + 12 + H2O,

in an acidic environment, it flows from left to right, and in an alkaline one (pH 9, created with NaHCO3) - from right to left.

There is the following rule for creating the reaction medium necessary for the optimal course of the process: if H + or OH - ions accumulate as a result of the redox reaction, then it is necessary to create a medium that has the opposite properties (alkaline or acidic, respectively).

In addition, for the reaction, it is necessary to take components (oxidizing agent and reducing agent) that react in the same medium. Otherwise, the process may stall.

As an example, which illustrates the effect of a poorly soluble compound on the potential and direction of the reaction, one can cite the reaction of the interaction of Cu2+ and I-.

The standard potential of the Cu/Cu redox pair is 0.159 V, and for the 12/21- pair it is 0.536 V. The given data indicate that Cu2+ ions cannot oxidize I- ions. However, in the course of this reaction, poorly soluble copper single iodine Cu1 is formed (PRCu1 = 1.1T0-

). The formation of this precipitate sharply reduces the concentration of Cu ions in the solution. In this case, the activity of Cu2+ ions in solution can be determined from the equation:

Using the Nernst equation, one can show that

E ° cu- / oui \u003d E0cu2 + / cu + - 0.059lgnP CuI \u003d 0.865 V.

Thus, the Cu /CuI redox pair has a potential that exceeds the potential of the I2/I- pair, as a result of which this transformation becomes possible.

The equilibrium constant of the redox reaction

In many cases, it is necessary to know not only the direction of the redox reaction, but also how complete it is. So in quantitative analysis, those reactions are used that practically proceed by 100% (or approach it). The extent of the reaction is determined by the equilibrium constant.

If the equation of the redox reaction is presented in general form by the scheme:

aOx1+ pRed2 ^ yRed1+ 5Ox2 ,

then its equilibrium constant (K) will look like:

K = a a ■ I'm in-.

Using the Nernst equation for the redox pairs involved in the reaction, it can be shown that at 25°C:

1 m^ Ox, /Re d, Ox2/Red2/

or in general:

T, (E ox - E Re d) ^n

where: ERed - standard potentials of pairs acting in a given re-

shares as an oxidizing agent and a reducing agent, respectively; n is the number of electrons involved in the process. It follows from the last equation that the greater the potential difference (EOx-ERed), the greater the equilibrium constant and the more complete the reaction will proceed from left to right. However, it should be remembered that great importance the equilibrium constant does not indicate a high rate of the process.

The reaction rate is affected by such factors as the reaction mechanism, the concentration of reactants, the temperature of the solution, the presence of catalysts or inhibitors.

The reaction rate increases with an increase in the concentration of the reactants, as well as with an increase in temperature. Usually, an increase in temperature by 10° leads to an acceleration of the reaction by a factor of 2-4. So, for example, when potassium permanganate interacts with a solution of oxalic acid, the reaction proceeds slowly and the solution is heated to speed it up. In addition, this reaction is autocatalytic (the catalyst is one of the reaction products - Mn ions).

In some cases, so-called conjugated or induced redox reactions take place.

For example, during the oxidation of Fe2+ ions with potassium permanganate in a solution acidified hydrochloric acid, part of potassium permanganate is spent on the oxidation of chloride:

2Mn04- + 10C1- + 16H+ - 2Mn2+ + 5C12 + 8H2O.

In the absence of Fe2+ ions, this reaction does not occur, although the difference in standard potentials allows it to proceed. To prevent a side reaction, this process is carried out in solutions acidified with sulfuric acid.

Application of redox reactions in analytical chemistry

Redox reactions are widely used in qualitative and quantitative analysis.

In qualitative analysis, redox reactions are used to:

Transferring compounds from lower oxidation states to higher ones and vice versa;

Transfer of poorly soluble compounds into solution;

Ion detection;

Ion removal.

Thus, the reactions of oxidation with hydrogen peroxide in an alkaline medium are used in the analysis of cations of the IV analytical group for the conversion of compounds Sn(I), Av(III), Cr(III) into hydroxo- and oxoanions of these elements in higher oxidation states.

For example, anions [Cr (0H) 6] - are oxidized into chromate ions:

2[Cr(0H)b]3- + 3H20^2Cr042- + 8H2O + 2OH-. Anions [Sp(0H)6]4- - to hydroxoanions [Sp(0H)6]2-: [Sp(0H)6]4- + H2O2- [3p(0H)6]2- + 2OH-.

Arsenite ions Ab03s- - to arsenate ions Ab04s-:

Ab03- + H202^=* Ab043- + H20.

At the same time, this reagent simultaneously oxidizes some ions of the V and VI analytical groups: Mn2+ - Mn02-pH20, Co2+-Co3+.

The oxidizing properties of hydrogen peroxide in an acidic environment are used to detect chromate and dichromate ions by the formation of perchromic acid (blue solution):

Cr2072- + 4H202 + 2H+ -2H2Cr06 + 3H20.

Hydrogen peroxide in a nitric acid medium is used as a reducing agent for transferring MnO2-pH20 and Co(OH)3 precipitates into solution in the form of Mn2+ and Co2+ cations, respectively:

MnO2-pH20 + H2O2 + 2H+^=* Mn2+ + |O2 + (n+2)H20, 2Co(OH)3 + H2O2 + 4H+^ 2Co2+ + |O2 + 6H20.

To detect the ions Av(III) and Ab(Z) the reaction of their reduction with metallic zinc to the gaseous product AbH3 is used.

In the analysis of group II cations, the reaction of disproportionation of mercury compounds is used:

HB2S12 + M1z- [NvMDO! + Hv |.

To detect Bi cations, their reducing properties with respect to divalent mercury compounds are used:

Bp2+ + + 4 C1 - + [BpC16]2-.

Ions Sb^) are detected by their reduction to Sb0 by the action of zinc metal on a nickel plate on their anions (formation of a galvanic pair):

2[SbC16]- + 57p- 57p2+ + 2Sb| +12С1-.

N cations are detected by reducing them in an alkaline medium to metallic bismuth with hexahydroxostannite (II) ions:

2N(0H)s| + 3[Bp(0H)6]4- - 2BC + 32- +60H-.

The detection of Mn ions is based on their oxidation to red-colored MnO4 ions with ammonium persulfate:

2Mp2+ +5B2082- + 8N2<0^ 2Мп04- + 16Н+ + 10Б042-.

Oxidizing agents such as HNO3, chlorine and bromine water are often used in qualitative analysis.

Thus, the mixture of precipitates of Cu2S and HgS sulfides formed in the systematic course of the analysis is separated by treating it with dilute nitric acid when heated:

3Cu2S + 4NO3" +16H+ - 6Cu2+ + 3S| + 4N0| + 8H20.

In this case, under these conditions, the HgS precipitate, in contrast to Cu2S, does not dissolve.

The HgS precipitate dissolves in bromine water in the presence of hydrochloric acid or in a mixture of concentrated HNO3 and concentrated HCl:

HgSj + Bg2 + 2Cl- - + 2Br- + Sj, 2HgS| + 2HNO3 + 6HC1 - 3 + 3S|+2N0 + 4H2O.

The black precipitate of metallic antimony obtained upon its detection (see above) is oxidized with nitric acid to form a white precipitate of metaantimony acid:

3Sb| + 5MV + 5H+ - 3HSb03| + 5N0|+ H20.

Redox reactions are widely used not only in the qualitative analysis of cations, but also in the analysis of anions.

So, during the analysis for anions, a test for oxidizing anions (Cr2072-, As043-, NO3-) is performed by the action of a KI solution in an acidic medium in the presence of chloroform. In this case, free iodine is formed, which stains the chloroform layer in a red-violet color.

In addition, a test is carried out for reducing anions (C204 S203 S2-, S032-, As033-, I-, N02-), based on the discoloration of the iodine solution in a slightly acidic medium (with the exception of As033-, which are found in a slightly alkaline medium).

A sample with concentrated H2S04 (Cl-, Br-, I-, Cr042-, N03-, S032-, S2032Ti etc.) is based on redox reactions, while gaseous products (Cl2, I2, C02, etc.) are released.

To detect reducing anions such as S03 "and C204 N02-, reactions are used, as a result of which KMn04 solutions become colorless:

5S02 + 2H20 +2Mn04- ^ 5S022- + 4H+ + 2Mn2+ 5C2042- +2Mn04- + 16H+ ^ 10C02t + 2Mn2+ + 8H20 5N02-+2Mn04- + 6H+^ 2Mn2+ + 5N03- + 3H20.

To remove nitrite ions when nitrate ions are detected, a reaction with crystalline ammonium chloride is used:

N02-+ Mi/^ N2| + 2H20.

Analytical group V cations Te2+, Te3+, M^+, Mn2+, 8b (III), 8b (V), B13+

general characteristics

Analytical group V includes cations of D-elements - Mg2+, p-elements - Bb (III), Bb (V), N3+ and e-elements - Fe3+, Fe2+, Mn2+. Due to the strong polarizing effect of cations of the V analytical group, many of their compounds (hydroxides, sulfides, phosphates) do not dissolve in water. Chlorides, bromides, nitrites, nitrates, acetates, sulfates of cations of the V analytical group dissolve in water.

The group reagent for cations of the V analytical group is a concentrated solution of ammonia, which precipitates them in the form of hydroxides, insoluble in excess of the reagent.

The further course of the analysis of cations of the V analytical group is based on the different solubility of the hydroxides of these cations in concentrated solutions of ammonium salts, acids, as well as on the use of various redox reactions and precipitation reactions of these cations.

Be compounds are yellow-brown in color, and Be are light green; solutions of compounds of other cations are colorless.

Reactions of cations of the V analytical group

The action of solutions of sodium hydroxide or potassium hydroxide

With solutions of N0H or K0H, magnesium, manganese, bismuth and antimony cations form white amorphous precipitates of hydroxides, green - iron (II) hydroxides and red-brown - iron (III) hydroxides:

Mv2+ + 20H- - Mv(0H)2^, Mn2+ + 20H- ^ Mn(0H)2^, Fe2+ + 20H- - Fe(0H)2^, Fe3+ + 30H- - Fe(0H)3^, [SbC16 ]3- + 30H- - Sb(0H)3^ + 6C1-, [SbC16]- + 50H- - Sb(0H)5^ + 6C1-, Sb(0H)5^ - Hb03^ + 2H20, N3+ + 30H - ^ N(0H)3^.

All hydroxides of cations of the V analytical group are soluble in acids, for example:

Fe(0H)3^ + 3H+ - Fe3+ + 3H20.

Hydroxides 8b (III) and 8b (V) dissolve in an excess of alkalis due to their amphoteric properties:

8b (OI ^ + 3OI "- 3-, ^ + OI- - -.

Hydroxides of magnesium, manganese (II) and iron (II) are also soluble in a saturated solution of NaI4C1, for example:

Nv(OI)2^ + 2MI4+ - Nv2+ + 2MI3-I2O.

This property is used to separate magnesium hydroxide from other hydroxides of analytical group V cations in a systematic course of analysis.

The action of the ammonia solution

Under the action of an ammonia solution on solutions of cations of the V analytical group, precipitation of the corresponding hydroxides precipitates:

Nv2+ + 2MuU2O - Nv(OI)2^ + 2M1/, - + 5Mіs-U2O - WSO3| + 6C1- + 5M/4+ + 2^O, 3- + 3Ks-H2O - 8L(OH)3^ + 6C1- + 3M/4+, Mn2+ + 2M/3-H2O ^ Mn(OI)2^ + 2M/, Fe2+ + 2MuU2O ^ Fe (OI)2^ + 2MI+, Fe3+ + 3MI3-I2O ^ Fe(OI)3^ + 3M1L

Bismuth cations under the action of ammonia solution form a white precipitate of the basic salt, the composition of which varies depending on the concentration of the solution, temperature:

Ві3+ + 2Мі3-Н2О + С1- - Ві(ОI)2С1^ + 2МІ/, Ві(ОI)2С1^ - BiOC1^ + И2О.

Hydrolysis of 3b(III), 3b(V) and bismuth salts

Salts of bismuth, antimony (W^) are hydrolyzed with the formation of white precipitates of basic salts:

3- + H2O - 8CO1^ + 5C1- + 2H+, - + 2 H2O - 8LO2C1^ + 5C1- + 4H+, Bi3+ + H2O + W)3- ^ BiOZ)3^ + 2H+.

All precipitates are soluble in acids.

The reaction of magnesium ions

Action of sodium hydrogen phosphate solution Na2HPO4

Magnesium cations form a white crystalline precipitate with a solution of sodium hydrogen phosphate in the presence of an ammonia buffer solution:

Mg2+ + HPO42- + NH3-H2O ^ MgNHPO^ + H2O.

This reaction can be performed as microcrystalloscopic. MgNH4PO4 crystals formed during rapid crystallization have a characteristic shape.

The reaction of iron (II) ions

Action of a solution of potassium hexacyanoferrate (III) K3

Iron (II) cations form with a solution of potassium hexacyanoferrate (III) a blue precipitate ("turnbull blue"):

3Fe2+ + 23-^Fe32^.

The reaction is specific and makes it possible to detect Fe cations by the fractional method.

The precipitate is insoluble in acids. In alkalis, it decomposes: Fe32^ + 6OH- - 3Fe(OH)2^ + 23-.

Reactions of iron (III) ions

1. Action of solutions of potassium hexacyanoferrate (II) K4

Iron (III) cations form a dark blue precipitate ("Prussian blue") with a solution of potassium hexacyanoferrate (II):

4Fe3+ + 34- - Fe43^

The reaction of Fe with potassium hexacyanoferrate (III) is specific and makes it possible to open them by the fractional method.

The reaction must be carried out in an acidic medium at pH=3. However, with strong acidification or the addition of an excess of the reagent, the precipitate dissolves.

In alkalis, the precipitate decomposes:

Fe43^ + 12OH- - 4Fe(Offb^ + 34-.

2. Action of thiocyanate ions

Iron (III) cations form iron (III) complex compounds with thiocyanate ions, which color the solution red, for example:

Fe3+ + 3NCS- - .

With an excess of thiocyanate ions, complex ions of various composition are formed:

-; 2-; 3-.

It is necessary to carry out the reaction in an acidic medium at pH=2. Discovered

The reduction of Fe with thiocyanate ions is interfered with by anions (F, PO4 ", etc.), which form more stable complexes with Fe3 +, for example:

6F- -- 3- + 3NCS-.

Reactions of iron ions (III), (II)

The action of a solution of sulfosalicylic acid

Iron (III), (II) cations form complexes of different colors with sulfosalicylic acid depending on the pH of the solution. At pH = 1.8-2.5, a violet complex is formed:

At pH = 4-8, a red complex is formed:

At pH = 8-11, a yellow complex is formed:

Reactions occurring in the systematic course of the analysis upon the detection of manganese ions

Manganese (II) hydroxide is easily oxidized by hydrogen peroxide, and a dark brown precipitate H2MnO3 (MnO2-pH2O) is formed.

2e + Mn(OH)2^ + H2O ^ H2MnO3^ + 2H+ 1

2е + Н2О2 + 2Н+ - 2Н2О__

Mn(OH)2^ + H2O2 -> H2MnO3 + H2O^

Manganese (IV) is also reduced to manganese (II) in a sulfuric or nitric acid medium under the action of H2O2:

2e + H2MnO3^ + 4H+ ^ Mn2+ + 3H2O 1 2e + H2O2 - O2T + 2H+__

H2MnO3^ + H2O2 + 2H+ -> Mn2+ + O2T+ 3H2O

Strong oxidizing agents, for example, ammonium persulfate (M14) 282O8, oxidize Mn (II) to MnO4- ions, which color the solution red:

5e + Mn2+ + 4H2O ^ MnO4- + 8H+ 2

2Mn2+ + 5B2O82- + 8H2O -> 2MnO4- + 16H+ + 10BO42-

The reaction proceeds when heated and in the presence of silver salts (catalyst). This reaction is used to detect manganese (II) ions in the systematic course of the analysis.

Reactions of bismuth ions

The action of freshly prepared sodium hexahydroxostanite (P) Na44-.

Hexahydroxostanite (11) ions reduce Bi ions to black bismuth metal. Ions 4- are stable only in alkaline solutions. Bismuth cations under these conditions form a white precipitate of Bi(OH)3:

Sn2+ + 2OH- - Sn(OH)2^, Sn(OH)2^ + 4OH- - 4-, Bi3+ + 3OH- ^ Bi(OH)3^.

3e + Bi(OH)3^ ^Bil + 3OH- 2

2e + 4- - 2-_^ _ | 3

2Bi(OH)3^ + 34- - 2Bi^ + 32- + "6OH-

When performing the reaction for the detection of bismuth cations, an excess of concentrated alkali and heating should be avoided, since under these conditions a black precipitate Snl- may form due to the disproportionation reaction:

24- - 2- + Sn^ + 6OH-.

Reaction of antimony ions Action of metallic zinc

Metallic zinc on a nickel plate reduces antimony (III) and (V) ions to metallic antimony:

23- + 3Zn - 2SU + 3Zn2+ + 12C1-.

The nickel plate forms a galvanic pair with zinc, in which the positive electrode is Ni, and the negative electrode is Zn. Nickel receives electrons, which zinc donates to it, and transfers them to antimony ions, which are reduced to metal. A deposit of metallic antimony on a nickel plate does not dissolve in hydrochloric acid, but does dissolve in nitric acid:

5e + Sb^ + 3H2O ^ HSbO3^ + 5H+ 3

3e + NO3- + 4H+ - NOt + 2H2O_\_5_

3Sb^ + 5NO3- + 5H+ - 3HSbO3^ + 5NOt + H2O

Systematic course of the analysis of cations of the V analytical group

Iron (II) and iron (III) cations are detected by the fractional method in separate samples by the action of solutions of potassium hexacyanoferrate (III) and hexacyanoferrate (II), respectively:

3Fe2+ + 23- - Fe32^, 4Fe3+ + 34- ^ Fe43^.

Antimony (III) and antimony (V) cations interfere with the detection of all cations of the V analytical group. Antimony (III) is oxidized with a solution of HNO3 to antimony (V), which precipitates as HSbO3. The precipitate is separated by centrifugation and dissolved in concentrated hydrochloric acid. Antimony (V) in solution is detected by the action of zinc on a nickel plate.

All Group V cations remaining in the centrifuge are precipitated as the corresponding hydroxides by the action of concentrated ammonia solution.

To separate magnesium hydroxides, a saturated solution of ammonium chloride and a 3% solution of H2O2 are added to the precipitate. In this case, the Mg(OH)2 precipitate dissolves, and the Mn2+ cations are oxidized to H2MnO3 (MnO2-nH2O). The precipitate consisting of Fe(OH)3, BiONO3, H2MnO3 (MnO2-nH2O) is separated by centrifugation. Magnesium cations are detected in the centrifuge by the action of a solution of sodium hydrogen phosphate in the presence of an ammonia buffer solution.

Under the action of nitric acid on the precipitate, the hydroxides of iron (III) and bismuth dissolve, and H2MnO3 (MnO2-nH2O) ^ remains in the precipitate, which is separated by centrifugation.

Bismuth cations are detected in the centrifuge by the action of a freshly prepared Na4 solution.

The precipitate H2MnO3 (MnO2-nH2O) is dissolved in nitric acid in

the presence of H2O2, while MnO2-nH2O is reduced to Mn. They are detected by the action of ammonium persulfate solution.

The systematic course of the analysis of a mixture of cations of the V analytical group is performed in accordance with the scheme.

SCHEME OF THE SYSTEMATIC PROCESS OF THE ANALYSIS OF CATIONS OF THE V ANALYTICAL GROUP

Fe2+, Fe3+, Mn2+, M^+, Bi3+, 8b (III), 8b (V)

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Lecture plan: The use of OVR in analytical chemistry. Types of OVR. Quantitative description of OVR. OVR equilibrium constant. Stability of aqueous solutions of oxidizing and reducing agents.

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The use of OVR in analytical chemistry During sample preparation for transferring a sample into a solution. To separate a mixture of ions. For masking. For conducting cation and anion detection reactions in qualitative chemical analysis. in titrimetric analysis. In electrochemical methods of analysis.

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For example, during hypoxia (a state of oxygen starvation), the transport of H + and e - in the respiratory chain slows down and the reduced forms of compounds accumulate. This shift is accompanied by a decrease in the OB potential (ORP) of the tissue, and as the ischemia deepens (local anemia, insufficient blood content in the organ or tissue), the ORP decreases. This is due both to the inhibition of oxidation processes due to a lack of oxygen and a violation of the catalytic ability of redox enzymes, and to the activation of recovery processes during glycolysis.

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Types of OVR 1. Intermolecular - the oxidation states (CO) of the atoms of the elements that make up different substances change:

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2. Intramolecular - oxidizing agent and reducing agent - atoms of one molecule:

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3. Self-oxidation - self-healing (disproportionation) - the same element increases and lowers S.O. Cl2 - is an oxidizing agent and a reducing agent.

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Quantitative description of OVR For example, the stronger the base, the greater its affinity for the proton. Also, a strong oxidizing agent has a high electron affinity. For example, in acid-base reactions, a solvent (water) is involved, giving and receiving a proton, and in OVR, water can also lose or gain an electron. For example, acid-base reactions require both an acid and a base, while OVR requires both an oxidizing agent and a reducing agent.

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Considering the OB pair as a whole, we can write a schematic equation for the reaction: Ox + nē = Red

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At a temperature of 298 K, the Nernst equation takes the form:

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It is difficult to directly measure the electrode potential, so all electrode potentials are compared with any one ("reference electrode"). The so-called hydrogen electrode is usually used as such an electrode.

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In the Nernst equation, instead of the activities of ions, their concentrations can be used, but then it is necessary to know the coefficients of ion activities:

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The strength of the oxidizing agent and reducing agent can be influenced by: pH value, precipitation reactions, complexation reactions. Then the properties of the redox pair will be described by the real potential.

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To calculate the real potential of half-reactions obtained by a combination of OVR and precipitation reactions, the following formulas are used:  if the oxidized form is a poorly soluble compound:

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 if the reduced form is a poorly soluble compound:

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Combination of OVR and complexation reactions

 if the oxidized form is bound into a complex:

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 if the restored form is connected in a complex:

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 if both forms are connected in a complex:

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Combination of OVR and protonation reactions

if the oxidized form is protonated:

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if the reduced form is protonated:

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if both forms are protonated:

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if the reaction proceeds according to the following equation: Ox + mH+ + nē = Red + H2O then