Redox titration. Essence and classification of redox titration methods

oxidation states

For example:

Methods for establishing T.E.

To determine the equivalence point during redox titration, use:

a) non-indicator methods. In the case where the solution of the titrated substance or titrant is colored, TE can be determined by the disappearance or appearance of this color, respectively;

b) specific indicators - changing color when the titrant appears or the substance being determined disappears. For example, for the J 2 /2J - system, the specific indicator is starch, which colors solutions containing J 2 blue, and for Fe 3+ ions the specific indicator is SCN - ions (thiocyanate ions), the resulting complex is colored blood-red ;

c) RH (redox) indicators – changing color when the RH potential of the system changes. Single-color indicators are diphenylamine, two-color indicators are ferroin.

Redox indicators exist in two forms - oxidized (Ind ok) and reduced (Ind rec), and the color of one form is different from the other. The transition of an indicator from one form to another and a change in its color occurs at a certain transition potential, which is observed when the concentrations of the oxidized and reduced forms of the indicator are equal and according to the Nernst-Peters equation:

The transition interval of redox indicators is very short, unlike acid-base indicators.

RH titration curves

RH titration curves depict the change in the RH potential of the system as the titrant solution is added.

Reductometry, when a solution of an oxidizing agent is titrated with a standard solution of a reducing agent

In reductometry, titration curves are calculated:

Oxidimetry, when a reducing agent solution is titrated with a standard oxidizing agent solution

In oxidimetry, titration curves are calculated:

Example. Let's calculate the titration curve of a 100 cm 3 solution of FeSO 4 with a molar concentration equivalent to 0.1 mol/dm 3 with a KMnO 4 solution of the same concentration.

Reaction equation:

The equilibrium constant of this reaction is

A large numerical value of the equilibrium constant indicates that the equilibrium of the reaction is almost entirely shifted to the right. After adding the first drops of titrant, two OM pairs are formed in the solution: , the potential of each of which can be calculated using the Nernst equation:

In this case, the reducing agent solution is titrated with an oxidizing agent solution, i.e. Titration refers to the oxidimetry method; the titration curve is calculated according to the appropriate scheme.

3) After T.E.

Calculation data for constructing a titration curve

No.	τ	Calculation formula	E, B
	0,10		0,71
0,50	0,77
0,90	0,83
0,99	0,89
0,999	0,95
			1,39
	1,001		1,47
1,01	1,49
1,10	1,50
1,50	1,505

Using the table data, we construct a titration curve:

For titration error ±0.1% titration jump

∆E = E τ =1.001 - E τ =0.999 = 1.47 – 0.95 = 0.52.

For titration error ± 1.0% titration jump

∆E = E τ =1.01 - E τ =0.99 = 1.49 – 0.89 = 0.60.

In the region of TE, when moving from a solution undertitrated by 0.1% to a solution overtitrated by 0.1%, the potential changes by more than 0.5 V. The potential jump makes it possible to use directly potentiometric measurements or RH indicators, the color of which changes with change in potential. In addition, in this case, a colored solution is used as a titrant, therefore T.E. can be determined by the appearance of a faint pink color from one excess drop of potassium permanganate.

PERMANGANOMETRY

The method is based on the oxidation of solutions of reducing agents with potassium permanganate KMnO 4. The oxidation of reducing agents can be carried out in various environments, and manganese (VII) is reduced in an acidic environment to Mn 2+ ions, in a neutral environment to manganese (IV) and in an alkaline environment to manganese (VI). Typically, in the permanganatometry method, the reaction is carried out in an acidic environment. In this case, a half-reaction occurs

A titrated solution cannot be prepared using an exact weighing, because it contains . Therefore, first prepare a solution of approximately the required concentration, leave it in a dark bottle for 7-10 days, filter off the precipitate, and then set the exact concentration of the resulting solution. Standardization of the solution is carried out using a titrated solution of oxalic acid ( ) or sodium oxalate ().

The indicator is the permanganate itself, colored red-violet. The end of the reaction is easily determined by the change in color from one excess drop of permanganate. In an acidic environment, the titrated solution turns pink due to excess MnO 4 - ions. A big disadvantage of redox reactions is their low speed, which complicates the titration process. Heat is used to speed up slow reactions. As a rule, with every 10° increase in temperature, the reaction rate increases by 2-3 times. The oxidation reaction with oxalic acid permanganate is carried out at a temperature of 70-80 °C. Under these conditions, titration proceeds normally, since the reaction rate increases significantly.

If heating cannot be used (volatilization of one of the substances, decomposition, etc.), the concentrations of the reacting substances are increased to speed up the reaction. The reaction rate can be affected by the introduction of a catalyst into the solution.

The oxidation reaction of oxalic acid permanganate can be catalytically accelerated by the addition of MnSO 4, the role of which is as follows:

The resulting manganese dioxide oxidizes oxalic acid, reducing to manganese (III):

Thus, manganese (II) added to the solution is completely regenerated and is not consumed in the reaction, but greatly accelerates the reaction. In permanganatometry, one of the products of the oxalic acid oxidation reaction is Mn 2+ ions, which, as they form in solution, accelerate the reaction process. Such reactions are called autocatalytic. The first drops of permanganate during the titration of a hot acidified solution of oxalic acid become discolored slowly. As a small amount of Mn 2+ ions is formed, further discoloration of the permanganate occurs almost instantly, since the formed Mn 2+ ions play the role of a catalyst.

Redox titration

Redox processes include chemical processes that are accompanied by changes oxidation states atoms of substances participating in the reaction.

Substances whose atoms reduce their oxidation state during a reaction due to the addition of electrons are called oxidizing agents, i.e. they are electron acceptors. In this case, the oxidizing agents themselves are reduced. Reducing agents, being electron donors, are oxidized.

The product of reduction of an oxidizing agent is called the reduced form, and the product of oxidation of a reducing agent is its oxidized form. The oxidizing agent with its reduced form constitutes a half-pair of the redox system, and the other half-pair is the reducing agent with its oxidized form. Thus, a reducing agent with an oxidized form and an oxidizing agent with its reduced form constitute two semi-pairs (redox pairs) of the redox system.

All OM processes (redox reactions) can be divided into three types

a) intermolecular, when during the OB reaction the transfer of electrons occurs between particles of different substances. For example

In this reaction, the role of the oxidizing agent in the presence of H 3 O + is played by ions, and the ions act as a reducing agent

b) dismutation (disproportionation), during which the transfer of electrons occurs between particles of the same substance. As a result of disproportionation, the oxidation state of one part of the atoms decreases at the expense of another part of the same atoms, the oxidation state of which becomes greater.

For example:

c) intramolecular, in which the transfer of electrons occurs between two atoms that are part of the same particle of a substance, leading to the decomposition of the substance into simpler ones.

Redoxometry methods are based on oxidation-reduction reactions. A lot of methods have been developed. They are classified according to the standard (working, titrant) solution used. The most commonly used methods are:

Permanganatometry - a method that is based on the oxidizing ability of a working solution of potassium permanganate KMnO4. Titration is carried out without an indicator. It is used to determine only reducing agents during direct titration. Permanganatometry is based on the oxidation reaction of various reducing agents with a working solution of potassium permanganate, i.e. MnO4- ion. Oxidation with potassium permanganate can be carried out in acidic, neutral and alkaline environments. In a strongly acidic environment, permanganate ions (MnO4-) have a high redox potential, being reduced to Mn2+, and they are used to determine many reducing agents: MnO4- + 8H+ + 5e = Mn2+ + 4H2O

In an alkaline environment, MnO4- is reduced to manganate ion: MnO4- + e = MnO42-

In a neutral or slightly alkaline environment, the permanganate ion is reduced to permanganic acid MnO(OH)2 or to MnO2: MnO4- + 2H2O + 3e = MnO2v + 4OH-

The KMnO4 solution is a titrant with a set titer. In this regard, before using it in the analysis, KMnO4 solution is used as a titrant.

Iodometry- a method in which the working titrated solution is a solution of free iodine in CI. The method allows the determination of both oxidizing agents and reducing agents. Starch serves as an indicator. The iodometric method of titrimetric analysis is based on the reaction: I2 + 2e = 2I-

When determining oxidizing agents, a solution of sodium thiosulfate is used as a titrant, which reacts with the released iodine (substituent) in an equivalent amount. Na 2 S 2 O 3 -thiosulfate

32. Potentiometry- a research method based on the thermodynamic relationships between the EMF of electrochemical circuits, on the one hand, and the physicochemical and parameters of solutions and chemical reactions, on the other.

Inert electrodes- a plate or wire made of difficult-to-oxidize metals - platinum, gold, palladium. They are used to measure E in solutions containing a redox couple (for example, /).

Membrane electrodes of various types have a membrane on which membrane potential E arises. The value of E depends on the difference in concentrations of the same ion on different sides of the membrane. The simplest and most commonly used membrane electrode is the glass electrode.

Mixing insoluble salts such as AgBr, AgCl, AgI and others with some plastics (rubbers, polyethylene, polystyrene) led to the creation of ion-selective electrodes that selectively adsorb these ions from solution due to the Paneth-Faience-Hahn rule. Since the concentration of detectable ions outside the electrode differs from that inside the electrode, the equilibria on the membrane surfaces are different, which leads to the appearance of a membrane potential.

Most often, potentiometers are used for direct measurements of pH, indicators of the concentrations of other ions pNa, pK, pNH₄, pCl and mV. Measurements are carried out using appropriate ion-selective electrodes.

To measure pH, a glass electrode and a reference electrode - silver chloride - are used. Before carrying out analyses, it is necessary to check the calibration of pH meters using standard buffer solutions, the fixation of which is attached to the device.

In addition to direct determinations of pH, pNa, pK, pNH₄, pCl and others, pH meters allow potentiometric titration of the ion being determined.

Potentiometric titration.

Potentiometric titration is carried out in cases where chemical indicators cannot be used or when a suitable indicator is not available.

In potentiometric titration, potentiometer electrodes placed in the titrated solution are used as indicators. In this case, electrodes are used that are sensitive to titrated ions. During the titration process, the ion concentration changes, which is recorded on the measuring scale of the potentiometer. Having recorded the potentiometer readings in pH or mV units, plot their dependence on the titrant volume (titration curve), determine the equivalence point and the volume of titrant consumed for titration. Based on the data obtained, a potentiometric titration curve is constructed.

The potentiometric titration curve has a form similar to the titration curve in titrimetric analysis. The titration curve is used to determine the equivalence point, which is located in the middle of the titration jump. To do this, tangents are drawn to sections of the titration curve and the equivalence point is determined in the middle of the tangent of the titration jump. The change in ∆рН/∆V acquires the greatest value at the equivalence point.

The equivalence point can be determined even more accurately by the Gran method, which plots the dependence of ∆V/∆E on the titrant volume. Using the Gran method, potentiometric titration can be carried out without bringing it to the equivalence point.

Potentiometric titration is used in all cases of titrimetric analysis.

Acid-base titration uses a glass electrode and a reference electrode. Since the glass electrode is sensitive to changes in the pH of the medium, when they are titrated, changes in the pH of the medium are recorded on the potentiometer. Acid-base potentiometric titration is successfully used in the titration of weak acids and bases (pK≤8). When titrating mixtures of acids, it is necessary that their pK differ by more than 4 units, otherwise part of the weaker acid is titrated together with the strong one, and the titration jump is not clearly expressed.

This allows the use of potentiometry to construct experimental titration curves, select indicators for titration and determine acidity and basicity constants.

In precipitation potentiometric titration, an electrode made of a metal that forms an electrode pair with the ions being determined is used as an indicator.

When complexometric titration is used: a) a metal electrode reversible to the ion of the metal being determined; b) a platinum electrode in the presence of a redox couple in the solution. When one of the components of the redox couple is bound by the titrant, its concentration changes, which causes changes in the potential of the indicator platinum electrode. Back titration of an excess EDTA solution added to a metal salt with a solution of an iron (III) salt is also used.

For redox titration, a reference electrode and a platinum indicator electrode, sensitive to redox couples, are used.

Potentiometric titration is one of the most used methods of instrumental analysis due to its simplicity, accessibility, selectivity and wide capabilities.

33. Electrode potentials and mechanisms of their occurrence. To determine the direction and completeness of redox reactions between redox systems in aqueous solutions, the values electrode potentials these systems. The mechanism of occurrence of electrode potentials, their quantitative determination, processes that are accompanied by the occurrence of electric current or caused by electric current are studied by a special branch of chemistry - electrochemistry. By combining an electrode representing the redox system under study with a standard hydrogen electrode, the electrode potential E of this system is determined. In order to be able to compare the redox properties of different systems based on their electrode potentials, it is necessary that the latter also be measured under standard conditions. These are usually an ion concentration of 1 mol/l, a pressure of gaseous substances of 101.325 kPa and a temperature of 298.15 K. Potentials measured under such conditions are called standard electrode potentials and are designated Eo. They are often also called redox potentials or redox potentials, representing the difference between the redox potential of the system under standard conditions and the potential of a standard hydrogen electrode. The standard electrode potential is the potential of a given electrode process at concentrations of all substances involved in it equal to unity. Standard electrode potentials for redox systems are given in reference literature. These systems are written in the form of equations of reduction half-reactions, on the left side of which there are atoms, ions or molecules that accept electrons (oxidized form). The electrochemical series of voltages characterizes the properties of metals in aqueous solutions: the lower the electrode potential of the metal, the easier it is to oxidize and the more difficult it is to be reduced from its ions; metals having negative electrode potentials, i.e. those standing in the voltage series to the left of hydrogen are capable of displacing it from dilute acid solutions; each metal is capable of displacing (reducing) from salt solutions those metals that have a higher electrode potential. Under conditions different from standard ones, the numerical value of the equilibrium electrode potential for the redox system, written in the form, is determined by Nernst equation: where and are the electrode and standard potentials of the system, respectively; R – universal gas constant; T – absolute temperature; F – Faraday constant; n is the number of electrons involved in the redox process. C(Red) and C(Ox) are the molar concentrations of the reduced and oxidized forms of the compound, respectively. For example, for a redox system, the Nernst equation has the form

(REDOXOMETRY, OXIDIMETRY)

Essence and classification of redox titration methods

Permanganatometry is a method that is based on the oxidizing ability of a working solution of potassium permanganate KMnO4. Titration is carried out without an indicator. Used to determine only reducing agents during direct titration.

Iodometry is a method in which the working titrated solution is a solution of free iodine in CI. The method allows the determination of both oxidizing agents and reducing agents. Starch serves as an indicator.

Dichromatometry is based on the use of potassium dichromate K2Cr2O7 as a working solution. The method can be used for both direct and indirect determination of reducing agents.

Bromatometry is based on the use of potassium bromate KBrO3 as a titrant in the determination of reducing agents.

Iodatometry uses a solution of potassium iodate KIO3 as a working solution when determining reducing agents.

Vanadatometry makes it possible to use the oxidizing ability of ammonium vanadate NH4VO3. In addition to the listed methods, such methods as cerimetry (Ce4+), titanometry and others are also used in laboratory practice.

To calculate the molar mass equivalent of oxidizing agents or reducing agents, the number of electrons taking part in the redox reaction is taken into account (Me = M/ne, where n is the number of electrons e). To determine the number of electrons, it is necessary to know the initial and final oxidation states of the oxidizing agent and the reducing agent.

Of the large number of redox reactions, only those reactions are used for chemical analysis that:

· flow to the end;

· pass quickly and stoichiometrically;

· form products of a certain chemical composition (formula);

· allow you to accurately fix the equivalence point;

· do not react with by-products present in the test solution.

The most important factors influencing the reaction rate are:

· concentration of reacting substances;

· temperature;

· pH value of the solution;

presence of a catalyst.

In most cases, the reaction rate is directly dependent on the temperature and pH of the solution. Therefore, many determinations by redox titration must be carried out at a certain pH value and under heating.

Redox titration indicators

oxidative reduction titration

When analyzing by redox titration methods, direct, reverse and substitution titration are used. The equivalence point of redox titration is fixed both with the help of indicators and without indicators. The indicator-free method is used in cases where the oxidized and reduced forms of the titrant differ. At the equivalence point, the introduction of 1 drop of excess titrant solution will change the color of the solution. Without indicators, determinations can be made using the permanganatometric method, because at the equivalence point, one drop of potassium permanganate solution turns the titrated solution pale pink.

In the indicator method of fixing the equivalence point, specific and redox indicators are used. Specific indicators include starch in iodometry, which in the presence of free iodine turns intense blue due to the formation of a blue adsorption compound. Redox indicators are substances whose color changes when a certain redox potential value is reached. Redox indicators include, for example, diphenylamine NH(C6H5)2. When exposed to colorless solutions by its oxidizing agents, it turns blue-violet.

Redox indicators have the following requirements:

· the color of the oxidized and reduced forms must be different;

· the color change should be noticeable with a small amount of indicator;

· the indicator must react at the equivalence point with a very small excess of reducing agent or oxidizing agent;

· its action interval should be as short as possible;

· the indicator must be resistant to environmental components (O2, air, CO2, light, etc.).

The action interval of the redox indicator is calculated by the formula:

E = Ео ± 0.058/n,

where Eo is the normal redox potential of the indicator (in the reference book), n is the number of electrons accepted in the process of oxidation or reduction of the indicator.

Permanganatometry

Permanganatometry is based on the oxidation reaction of various reducing agents with a working solution of potassium permanganate, i.e. MnO4- ion. Oxidation with potassium permanganate can be carried out in acidic, neutral and alkaline environments

In a strongly acidic environment, permanganate ions (MnO4-) have a high redox potential, being reduced to Mn2+, and they are used to determine many reducing agents:

MnO4- + 8H+ + 5e = Mn2+ + 4H2O

E0 MnO4- / Mn2+ = 1.51 V

In an alkaline environment, MnO4- is reduced to manganate ion:

MnO4- + e = MnO42-

In a neutral or slightly alkaline environment, the permanganate ion is reduced to permanganic acid MnO(OH)2 or to MnO2:

МnО4- + 2Н2О + 3е = МnО2↓ + 4ОН-

E0 MnO4- / MnO2 = 0.59 V

When titrating with permanganate, indicators are not used, since the reagent itself is colored and is a sensitive indicator: 0.1 ml of 0.01 M KMnO4 solution turns 100 ml of water pale pink. As a result of the reaction of potassium permanganate with a reducing agent in an acidic medium, colorless Mn2+ ions are formed, which makes it possible to clearly determine the equivalence point.

The KMnO4 solution is a titrant with a set titer. In this regard, before using it in the analysis as a titrant, the KMnO4 solution is standardized according to the concentration of solutions of the starting substances of shawelic acid or sodium oxalate. A solution of potassium permanganate is very difficult to obtain in pure form. It is usually contaminated with traces of manganese(IV) oxide. Additionally, pure distilled water usually contains traces of substances that reduce potassium permanganate to form manganese(IV) oxide:

4 KMnO4 + 2H2O = 4 MnO2↓ + 4OH- + 3O2

When stored in solid form, potassium permanganate decomposes under the influence of light, also becoming contaminated with MnO2:

КМnО4 = К2МnО4 + МnО2↓ + О2

A solution of potassium permanganate can be prepared from a standard titer and a sample taken on a technical scale. In the first case, the contents of the ampoule are transferred quantitatively into a 2-liter volumetric flask, rinsing the ampoule and funnel with warm distilled water. Add a small volume of hot water to the volumetric flask to dissolve the crystals, then cool the resulting solution to room temperature, bring the volume of the solution to the mark and stir. The molar concentration of the resulting solution is 0.05 mol/l.

In the second case, weigh out a sample of potassium permanganate weighing 1.6 g on a technical scale in a beaker or on a watch glass, place it in a beaker and dissolve it in hot distilled water while thoroughly mixing the resulting solution, trying to ensure that all KMnO4 crystals dissolve. Then carefully pour the solution through a funnel into a 1-liter volumetric flask and mix thoroughly, after closing the flask with a ground-in stopper (do not use a rubber stopper). Leave the prepared KMnO4 solution for 7-10 days, then filter the solution through a funnel with glass wool or carefully pour it into another bottle using a siphon. It is imperative to store the KMnO4 solution in dark bottles, protected from light, to prevent decomposition.

The titer of a potassium permanganate solution prepared from a sample can be determined using oxalic acid H2C2O4*2H2O or sodium oxalate Na2C2O4.

Determination of nitrite ions in solution

In a neutral or alkaline environment, nitrites do not react with potassium permanganate; in a hot acidic solution they are oxidized to nitrates:

5КNO3 + 2КМnО4 + 3Н2SO4 = 2MnSO4 + 5КNO2 + K2SO4 + 3H2O

When slowly titrating an acidified solution of sodium nitrite with a solution of potassium permanganate, reduced results are obtained because nitrites are easily oxidized by acids to form nitrogen oxides:

2NO2- + 2H+ → 2 HNO2 → NO2- + NO + H2O

Therefore, to avoid losses, you can use the back titration method or the Lynge method - titration of an acidified solution of potassium permanganate with a solution of sodium nitrite.

Determination of calcium in calcium carbonate

Determination of calcium in solution by permanganatometric titration is possible by reverse or substitution titration. In the first case, a precisely measured excess of a titrated solution of oxalic acid is introduced into a solution containing calcium. The resulting CaC2O4 + H2SO4 precipitate, CaC2O4, is filtered off, and the residue that is not included in the oxalic acid reaction is titrated with a standard solution of potassium permanganate. Based on the difference between the introduced volume and the residue, it is determined how much oxalic acid was required for the precipitation of Ca2+, which will be equivalent to the calcium content in the solution.

According to the method of substitution titration, Ca2+ is isolated in the form of a precipitate of CaC2O4, which is filtered, washed and dissolved in H2SO4 or HC1.

CaC2O4 + H2SO4 → H2C2O4 + CaSO4

The resulting oxalic acid is titrated with a standard solution of potassium permanganate, the amount of which is equivalent to the calcium content in the solution.

Iodometry

The iodometric method of titrimetric analysis is based on the reaction:

I2 + 2e= 2I- ; Ео I2/3I- = 0.545 V

(REDOXOMETRY, OXIDIMETRY)

Essence and classification of redox titration methods

Dichromatometry is based on the use of potassium dichromate K2Cr2O7 as a working solution. The method can be used for both direct and indirect determination of reducing agents.

Bromatometry is based on the use of potassium bromate KBrO3 as a titrant in the determination of reducing agents.

Iodatometry uses a solution of potassium iodate KIO3 as a working solution when determining reducing agents.

To calculate the molar mass equivalent of oxidizing agents or reducing agents, the number of electrons participating in the redox reaction is taken into account (Me = M/ne, where n is the number of electrons e). To determine the number of electrons, it is necessary to know the initial and final oxidation states of the oxidizing agent and the reducing agent.

Of the large number of redox reactions, only those reactions are used for chemical analysis that:

· flow to the end;

· pass quickly and stoichiometrically;

· form products of a certain chemical composition (formula);

· allow you to accurately fix the equivalence point;

· do not react with by-products present in the test solution.

The most important factors influencing the reaction rate are:

· concentration of reacting substances;

· temperature;

· pH value of the solution;

presence of a catalyst.

Redox titration indicators

oxidative reduction titration

Redox indicators have the following requirements:

· the color of the oxidized and reduced forms must be different;

· the color change should be noticeable with a small amount of indicator;

· the indicator must react at the equivalence point with a very small excess of reducing agent or oxidizing agent;

· its action interval should be as short as possible;

· the indicator must be resistant to environmental components (O2, air, CO2, light, etc.).

The action interval of the redox indicator is calculated by the formula:

E = Ео ± 0.058/n,

where Eo is the normal redox potential of the indicator (in the reference book), n is the number of electrons accepted in the process of oxidation or reduction of the indicator.

Permanganatometry

In a strongly acidic environment, permanganate ions (MnO4-) have a high redox potential, being reduced to Mn2+, and they are used to determine many reducing agents:

MnO4- + 8H+ + 5e = Mn2+ + 4H2O

E0 MnO4- / Mn2+ = 1.51 V

In an alkaline environment, MnO4- is reduced to manganate ion:

MnO4- + e = MnO42-

In a neutral or slightly alkaline environment, the permanganate ion is reduced to permanganic acid MnO(OH)2 or to MnO2:

МnО4- + 2Н2О + 3е = МnО2↓ + 4ОН-

E0 MnO4- / MnO2 = 0.59 V

4 KMnO4 + 2H2O = 4 MnO2↓ + 4OH- + 3O2

When stored in solid form, potassium permanganate decomposes under the influence of light, also becoming contaminated with MnO2:

КМnО4 = К2МnО4 + МnО2↓ + О2

The titer of a potassium permanganate solution prepared from a sample can be determined using oxalic acid H2C2O4*2H2O or sodium oxalate Na2C2O4.

Determination of nitrite ions in solution

In a neutral or alkaline environment, nitrites do not react with potassium permanganate; in a hot acidic solution they are oxidized to nitrates:

5КNO3 + 2КМnО4 + 3Н2SO4 = 2MnSO4 + 5КNO2 + K2SO4 + 3H2O

2NO2- + 2H+ → 2 HNO2 → NO2- + NO + H2O

Therefore, to avoid losses, you can use the back titration method or the Lynge method - titration of an acidified solution of potassium permanganate with a solution of sodium nitrite.

Determination of calcium in calcium carbonate

According to the method of substitution titration, Ca2+ is isolated in the form of a precipitate of CaC2O4, which is filtered, washed and dissolved in H2SO4 or HC1.

CaC2O4 + H2SO4 → H2C2O4 + CaSO4

The resulting oxalic acid is titrated with a standard solution of potassium permanganate, the amount of which is equivalent to the calcium content in the solution.

Iodometry

The iodometric method of titrimetric analysis is based on the reaction:

I2 + 2e= 2I-; Ео I2/3I- = 0.545 V

This equation is written schematically, since in practice, to increase the solubility of I2, a solution of KI is used, which forms a complex K with I2. Then the equation for iodometric determination looks like this:

The amount of the substance being determined is judged by the amount of iodine absorbed or released. Substances whose redox potential is below 0.545 V will be reducing agents (SO2, Na2S2O3, SnCl2, etc.) and, therefore, a reaction will occur with the absorption of iodine. The balance will shift to the right. Substances whose redox potential is greater than 0.545 V will be oxidizing agents (KMnO4, MnO2, K2Cr2O7, Cl2, Br2, etc.) and direct the reaction to the left, towards the release of free iodine.

In this regard, the iodometric method is used both for the determination of reducing agents and oxidizing agents. Iodometric determinations are carried out in an acidic environment, since in an alkaline environment a hypoiodide ion can be formed, the oxidizing ability of which is higher than that of iodine, which can contribute to the occurrence of side processes, in particular, oxidize the thiosulfate ion to sulfate and the results will be distorted.

When determining strong reducing agents (Eo is much greater than 0.545 V), direct titration is used, and weak ones (Eo is close to 0.545 V) are used by reverse titration. The working solution (titrant) is a solution of I2. Oxidizing agents are determined only by substitution titration, since When using potassium iodide as a working solution, it is impossible to fix the equivalence point (the moment the release of iodine ceases). When determining oxidizing agents, a solution of sodium thiosulfate is used as a titrant, which reacts with the released iodine (substituent) in an equivalent amount.

A freshly prepared 1% starch solution is used as an indicator in iodometry. When starch interacts with iodine, 2 processes occur - complexation and adsorption, which results in the formation of a blue compound. The sensitivity of the reaction with starch is high, but decreases sharply with increasing temperature. Starch should be added to the titrated solution only when the main amount of iodine has already been titrated, otherwise the starch forms such a strong compound with excess iodine that excess consumption of sodium thiosulfate is observed.

Standardization of sodium thiosulfate solution with potassium dichromate

It is impossible to titrate thiosulfate directly with potassium dichromate, since it reacts nonstoichiometrically with all strong oxidizing agents (dichromate, permanganate, bromate, etc.). Therefore, the substitution method is used, first using the stoichiometric reaction between dichromate and iodide:

Cr2O72- + 6I- + 14 H+ = 2Cr3+ + 3I2 + 7H2O (1)

Iodine, released in an amount equivalent to dichromate, is titrated with a standard solution of thiosulfate:

I2 + 2S2O32- = 2I- + S4O62- (2)

For reaction (1) to occur, a high concentration of hydrogen ions is required, because in an acidic environment, the redox potential of the Cr2O72-/ 2Cr3+ pair increases, i.e. the oxidizing ability of potassium dichromate is enhanced. Excess I- dissolves the released iodine and lowers the potential of the redox couple I3-/ 3I-, thus increasing the emf of reaction (1). Before titrating the released iodine, it is necessary to reduce the acidity of the solution by diluting it with water to prevent a side reaction from occurring:

2H+ + S2O32- = H2S2O3 = H2O + SO2 + S

Dichromatometry

The essence of dichromatometric titration

Dichromatometric titration is one of the methods of redox titration, based on the use of potassium dichromate K2Cr207 as an oxidizing agent. When exposed to reducing agents, the dichromate ion Cr2O72- acquires six electrons and is reduced to Cr3+

Сr2О72- + 6е + 14Н+ = 2Сr3+ + 7Н20

Therefore, the molar mass of potassium dichromate equivalent is 1/6 of the molar mass. From the reaction equation it is clear that the reduction of Cr2O72- anions to Cr3+ cations occurs in the presence of H+ ions.

Therefore, titrate with dichromate in an acidic medium. The redox potential of the Сr2О72-/2Сr3+ system is 1.36 V. At [H+] = 1 mol/l. Therefore, in an acidic environment, potassium dichromate is a strong oxidizing agent. Therefore, dichromatometry is successfully used to determine almost all reducing agents determined permanganometrically. Dichromatometry even has some advantages over permanganatometry.

Potassium dichromate is easily obtained in chemically pure form by recrystallization. Therefore, its standard solution is prepared by dissolving an accurate sample. Solutions of potassium dichromate are extremely stable when stored in closed containers; it does not decompose even when the acidified solution is boiled and practically does not change when the solution stands.

In addition, potassium dichromate is more difficult than permanganate to be reduced by organic substances. Therefore, it does not oxidize random impurities of organic substances. This also determines the constancy of its titer in solution. Potassium dichromate does not oxidize (without heating) chloride ions. This allows them to titrate reducing agents in the presence of HCl.

The indicator most often used in dichromatometric titration is diphenylamine, which turns the solution blue at the slightest excess of dichromate. Diphenylamine belongs to the group of so-called redox indicators (redox indicators). They are redox systems that change color when a reduced form changes to an oxidized one, or vice versa.

If we designate the oxidized form of the indicator Indoxid. the reduced form Indrestored, and the number of transferred electrons is n, then the transformation of one form of such an indicator into another can be depicted by a diagram;

Indoxid. ↔Ind restored -ne-

Each redox indicator is characterized by a certain redox potential. For dephenylamine it is +0.76 V. The oxidized form of diphenylamine is blue, and the reduced form is colorless.

In addition to diphenylamine, redox indicators include ferroin, sodium diphenylaminosulfonate, phenylanthranilic acid, etc.

Fe2+ ions are determined dichromatometrically in HCl solutions or in sulfuric acid solutions. Chloride ions do not interfere with determination if their concentration does not exceed 1 mol/l.

However, when Fe2+ salts are titrated with dichromate, Fe3+ cations accumulate in the solution, the redox potential of the Fe3+↔Fe2+ system increases and diphenylamine is oxidized. Therefore, a blue color may appear when the equivalence point has not yet been reached.

To lower the redox potential of the Fe2+ ↔ Fe3+ system, in addition to diphenylamine and hydrochloric acid, orthophosphoric acid is added to the solution. The latter masks interfering Fe3+ ions, binding them into a stable, colorless Fe (HP04)+ complex.

Preparation of a standard solution of potassium dichromate

A standard solution is prepared by dissolving an accurately weighed portion of potassium dichromate (reagent grade) in a volumetric flask. Potassium dichromate must first be recrystallized from an aqueous solution and dried at 150°C.

Prepare 100 ml of approximately 0.1 N potassium dichromate standard solution. It was noted above that when interacting with reducing agents in an acidic environment, the dichromate ion Cr2O72- acquires six electrons. Therefore, the molar mass of the equivalent of K2Cr207 is 294.20:6 = 49.03 g/mol and to prepare 0.1 l of 0.1 N solution, 49.03 * 0.1 * 0.1 = 0.4903 g of potassium dichromate will be required .

Take about 0.5 g of freshly recrystallized potassium dichromate in a small test tube and weigh it on an analytical balance. Using a funnel, transfer the contents of the test tube into a 100 ml volumetric flask. Weigh the test tube again and use the difference to find the mass of the sample

Dissolve a sample of potassium dichromate in distilled water, remove the funnel and, using a pipette, bring the volume of the solution in the flask to the mark. Calculate the titer and normal concentration of potassium dichromate solution.

Let us assume that the portion of potassium dichromate was 0.4916 g. Then the titer of the solution

T= m/V= 0.4916/100 = 0.004916 g/ml,

and normal concentration (molar concentration equivalent)

c = 0.004916*1000 /49.03 = 0.1003.

Determination of iron (II) content in solution

Iron is determined dichromatometrically mainly in ores, alloys, slags and other materials. However, when they dissolve, iron partially transforms into Fe3+ ions. Therefore, before determination it is necessary to reduce Fe3+ to Fe2+. This is achieved by the action of metals (or their amalgams), for example, by the action of metallic zinc:

2Fe3++ Zn = 2Fe2+ + Zn2+

Excess zinc is removed from the solution by filtration (for example, through cotton wool). The essence of the reaction used for the dichromatometric determination of Fe2+ can be expressed by the equation

6Fe2++ Сr2О72- + 14Н+ → 6Fe3+ + 2Сr3+ + 7Н20

The determination consists of direct titration of the analyzed solution with a standard solution of potassium dichromate in the presence of diphenylamine:

6FeS04 + K2Сr207 + 7H2S04 = 3Fe 2 (S04)3 + Cr2 (S04)3 + K2S04 + 7H20

1 Сr2О72- + 14H+ + 6е = 2Cr3+ + 7Н20

6 Fe 2+ - e= Fe3+

Sulfuric acid is added to the test solution to maintain a high acidity of the medium and phosphoric acid to bind the accumulated Fe3+ ions, which can prematurely convert diphenylamine into an oxidized (colored) form.

Chemical elements with variable oxidation states can be quantified titrimetrically using an oxidation-reduction reaction (ORR). Oxidation-reduction (RO) titration methods or red-ox methods are titrimetric methods based on the use of redox reactions.

Redox titrations can be divided into:

1 According to the nature of the titrant:

– oxidimetric– methods for determining reducing agents using an oxidizing titrant;

– reductometric– methods for determining an oxidizing agent using a reducing titrant.

2 By the nature of the reagent (titrant) interacting with a certain substance:

– KMnO 4 – permanganatometry;

– KBrO 3 – bromatometry;

– KI,Na 2 S 2 O 3 – iodometry;

– I 2 – iodymetry;

– Br 2 – bromometry;

– Ce(SO 4) 2 – cerimetry

Depending on the analytical problem being solved, direct, reverse, and substitution titrations are used in redoximetry. Both inorganic and organic substances can be quantified redoximetrically. For example, by reduction with potassium permanganate in a strongly alkaline medium, methanol, formic, tartaric, citric, salicylic acids, as well as glycerol, phenol, formaldehyde, etc. can be determined.

Schematically, the ORR, taking into account the law of electroneutrality of the solution, can be depicted as follows:

mOX1+ nRed2↔ mRed1+ n OX2

Here, the indices 1 and 2 refer to substances 1 and 2 in oxidized (Ox1 and Ox2) and reduced (Red1 and Red2) forms. During ORR, substance Ox1 with a higher electron affinity (oxidizing agent) adds electrons, lowers its oxidation state, and is reduced, and substance Red2 with lower electron affinity (reducing agent) is oxidized.

The oxidized and reduced forms of substances reacting in ORR form redox (oxred-, redox) pairs Ox1/Red1 and Ox2/Red2, and transformations of the Ox+ze Red type are called oxed-(redox) transitions or redox half-reactions.

§2. Redox potential.

Nernst equation.

Redox processes, like all dynamic processes, are reversible to one degree or another. The direction of reactions is determined by the ratio of the electron-donating properties of the components of the system of one redox half-reaction and the electron-acceptor properties of the second (provided that the factors influencing the shift of equilibrium chemical reactions are constant). The movement of electrons during redox reactions gives rise to a potential. Thus, potential, measured in volts, serves as a measure of the redox ability of a compound.

To quantify the redox properties of redox pairs, redox (oxidation-reduction) potentials are used. When calculating the redox potential, use Nernst equation:

E (Ox/Red) = E 0 (Ox/Red) +

where E(Ox/Red) is the real or equilibrium redox potential, V;

E 0 (Ox/Red) - standard redox potential, equal to equilibrium at a(Ox) = a(Red) = 1 mol/dm 3 ;

R is the universal gas constant (8.31 J/K mol);

T - absolute temperature, K; F - Faraday number (96500 C/mol);

z is the number of electrons participating in the redox transition in OX+ze dRed;

a(OX) and a(Red) - activities of the oxidized and reduced forms of the substance, respectively, mol/dm 3.

When substituting the values R, F and T = 298 K into the Nernst equation, as well as passing to the decimal logarithm, we obtain

E(Ox/Red) =E 0 (Ox/Red) +

The redox potential also depends on the acidity of the medium, complexation or precipitation of one of the components of the redox pair during the redox transition. The greater the concentration of hydrogen ions in the solution, the greater the oxidizing ability of the oxidized form of the redox pair substance and the greater the E(Ox/Red).

When choosing a titrant substance in redoximetry, a qualitative and quantitative assessment of the possibility (direction) and completeness of the passage of the ORR between the titrant and the substance being determined is carried out.

A qualitative assessment is carried out by comparing the tabulated values of E 0 (Ox|Red) of the titrant substance and the substance being determined, given in analytical, chemical and physicochemical reference books.

Permanganatometry

The permanganatometric method of volumetric analysis is a method based on the oxidation of various substances with potassium permanganate (KMnO 4).

Depending on the conditions under which the oxidation-reduction reaction occurs, MnO 4 – ions can accept a different number of electrons.

In an acidic environment:

In a neutral environment:

In an alkaline environment:

The normal potential of the system is E 0 (MnO 4 – ⁄Mn 2+) = +1.52 V, and E 0 (MnO 4 – ⁄ MnO 2) = +0.57 V, therefore potassium permanganate in an acidic environment has strong oxidizing properties and is capable of oxidizing many substances.

The equivalent of potassium permanganate in an acidic medium is:

M(1/zKMnO 4) = M(KMnO 4)/n e = 158/5 = 31.608 g/mol

In laboratory practice, potassium permanganate is used in the form of solutions of various concentrations. Usually a 0.1 N solution of KMnO 4 is used, although in some cases 0.01 N, 0.05 N, 0.2 N solutions are used.

Preparation of working solution KMnO 4

Potassium permanganate, used to prepare a working solution of KMnO 4, usually contains a number of impurities, the most significant of which are manganese (IV) compounds. In addition, in the first days after preparing the solution, KMnO 4 is reduced by organic impurities contained even in distilled water. As a result, the concentration of the KMnO 4 solution changes:

Therefore, first prepare a solution of approximate concentration. For example, to prepare 500 ml of a 0.1 N KMnO 4 solution, calculate the required weight of the substance using the formula:

m(KMnO 4) = N(KMnO 4) M(1/zKMnO 4) V

m=31.608 0.1 0.5≈1.58g.

The sample is dissolved in a 0.5 liter volumetric flask. The solution is poured into a dark glass bottle and left in a dark place for at least a week. During this time, the permanganate will oxidize all the impurities contained in the water, and the manganese dioxide MnO 2 formed as a result of partial reduction of the permanganate will settle to the bottom of the bottle. The solution is filtered from MnO 2 and stored in dark flasks. Obviously, after this they begin to standardize the solution.

Ammonium oxalate (NH 4) 2 C 2 O 4 H 2 O, sodium oxalate Na 2 C 2 O 4 and oxalic acid H 2 C 2 O 4 2H 2 O are usually used as starting materials to set the exact concentration of the KMnO 4 solution The most convenient is sodium oxalate, because... it crystallizes without water and is not hygroscopic.

The reaction is autocatalytic, so the solution should be heated to speed up the process.

The potential difference for this reaction is determined by subtracting from the normal potential of the MnO 4 – /Mn 2+ system (E 0 = +152V) the normal potential of the 2CO 2 /C 2 O 4 2– system (E 0 = –0.49V), then E = +1.52–(–0.49)=2.01V

A large potential difference indicates that the reaction is irreversible.

All products of this reaction are colorless, while the KMnO 4 solution is red-violet. Therefore, the reaction must be accompanied by discoloration of the added permanganate solution. If you add 2-3 drops of KMnO 4 solution to an acidic sodium oxalate solution, the colorless solution will turn pink, indicating the presence of unreacted KMnO 4 . The color disappears only after a few minutes. This indicates that the reaction rate is initially low. Discoloration of the solution after adding subsequent drops of KMnO 4 solution occurs faster and faster, and finally will occur almost instantly up to the equivalence point. An extra drop of KMnO 4 will color the titrated solution a permanent pink color.

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