LECTURE No. 9

Lecture outline:

1. Redox systems, their characteristics.

2. Redox potentials, their experimental measurement. Standard redox potential as a measure of force

oxidizing agent and reducing agent.

3. Application of standard redox potentials to determine products, direction and sequence of redox reactions.

4. Real redox potentials. Nernst equation.

Redox systems, their characteristics.

Many reactions of interest in analytical chemistry are redox and are used in both qualitative and quantitative analysis.

Oxidation-reduction reactions (ORR) are reactions involving a change in the oxidation state of the reacting substances. In this case, the change in the oxidation state occurs with the addition and loss of electrons.

The processes of electron gain and loss are considered as half-reactions of reduction and oxidation, respectively:

aOk1 + ne cBoc1 (reduction) bBoc2 – ne dOk2 (oxidation) In each half-reaction, a substance in a higher oxidation state is called an oxidized form (Ok), and a substance in a lower oxidation state is called a reduced form (Boc).

The oxidized and reduced forms of a substance represent a conjugate redox couple (redox couple). In a redox couple, the oxidized form (Ok) is an electron acceptor and is reduced, the reduced form (Boc) acts as an electron donor and is oxidized.

The half-reactions of oxidation and reduction are not possible from one another - if there is an electron donor, then there must also be an acceptor. The overall redox reaction actually occurs:

aOk1 + bBoc2 cBoc1 + dOk In this case, the number of given and received electrons must be the same.

For example, consider a redox reaction:

2Fe3+ + Sn2+ 2Fe2+ + Sn4+ The corresponding half-reactions can be written as:

2Fe3+ + 2e 2Fe2+ Sn2+ – 2e Sn4+ This redox reaction involves two electrons and there are two redox pairs Fe3+/Fe2+ and Sn4+/Sn2+, each of which contains oxidized (Fe3+, Sn4+) and reduced (Fe2+, Sn2+) forms .

Redox potentials, their experimental measurement. Standard redox potential as a measure of the strength of an oxidizing agent and a reducing agent.

The effectiveness of the oxidative or reduction properties of a given substance (the ability to donate or accept electrons) depends on its nature, the conditions for the redox reaction, and is determined by the value of the redox potential (ORP) of the half-reaction (redox pair). This potential is experimentally measured using a redox electrode consisting of an inert material M (platinum, gold, graphite, glassy carbon) immersed in an aqueous solution in which there are oxidized and reduced forms of this substance. Such an electrode is designated as follows:

M | Ok, Vos On the surface of such a reversible electrode the following reaction occurs:

Ok + ne Boc, as a result of which a potential arises equal to the redox potential of the redox pair under study.

For example, if a platinum electrode is immersed in a solution containing iron(III) (oxidized form) and iron(II) (reduced form) chlorides (Pt | FeCl3, FeCl2), then the redox reaction Fe3+ + e Fe2+ and an electrode potential arises equal to the redox potential of the Fe3+/Fe2+ redox couple.

It is not possible to measure the absolute value of the redox potential; therefore, in practice, the ORP value of the redox pair under study is measured relative to any standard reference half-reaction and an electrode created on its basis (reference electrode). The standard half-reaction must be reversible, and the reference electrode must have a constant and reproducible potential and be of fairly simple design.

As a universal reference electrode for measuring redox potential, a standard hydrogen electrode (SHE) is used, which consists of a platinum plate, coated with a layer of finely dispersed platinum (platinum black), and immersed in a solution of hydrochloric (or sulfuric) acid with Pt( H2) (p =1 atm) | HCl, hydrogen, mol/l || unit – аН+ = 1:

ion activity a(H+) = 1 equal to H2 (gas) platinum plate, hydrogen molecules coated with finely dispersed hydrogen molecules adsorbed on platinum platinum plate HCl (platinum black) Pt H 2H+ + 2e Platinum is washed by a flow of hydrogen gas under a pressure of 1 atm (101.3 kPa), Standard conditions: t = 250C (298 K), p(H2) = 1 atm (101.3 kPa), which is sorbed on the porous surface of platinum black. Denoted stana(H+) = 1 mol/l ESVE = E2H /H = dart hydrogen electrode as follows: + Pt(H2) (p = 1 atm) | HCl (aH+ = 1) A half-reaction occurs on the surface of such a reversibly operating electrode:

the potential of which is conventionally assumed to be zero at any temperature, that is, the potential of a standard hydrogen electrode ESVE = 0.

It should be noted that a standard hydrogen electrode is not a redox electrode, but a galvanic cell is assembled. To measure ORP refers to the so-called electrodes of the first kind, the potential is composed of the SVE activity of the corresponding cations - in which it depends on the ORP pair (half-reaction) being studied.

case from the activity of hydrogen cations.

To measure the ORP of a half-reaction, it is necessary to compose a galvanic cell from the redox couple (half-reaction) - this is the EMF of the galvanic dart hydrogen electrode and the electrode on which the half-reaction under study flows, composed of this ORP of the half-reaction and SVE.

In this case, the recording diagram of the galvanic cell looks like this:

In this scheme, a vertical bar (|) means a potential jump at the “electrode-solution” phase boundary, and a double vertical bar (||) means elimination of the diffusion potential using a salt bridge.

The electromotive force (EMF) of a given galvanic circuit, that is, the potential difference between the half-reaction under study and the standard hydrogen electrode, is equal to the redox potential of the redox pair under study:

If the potential of the redox couple under study is measured under standard conditions - temperature 250C (298 K), pressure 1 atm (101.3 kPa) and the activities of the oxidized and reduced forms are equal to unity (aOk = aBoc = 1 mol/l), then it is called standard redox potential and denoted E0Ok/Boc.

Standard ORPs of many redox pairs are measured and their values ​​in volts are given in tables, for example:

The greater the E0Ok/Boc, the stronger the oxidized form is as an oxidizing agent and the weaker the reducing agent is the reduced form. And, conversely, the lower the E0Ok/Boc, the stronger the reducing agent is the reduced form and the weaker the oxidizing agent is the oxidized form.

From the data given in the table it is clear that molecular fluorine has the greatest oxidizing properties, and metallic magnesium has the greatest reducing properties. At the same time, fluorine and magnesium ions have practically no reducing and oxidizing properties, respectively.

A positive sign of the potential indicates the spontaneous occurrence of a reduction reaction in a pair with UHE, a negative sign indicates the spontaneous occurrence of an oxidation reaction. Thus, the potentials of strong oxidizing agents are always positive, and those of strong reducing agents are always negative. The sign convention was adopted in 1953 at the congress of the International Union of Pure and Applied Chemistry (IUPAC).

Application of standard redox potentials to determine products, direction and sequence of redox reactions.

From the thermodynamic theory of electromotive forces and electrode potentials it is known that the standard reaction potential E0 (standard EMF of the reaction), which is equal to the difference between the standard ORP of the redox pairs (half-reactions) participating in the reaction, is related to the standard change in the Gibbs energy G0 of the reaction by the equation:

where: n is the number of electrons participating in the redox reaction F is the Faraday number, 96500 C/mol From the thermodynamics of equilibrium processes it is also known that if the change in the Gibbs energy in any chemical reaction is less than zero, then this reaction spontaneously proceeds in the forward direction in in accordance with the notation of the reaction equation; if more than zero - in the opposite direction.

From here it is easy to see that with a positive difference between the standard redox pairs (half-reactions) participating in any redox reaction aOk1 + bBoc2 cBoc1 + dOk2, the change in the standard Gibbs energy is less than zero and the reaction under standard conditions proceeds in the forward direction:

In the case of a negative difference between the standard ORP of redox pairs (half-reactions) involved in a redox reaction, the change in the standard Gibbs energy is greater than zero and the reaction under standard conditions does not proceed in the forward direction, but proceeds in the opposite direction:

In other words, the redox reaction proceeds in the direction from stronger oxidizing and reducing agents to weaker ones. In this case, the reaction proceeds until a state of equilibrium is established.

For example, is it possible to oxidize iron(II) ions with stannous salt?

The proposed oxidation reaction proceeds according to the equation:

The standard ORP of redox pairs are: ESn4+/Sn2+ +0.15 V, EFe3+/Fe2+ +0.77 V. Then, according to the above, E0 = 0.15 – 0.77 = -0.62 V 0). This means that the reaction under standard conditions does not proceed in the forward direction, that is, it is impossible to oxidize iron(II) ions with tetravalent tin ions. On the contrary, the reaction proceeds in the opposite direction and oxidation of tin(II) ions by iron ions() is possible:

In this case, the standard reaction potential is positive E0 = 0.77 – 0.15 = 0.62 V > 0, and the change in the standard Gibbs energy is less than zero (G0

Thus, in accordance with standard redox potentials, the reaction proceeds in the direction from stronger oxidizing and reducing agents (Fe3+ and Sn2+) to weaker ones (Fe2+ and Sn4+).

Using standard redox potentials, it is possible to determine not only the direction, but also the sequence of redox reactions. In the case of several OVRs, the one whose standard potential E0 is the greatest comes first.

For example, when chlorine water acts on a solution containing iodide and bromidiones, the following reactions may occur:

The standard ORP of redox pairs involved in reactions are:

In this case, the strong oxidizing agent Cl2 (high standard ORP) will react first with the strongest reducing agent iodide ion (small standard ORP) and then with the bromide ion. This is indicated by the higher value of the standard potential for the reaction of chlorine with iodide (E0 = 1.36 – 0.54 = 0.82 V) than with bromide (E0 = 1.36 – 1.08 = 0.28 V).

Standard ORPs can also be used to determine the products of redox reactions.

For example, when tin(IV) chloride reacts with metallic iron, it is possible to reduce tin to Sn2+ or Sn0 and oxidize iron to Fe2+ or Fe3+. Wherein:

From the given values ​​of standard ORP it is clear that the Sn4+ ion exhibits greater oxidizing properties when reduced to Sn2+, and metallic iron is a stronger reducing agent when oxidized to the Fe2+ ion. Therefore, the reaction under study proceeds according to the equation:

This reaction also corresponds to the largest value of the standard potential equal to:

Thus, the products of the reaction between tin(IV) chloride and metallic iron are tin(II) and iron(II) chlorides:

Real redox potentials. Nernst equation.

The situation when all participants in the redox reaction are simultaneously in standard states (their activities, concentrations and activity coefficients are equal to unity) is often practically unrealistic and should not be considered hypothetical.

A redox reaction occurring under real conditions is characterized by work A, which is spent on the electrochemical transformation of one mole of a substance:

where: n is the number of electrons participating in the redox reaction F is the Faraday number, 96500 C/mol For the spontaneous reaction aOk1 + bBoc2 cBoc1 + dOk2 this work is the Gibbs energy:

Knowing that dividing by nF, changing signs and substituting the expression for K0, we get:

When the activities of all components are equal to unity, E = E0, that is, the reaction potential is equal to the standard potential.

The potential of any redox reaction (real E or standard E0) is equal to the difference in the corresponding redox potentials of the half-reactions of its components, then:

If the second half-reaction is the half-reaction 2Н+ + 2е Н2 (aH+ = 1, p = 1 atm) occurring under standard conditions, for which E2H+ /H E2H+ /H 0, then the reaction potential will be equal to the potential of the first half-reaction:

Then the expression for the redox potential of any half-reaction aOk + ne cBoc has the form:

where: EOk/Boc – real redox potential of the half-reaction E0Ok/Boc – standard redox potential of the half-reaction R – universal (molar) gas constant, 8.314 J/molK T – absolute temperature, K n – number of electrons involved in oxidation reduction reaction F – Faraday number, 96500 C/mol This expression is called the Nernst equation. Often, constant quantities in the Nernst equation are combined into one constant, and the natural logarithm is replaced by a decimal one (ln = 2.3lg). Then at 250C (298 K):

From the Nernst equation it follows that the standard redox potential is equal to the real redox potential of the half-reaction (redox pair) with the activities of all particles participating in the equilibrium equal to unity:

For example, for a half-reaction:

The standard redox potential depends only on temperature, pressure and the nature of the solvent.

In practice, it is more convenient to use concentrations rather than activities. In this case, the Nernst equation can be rewritten using the total concentrations of the oxidized (cOk) and reduced forms (cBoc). Since a = c (where is the activity coefficient and is the coefficient of the competing reaction), the Nernst equation takes the form:

where: EOk/Boc is the formal redox potential of the half-reaction. The formal ORP is equal to the real redox potential at total concentrations of oxidized and reduced forms equal to 1 mol/l, and given concentrations of all other substances present in the system:

For example, for a half-reaction:

Thus, the formal redox potential, in contrast to the standard one, depends not only on temperature, pressure and the nature of the solvent, but also on ionic strength, the occurrence of competing reactions and the concentration of particles that are not oxidized or reduced forms, but take part in the half-reaction (in this example H+).

When calculating redox potentials, the influence of ionic strength is often neglected, taking the ratio of activity coefficients equal to unity, and instead of activities in the Nernst equation, equilibrium concentrations are used ([Ok] = Ok cOk; [Boc] = Boc cBoc). Then:

All subsequent examples are recorded and calculated using this assumption.

When writing the Nernst equation for any redox half-reaction, a certain order and rules should be followed:

Correctly write down the redox half-reaction in compliance with stoichiometric coefficients and determine the number of electrons involved in it;

- determine the oxidized and reduced form;

Determine the components in the standard state (solid forms, poorly soluble gases with p = 1 atm, solvent molecules) and exclude them from writing the Nernst equation, since their activities are equal to unity;

- write down the Nernst equation taking into account stoichiometric coefficients and accompanying ions.

For example, write Nernst equations for the following redox pairs:

a) Cr2O72-/Cr3+ (in an acidic environment) - write the half-reaction: Cr2O72- + 14H+ + 6e 2Cr3+ + H2O (n = 6) - in this half-reaction Cr2O72- is the oxidized form, Cr3+ is the reduced form - H2O (solvent) in the standard state (a = 1) - we write the Nernst equation taking into account the stoichiometric coefficients and accompanying H+ ions:

b) AgCl/Ag - in this half-reaction AgCl is the oxidized form, Ag is the reduced form - AgCl and Ag0 in solid form, that is, in the standard state (a = 1) - we write the Nernst equation taking into account the stoichiometric coefficients and accompanying Cl- ions:

c) O2/H2O2 (in an acidic medium) - in this half-reaction O2 is the oxidized form, H2O2 is the reduced form - gaseous O2 in the standard state (a = 1) - we write the Nernst equation taking into account the stoichiometric coefficients and accompanying H+ ions:

d) O2/H2O2 (in an alkaline medium) - write the half-reaction: O2 + 2H2O + 2e H2O2 + 2OH- (n = 2) - in this half-reaction O2 is the oxidized form, H2O2 is the reduced form - gaseous O2 and H2O (solvent) in standard state (a = 1) - we write the Nernst equation taking into account the stoichiometric coefficients and accompanying OH- ions:

e) SO42-/SO32- (in an alkaline medium) - write the half-reaction: SO42- + H2O + 2e SO32- + 2OH- (n = 2) - in this half-reaction SO42- is the oxidized form, SO32- is the reduced form - H2O ( solvent) in the standard state (a = 1) - we write the Nernst equation taking into account the stoichiometric coefficients and accompanying OH- ions:

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There are three main types of redox reactions:

1. Intermolecular (intermolecular oxidation - reduction).

This type includes the most numerous reactions in which the atoms of an oxidizing element and a reducing element are located in different molecules of substances. The reactions discussed above belong to this type.

2. Intramolecular (intramolecular oxidation - reduction).

These include reactions in which an oxidizing agent and a reducing agent in the form of atoms of different elements are contained in the same molecule. Thermal decomposition reactions of compounds proceed according to this type, for example:

2KCIO 3 = 2KCI + 3O 2.

3. Disproportionation (auto-oxidation - self-healing).

These are reactions in which the oxidizing and reducing agent is the same element in the same intermediate oxidation state, which as a result of the reaction both decreases and increases. For example:

3CI 0 2 + 6 KOH = 5 KCI + KCIO 3 + 3H 2 O,

3HCIO = HCIO 3 + 2HCI.

Redox reactions play an important role in nature and technology. Examples of OVR occurring in natural biological systems include the photosynthesis reaction in plants and respiration processes in animals and humans. Fuel combustion processes occurring in the furnaces of boilers at thermal power plants and in internal combustion engines are an example of OVR.

ORRs are used in the production of metals, organic and inorganic compounds, and carry out the purification of various substances, natural and waste waters.

9.5. Oxidation – reduction (electrode) potentials

A measure of the redox potential of substances is their electrode or redox potential j ox / Red (redox potentials).1 The redox potential characterizes the redox system, consisting of the oxidized form of the substance (Ox), the reduced form (Red) and electrons. It is customary to write redox systems in the form of reversible reduction reactions:

Ox + ne - D Red.

Mechanism of occurrence of electrode potential. We will explain the mechanism of occurrence of the electrode or redox potential using the example of a metal immersed in a solution containing its ions. All metals have a crystalline structure. The crystal lattice of a metal consists of positively charged Men + ions and free valence electrons (electron gas). In the absence of an aqueous solution, the release of metal cations from the metal lattice is impossible, because this process requires large energy costs. When a metal is immersed in an aqueous solution of a salt containing metal cations, polar water molecules, correspondingly oriented at the surface of the metal (electrode), interact with the surface cations of the metal (Fig. 9.1).

As a result of the interaction, the metal is oxidized and its hydrated ions go into solution, leaving electrons in the metal:

Me (k) + m H 2 Oxidation Me n+ *m H 2 O (p) + ne-

The metal becomes negatively charged, and the solution becomes positively charged. Positively charged ions from the solution are attracted to the negatively charged metal surface (Me). An electric double layer appears at the metal-solution interface (Fig. 9.2). The potential difference arising between the metal and the solution is called electrode potential or redox potential of the electrode φ Ме n + /Ме(φ Ox / Red in general). A metal immersed in a solution of its own salt is an electrode (Section 10.1). The symbol of the metal electrode Me/Me n + reflects the participants in the electrode process.

As ions pass into solution, the negative charge of the metal surface and the positive charge of the solution increase, which prevents oxidation (ionization) of the metal.

In parallel with the oxidation process, a reverse reaction occurs - the reduction of metal ions from solution to atoms (metal precipitation) with the loss of the hydration shell on the metal surface:

Me n+ * m H 2 O(p) + ne- reduction Me(k) + m H 2 O.

As the potential difference between the electrode and the solution increases, the rate of the forward reaction decreases, and the rate of the reverse reaction increases. At a certain value of the electrode potential, the rate of the oxidation process will be equal to the rate of the reduction process, equilibrium is established:

Me n + * m H 2 O (p) + ne - D Me (k) + m H 2 O.

For simplicity, hydration water is usually not included in the reaction equation and it is written as

Ме n + (р) + ne - D Ме (к)

or in general form for any other redox systems:

Ox + ne - D Red.

The potential established under equilibrium conditions of the electrode reaction is called equilibrium electrode potential. In the case considered, the ionization process in solution is thermodynamically possible, and the metal surface becomes negatively charged. For some metals (less active), the process of reduction of hydrated ions to metal is thermodynamically more probable, then their surface is charged positively, and the layer of adjacent electrolyte is charged negatively.

Hydrogen electrode device. The absolute values ​​of electrode potentials cannot be measured; therefore, their relative values ​​are used to characterize electrode processes. To do this, find the potential difference between the measured electrode and the reference electrode, the potential of which is conventionally assumed to be zero. A standard hydrogen electrode, which is classified as a gas electrode, is often used as a reference electrode. In general, gas electrodes consist of a metal conductor in contact simultaneously with a gas and a solution containing an oxidized or reduced form of an element included in the gas. The metal conductor serves to supply and remove electrons and, in addition, is a catalyst for the electrode reaction. The metal conductor should not send its own ions into the solution. Platinum and platinum metals satisfy these conditions.

The hydrogen electrode (Fig. 9.3) is a platinum plate coated with a thin layer of loose porous plate (to increase surface of the electrode) and immersed in an aqueous solution of sulfuric acid with an activity (concentration) of H + ions equal to unity.

Hydrogen is passed through a solution of sulfuric acid under atmospheric pressure. Platinum (Pt) is an inert metal that practically does not interact with solvents or solutions (does not send its ions into the solution), but it is capable of adsorbing molecules, atoms, and ions of other substances. When platinum comes into contact with molecular hydrogen, hydrogen is adsorbed on the platinum. Adsorbed hydrogen, interacting with water molecules, goes into solution in the form of ions, leaving electrons in the platinum. In this case, the platinum is charged negatively, and the solution – positively. A potential difference arises between the platinum and the solution. Along with the transition of ions into solution, the reverse process occurs - the reduction of H + ions from solution with the formation of hydrogen molecules . Equilibrium at the hydrogen electrode can be represented by the equation

2H + + 2e - D H 2.

Symbol for hydrogen electrode H 2 , Pt│H + . The potential of the hydrogen electrode under standard conditions (T = 298 K, P H2 = 101.3 kPa, [H + ] = 1 mol/l, i.e. pH = 0) is assumed to be conditionally equal to zero: j 0 2H + / H2 = 0 V.

Standard electrode potentials . Electrode potentials measured relative to a standard hydrogen electrode under standard conditions(T = 298 K; for dissolved substances, concentration (activity) C Red = C o x = 1 mol/l or for metals C Me n + = 1 mol/l, and for gaseous substances P = 101.3 kPa), are called standard electrode potentials and are designated j 0 O x / Red. These are reference values.

The higher the algebraic value of their standard electrode (redox) potential, the higher the oxidative capacity of substances. On the contrary, the lower the standard electrode potential of the reactant, the more pronounced its reducing properties. For example, comparison of standard system potentials

F 2 (g.) + 2e - D 2F (p.) j 0 = 2.87 V

H 2 (r.)+ 2e - D 2H (r.) j 0 = -2.25 V

shows that F 2 molecules have a strongly pronounced oxidative tendency, while H ions have a reducing tendency.

A range of metal stresses. By arranging metals in a series as the algebraic value of their standard electrode potentials increases, the so-called “Series of standard electrode potentials” or “Series of voltages”, or “Series of activity of metals” is obtained.

The position of the metal in the “Series of Standard Electrode Potentials” characterizes the reducing ability of metal atoms, as well as the oxidizing properties of metal ions in aqueous solutions under standard conditions. The lower the value of the algebraic value of the standard electrode potential, the greater the reducing properties of a given metal in the form of a simple substance, and the weaker the oxidizing properties of its ions and vice versa. .

For example, lithium (Li), which has the lowest standard potential, is one of the strongest reducing agents, while gold (Au), which has the highest standard potential, is a very weak reducing agent and only oxidizes when interacting with very strong oxidizing agents. From the data of the “Voltage Series” it is clear that the ions of lithium (Li +), potassium (K +), calcium (Ca 2+), etc. - the weakest oxidizing agents, and the strongest oxidizing agents include ions of mercury (Hg 2+), silver (Ag +), palladium (Pd 2+), platinum (Pt 2+), gold (Au 3+, Au +).

Nernst equation. Electrode potentials are not constant. They depend on the ratio of concentrations (activities) of the oxidized and reduced forms of the substance, on temperature, the nature of the dissolved substance and solvent, pH of the medium, etc. This dependence is described Nernst equation:

,

where j 0 О x / Red is the standard electrode potential of the process; R – universal gas constant; T – absolute temperature; n is the number of electrons participating in the electrode process; a oh, a Red – activity (concentration) of the oxidized and reduced forms of the substance in the electrode reaction; x and y are stoichiometric coefficients in the electrode reaction equation; F is Faraday's constant.

For the case when the electrodes are metal and the equilibria established on them are described in general form

Me n + + ne - D Me,

Nernst's equation can be simplified by taking into account that for solids the activity is constant and equal to unity. For 298 K, after substituting a Me =1 mol/l, x=y=1 and constant values ​​R=8.314 J/K*mol; F = 96485 C/mol, replacing the activity of a Me n + with the molar concentration of metal ions in a solution of C Me n + and introducing a factor of 2.303 (transition to decimal logarithms), we obtain the Nernst equation in the form

j Ме n + / Ме = j 0 Ме n + / Ме + lg С Ме n + .

General chemistry: textbook / A. V. Zholnin; edited by V. A. Popkova, A. V. Zholnina. - 2012. - 400 pp.: ill.

Chapter 8. REDOX REACTIONS AND PROCESSES

Chapter 8. REDOX REACTIONS AND PROCESSES

Life is a continuous chain of redox processes.

A.-L. Lavoisier

8.1. BIOLOGICAL SIGNIFICANCE OF REDOX PROCESSES

The processes of metabolism, respiration, decay, fermentation, photosynthesis are basically redox processes. In the case of aerobic metabolism, the main oxidizing agent is molecular oxygen, and the reducing agent is organic substances in food products. An indicator that the vital activity of the body is based on redox reactions is the bioelectric potential of organs and tissues. Biopotentials are a qualitative and quantitative characteristic of the direction, depth and intensity of biochemical processes. Therefore, recording the biopotentials of organs and tissues is widely used in clinical practice when studying their activities, in particular, when diagnosing cardiovascular diseases, an electrocardiogram is taken, and when measuring the biopotentials of muscles, an electromyogram is taken. Registration of brain potentials - encephalography - allows us to judge pathological disorders of the nervous system. The source of energy for the vital activity of cells is the membrane potential equal to 80 mV, caused by the occurrence of ion asymmetry, i.e. unequal distribution of cations and anions on both sides of the membrane. The membrane potential is ionic in nature. In multinuclear complexes, processes occur related to the transfer of electrons and protons between particles that resist

are driven by a change in the degree of oxidation of the reacting particles and the appearance of a redox potential. The redox potential is electronic in nature. These processes are reversible and cyclic in nature and underlie many important physiological processes. Michaelis noted the important role of redox processes in life: “The redox processes occurring in living organisms are among those that not only catch the eye and can be identified, but are also the most important for life, both biologically and from a philosophical point of view."

8.2. ESSENCE

REDOX PROCESSES

In 1913 L.V. Pisarzhevsky came up with an electronic theory of redox processes, which is currently generally accepted. This type of reaction is carried out due to the redistribution of electron density between the atoms of the reacting substances (transfer of electrons), which manifests itself in a change in the oxidation state.

Reactions that result in changes in the oxidation states of the atoms that make up the reacting substances due to electron transfer between them are called redox reactions.

The redox process consists of 2 elementary acts or half-reactions: oxidation and reduction.

Oxidation- this is the process of loss (donation) of electrons by an atom, molecule or ion. During oxidation, the oxidation state of particles increases:

A particle that donates electrons is called reducing agent. The oxidation product of a reducing agent is called its oxidized form:

The reducing agent and its oxidized form form one pair of the redox system (Sn 2 + / Sn 4 +).

A measure of the reducing ability of an element is ionization potential. The lower the ionization potential of an element, the stronger the reducing agent it is; s-elements and elements in the lowest and intermediate oxidation states are strong reducing agents. The ability of a particle to donate electrons (donor ability) determines its reducing properties.

Recovery - This is the process of adding electrons to a particle. During reduction, the oxidation state decreases:

The particle (atoms, molecules or ions) that gains electrons is called oxidizing agent The reduction product of the oxidizing agent is called its restored form:

The oxidizing agent with its reduced form constitutes another pair (Fe 3+ /Fe 2+) of the redox system. A measure of the oxidative capacity of particles is electron affinity. The greater the electron affinity, i.e. electron-withdrawing ability of a particle, the more powerful an oxidizing agent it is. Oxidation is always accompanied by reduction, and, conversely, reduction is associated with oxidation.

Let's consider the interaction of FeCl 3 with SnCl 2. The process consists of two half-reactions:

A redox reaction can be represented as a combination of two conjugate pairs.

During reactions, the oxidizing agent is converted into a conjugate reducing agent (reduction product), and the reducing agent is converted into a conjugate oxidizing agent (oxidation product). They are considered as redox couples:

Therefore, redox reactions represent the unity of two opposing processes of oxidation and reduction, which in systems cannot exist one without the other. In this we see a manifestation of the universal law of unity and struggle of opposites. A reaction will occur if the electron affinity of the oxidizing agent is greater than the ionization potential of the reducing agent. For this purpose, the concept was introduced electronegativity - a quantity characterizing the ability of atoms to give or accept electrons.

The equations of redox reactions are drawn up using the electron balance method and the half-reaction method. The half-reaction method should be preferred. Its use is associated with the use of ions that actually exist; the role of the medium is visible. When drawing up equations, it is necessary to find out which of the substances entering the reaction act as an oxidizing agent and which ones act as a reducing agent, the influence of the pH of the medium on the course of the reaction, and what are the possible reaction products. Redox properties are exhibited by compounds that contain atoms with a large number of valence electrons with different energies. Compounds of d-elements (IB, VIIB, VIIIB groups) and p-elements (VIIA, VIA, VA groups) have these properties. Compounds that contain an element in the highest oxidation state exhibit only oxidizing properties(KMnO 4, H 2 SO 4), in the lowest - only restorative properties(H2S), in the intermediate - they can behave in two ways(Na 2 SO 3). After composing the half-reaction equations, the ionic equations create the reaction equation in molecular form:

Checking the correctness of the equation: the number of atoms and charges on the left side of the equation must be equal to the number of atoms and charges on the right side of the equation for each element.

8.3. CONCEPT OF ELECTRODE POTENTIAL. MECHANISM OF ELECTRODE POTENTIAL APPEARANCE. GALVANIC CELL. NERNST EQUATION

A measure of the redox ability of substances is the redox potential. Let us consider the mechanism of potential emergence. When a reactive metal (Zn, Al) is immersed in a solution of its salt, for example Zn in a solution of ZnSO 4, additional dissolution of the metal occurs as a result of the oxidation process, the formation of a pair, a double electrical layer on the surface of the metal and the emergence of a potential of the Zn 2 +/Zn° pair .

A metal immersed in a solution of its salt, for example zinc in a solution of zinc sulfate, is called an electrode of the first kind. This is a two-phase electrode that charges negatively. The potential is formed as a result of an oxidation reaction (according to the first mechanism) (Fig. 8.1). When low-active metals (Cu) are immersed in a solution of their own salt, the opposite process is observed. At the interface of the metal with the salt solution, the metal is deposited as a result of the reduction process of an ion that has a high electron acceptor ability, which is due to the high charge of the nucleus and the small radius of the ion. The electrode becomes positively charged, excess salt anions form a second layer in the near-electrode space, and an electrode potential of the Cu 2 +/Cu° pair arises. The potential is formed as a result of the recovery process according to the second mechanism (Fig. 8.2). The mechanism, magnitude and sign of the electrode potential are determined by the structure of the atoms of the participants in the electrode process.

So, the potential arises at the interface between the metal and the solution as a result of oxidation and reduction processes occurring with the participation of the metal (electrode) and the formation of a double electrical layer is called the electrode potential.

If electrons are transferred from a zinc plate to a copper plate, then the equilibrium on the plates is disrupted. To do this, we connect the zinc and copper plates, immersed in solutions of their salts, with a metal conductor, and the near-electrode solutions with an electrolyte bridge (a tube with a K 2 SO 4 solution) to close the circuit. An oxidation half-reaction occurs on the zinc electrode:

and on copper - the reduction half-reaction:

The electric current is caused by the total redox reaction:

Electric current appears in the circuit. The reason for the occurrence and flow of electric current (EMF) in a galvanic cell is the difference in electrode potentials (E) - fig. 8.3.

Rice. 8.3. Electrical circuit diagram of a galvanic cell

Galvanic cell is a system in which the chemical energy of the redox process is converted

to electric. The chemical circuit of a galvanic cell is usually written in the form of a short diagram, where a more negative electrode is placed on the left, the pair formed on this electrode is indicated with a vertical line, and the potential jump is shown. Two lines indicate the boundary between solutions. The electrode charge is indicated in parentheses: (-) Zn°|Zn 2 +||Cu 2 +|Cu° (+) - diagram of the chemical circuit of a galvanic cell.

The redox potentials of the pair depend on the nature of the participants in the electrode process and the ratio of the equilibrium concentrations of the oxidized and reduced forms of the participants in the electrode process in the solution, the temperature of the solution, and are described by the Nernst equation. A quantitative characteristic of a redox system is the redox potential that occurs at the platinum-aqueous solution interface. The magnitude of the potential in SI units is measured in volts (V) and is calculated by Nernst-Peters equation:

where a(Ox) and a(Red) are the activity of the oxidized and reduced forms, respectively; R- universal gas constant; T- thermodynamic temperature, K; F- Faraday constant (96,500 C/mol); n- the number of electrons taking part in the elementary redox process; a - activity of hydronium ions; m- stoichiometric coefficient before the hydrogen ion in the half-reaction. The value φ° is the standard redox potential, i.e. potential measured under the conditions a(Ox) = a(Red) = a(H +) = 1 and a given temperature.

The standard potential of the 2H + /H 2 system is assumed to be 0 V. Standard potentials are reference values ​​and are tabulated at a temperature of 298K. A strongly acidic environment is not typical for biological systems, therefore, to characterize the processes occurring in living systems, the formal potential is more often used, determined under the condition a(Ox) = a(Red), pH 7.4 and temperature 310K (physiological level). When writing the potential of a pair, it is indicated as a fraction, with the oxidizing agent in the numerator and the reducing agent in the denominator.

For 25 °C (298K) after substituting constant values ​​(R = 8.31 J/mol deg; F= 96,500 C/mol) the Nernst equation takes the following form:

where φ° is the standard redox potential of the pair, V; with o.f. and with v.f. - product of equilibrium concentrations of oxidized and reduced forms, respectively; x and y are stoichiometric coefficients in the half-reaction equation.

The electrode potential is formed on the surface of a metal plate immersed in a solution of its salt and depends only on the concentration of the oxidized form [M n+ ], since the concentration of the reduced form does not change. The dependence of the electrode potential on the concentration of the ion of the same name is determined by the equation:

where [M n+ ] is the equilibrium concentration of the metal ion; n- the number of electrons participating in the half-reaction and corresponds to the oxidation state of the metal ion.

Redox systems are divided into two types:

1) in the system only electron transfer occurs Fe 3 + + ē = = Fe 2 +, Sn 2 + - 2ē = Sn 4 +. This isolated redox equilibrium;

2) systems when the transfer of electrons is complemented by the transfer of protons, i.e. observed combined equilibrium of different types: protolytic (acid-base) and redox with possible competition between two particles of protons and electrons. In biological systems, important redox systems are of this type.

An example of a system of the second type is the process of recycling hydrogen peroxide in the body: H 2 O 2 + 2H + + 2ē ↔ 2H 2 O, as well as the reduction in an acidic environment of many oxidizing agents containing oxygen: CrO 4 2-, Cr 2 O 7 2-, MnO 4 - . For example, MnO 4 - + 8H + + 5ē = = Mn 2 + + 4H 2 O. Electrons and protons participate in this half-reaction. The pair potential is calculated using the formula:

In a wider range of conjugate pairs, the oxidized and reduced forms of the pair are in solution in varying degrees of oxidation (MnO 4 - /Mn 2 +). As a measuring electrode

in this case, an electrode made of inert material (Pt) is used. The electrode is not a participant in the electrode process and only plays the role of an electron carrier. The potential generated due to the redox process occurring in a solution is called redox potential.

It is measured on redox electrode is an inert metal found in a solution containing oxidized and reduced forms of the pair. For example, when measuring E o Fe 3 + /Fe 2 + pairs use a redox electrode - a platinum measuring electrode. The reference electrode is hydrogen, the pair potential of which is known.

Reaction occurring in a galvanic cell:

Chemical chain diagram: (-)Pt|(H 2 °), H+||Fe 3 +, Fe 2 +|Pt(+).

Oxidation-reduction potential is a measure of the redox ability of substances. The values ​​of standard pair potentials are indicated in the reference tables.

The following patterns are noted in the series of redox potentials.

1. If the standard redox potential of a pair is negative, for example φ°(Zn 2+ (p)/Zn°(t)) = -0.76 V, then in relation to the hydrogen pair, whose potential is higher, this pair acts as reducing agent. The potential is formed by the first mechanism (oxidation reaction).

2. If the potential of the pair is positive, for example φ°(Cu 2 +(p)/ Cu(t)) = +0.345 V relative to a hydrogen or other conjugate pair whose potential is lower, this pair is an oxidizing agent. The potential of this pair is formed by the second mechanism (reduction reaction).

3. The higher the algebraic value of the standard potential of the pair, the higher the oxidizing ability of the oxidized form and the lower the reducing ability of the reduced form of this

couples. A decrease in the value of the positive potential and an increase in the negative corresponds to a decrease in oxidative activity and an increase in reduction activity. For example:

8.4. HYDROGEN ELECTRODE, MEASUREMENT OF REDOX POTENTIALS

The redox potential of a pair is determined by the potential of the electrical double layer, but, unfortunately, there is no method for measuring it. Therefore, they determine not the absolute, but the relative value, choosing some other pair for comparison. Potential measurement is carried out using a potentiometric setup, which is based on a galvanic element having a circuit: the electrode of the test pair (measuring electrode) is connected to the electrode of a hydrogen pair (H + /H°) or any other whose potential is known (reference electrode) . The galvanic cell is connected to an amplifier and an electric current meter (Fig. 8.4).

A hydrogen pair is formed at the hydrogen electrode as a result of the redox process: 1/2H 2 o (g) ↔ H + (p) + e - . The hydrogen electrode is a half-cell consisting

from a platinum plate coated with a thin, loose layer of platinum, dipped in a 1 N solution of sulfuric acid. Hydrogen is passed through the solution; in the porous layer of platinum, part of it becomes atomic. All this is enclosed in a glass vessel (ampoule). The hydrogen electrode is a three-phase electrode of the first kind (gas-metal). Analyzing the electrode potential equation for a hydrogen electrode, we can conclude that the potential of the hydrogen electrode increases linearly

Rice. 8.4. Hydrogen electrode

with a decrease in the pH value (increase in acidity) of the medium and a decrease in the partial pressure of hydrogen gas above the solution.

8.5. DIRECTION PREDICTION

BY CHANGE IN THE FREE ENERGY OF SUBSTANCES AND BY THE VALUES OF STANDARD REDOX POTENTIALS

The direction of the redox reaction can be judged by the change in the isobaric-isothermal potential of the system (Gibbs energy) and the free energy (ΔG) of the process. The reaction is fundamentally possible at ΔG o < 0. В окислительно-восстановительной реакции изменение свободной энергии равно электрической работе, совершаемой системой, в результате которой ē переходит от восстановителя к окислителю. Это находит отражение в формуле:

Where F- Faraday constant equal to 96.5 kK/mol; n- the number of electrons involved in the redox process, per 1 mole of substance; E o- the magnitude of the difference between the standard redox potentials of two conjugate pairs of the system, which is called the electromotive force of reactions (EMF). This equation reflects the physical meaning of the relationship E o and Gibbs free energy of the reaction.

For the spontaneous occurrence of a redox reaction, it is necessary that the potential difference of conjugated pairs be a positive value, which follows from the equation, i.e. a pair whose potential is higher can act as an oxidizing agent. The reaction continues until the potentials of both pairs become equal. Therefore, to answer the question whether a given reducing agent will be oxidized by a given oxidizing agent or, conversely, you need to know ΔE o : ΔE o = φ°oxid. - φ°recovery The reaction proceeds in a direction that results in the formation of a weaker oxidizing agent and a weaker reducing agent. Thus, by comparing the potentials of two conjugate pairs, it is possible to fundamentally resolve the issue of the direction of the process.

Task. Is it possible to reduce the Fe 3+ ion with T1+ ions according to the proposed scheme:

ΔE° reaction has a negative value:

The reaction is impossible, since the oxidized form of Fe 3+ of the Fe 3+ / Fe 2 + pair cannot oxidize the T1+ of the T1 3 + / T1 + pair.

If the emf of the reaction is negative, then the reaction proceeds in the opposite direction. The greater the ΔE°, the more intense the reaction.

Task. What is the chemical behavior of FeC1 3 in a solution containing:

a) NaI; b) NaBr?

We compose half-reactions and find potentials for pairs:

A) E reaction 2I - + 2Fe 3 + = I 2 + 2Fe 2 + will be equal to 0.771-0.536 = = 0.235 V, E has a positive meaning. Consequently, the reaction proceeds towards the formation of free iodine and Fe 2+.

b) E° reaction 2Br - + 2Fe 3 + = Br 2 + 2Fe 2 + will be equal to 0.771-1.065 = -0.29 V. Negative value E o shows that ferric chloride will not be oxidized by potassium bromide.

8.6. EQUILIBRIUM CONSTANT

REDOX REACTION

In some cases, it is necessary to know not only the direction and intensity of redox reactions, but also the completeness of the reactions (what percentage of the starting substances are converted into reaction products). For example, in quantitative analysis you can rely only on those reactions that practically proceed 100%. Therefore, before using this or that reaction to solve any problem, determine the constant equal to

news (K R) of a given island system. To determine the Kp of redox processes, use the table of standard redox potentials and the Nernst equation:

because the when equilibrium is reached, the potentials of the conjugate pairs of oxidizer and reducer of the redox process become the same: φ°oxid. - φ°recovery = 0, then E o= 0. From the Nernst equation under equilibrium conditions E o reaction is equal to:

Where n- the number of electrons involved in the redox reaction; P.S. cont. district and P.S. ref. c-c - respectively, the product of the equilibrium concentrations of reaction products and starting substances to the power of their stoichiometric coefficients in the reaction equation.

The equilibrium constant indicates that the equilibrium state of a given reaction occurs when the product of the equilibrium concentrations of the reaction products becomes 10 times greater than the product of the equilibrium concentrations of the starting substances. In addition, a large Kp value indicates that the reaction proceeds from left to right. Knowing Kp, it is possible, without resorting to experimental data, to calculate the completeness of the reaction.

8.7. REDOX REACTIONS IN BIOLOGICAL SYSTEMS

During life, electrical potential differences may arise in cells and tissues. Electrochemical transformations in the body can be divided into 2 main groups.

1. Redox processes due to the transfer of electrons from one molecules to others. These processes are of an electronic nature.

2. Processes associated with the transfer of ions (without changing their charges) and the formation of biopotentials. Biopotentials recorded in the body are mainly membrane potentials. They are ionic in nature. As a result of these processes, potentials arise between different layers of tissues that are in different physiological states. They are associated with different intensities of physiological redox processes. For example, potentials formed in the tissues of the leaf surface on the illuminated and unlit sides as a result of different rates of the photosynthesis process. The illuminated area turns out to be positively charged relative to the unlit area.

In redox processes of an electronic nature, three groups can be distinguished.

The first group includes processes associated with the transfer of electrons between substances without the participation of oxygen and hydrogen. These processes are carried out with the participation of electron transfer complexes - heterovalent and heteronuclear complexes. Electron transfer occurs in complex compounds of the same metal or atoms of different metals, but in different oxidation states. The active source of electron transfer are transition metals, which exhibit several stable oxidation states, and the transfer of electrons and protons does not require large energy costs; transfer can be carried out over long distances. The reversibility of processes allows for repeated participation in cyclic processes. These oscillatory processes are found in enzymatic catalysis (cytochromes), protein synthesis, and metabolic processes. This group of transformations is involved in maintaining antioxidant homeostasis and protecting the body from oxidative stress. They are active regulators of free radical processes, a system for recycling reactive oxygen species and hydrogen peroxide, and participate in the oxidation of substrates

such as catalase, peroxidase, dehydrogenase. These systems carry out antioxidant and antiperoxide effects.

The second group includes redox processes associated with the participation of oxygen and hydrogen. For example, oxidation of the aldehyde group of the substrate into an acidic one:

The third group includes processes associated with the transfer of protons and electrons from the substrate, which are pH-dependent in nature and occur in the presence of dehydrogenase (E) and coenzyme (Co) enzymes with the formation of an activated enzyme-coenzyme-substrate complex (E-Co-S ), adding electrons and hydrogen cations from the substrate, and cause its oxidation. Such a coenzyme is nicotinamide adenine dinucleotide (NAD +), which attaches two electrons and one proton:

In biochemical processes, combined chemical equilibria take place: redox, protolytic, and complexation processes. The processes are usually enzymatic in nature. Types of enzymatic oxidation: dehydrogenase, oxidase (cytochromes, free radical oxidation-reduction). The redox processes occurring in the body can be conditionally divided into the following types: 1) reactions of intramolecular dismutation (disproportionation) due to carbon atoms of the substrate; 2) intermolecular reactions. The presence of carbon atoms in a wide range of oxidation states from -4 to +4 indicates its duality. Therefore, in organic chemistry, redox dismutation reactions due to carbon atoms, which occur intra- and intermolecularly, are common.

8.8. MEMBRANE POTENTIAL

Since the time of R. Virchow it has been known that living cell is an elementary cell of biological organization that provides all the functions of the body. The occurrence of many physiological processes in the body is associated with the transfer of ions in cells and tissues and is accompanied by the appearance of a potential difference. A large role in membrane transport belongs to the passive transport of substances: osmosis,

filtration and bioelectrogenesis. These phenomena are determined by the barrier properties of cell membranes. The potential difference between solutions of different concentrations separated by a selectively permeable membrane is called membrane potential. The membrane potential is ionic rather than electronic in nature. It is caused by the occurrence of ion asymmetry, i.e. unequal distribution of ions on both sides of the membrane.

The cationic composition of the intercellular medium is close to the ionic composition of sea water: sodium, potassium, calcium, magnesium. In the process of evolution, nature created a special way of transporting ions, called passive transport, accompanied by the appearance of a potential difference. In many cases, the basis for the transfer of substances is diffusion, therefore the potential that forms on the cell membrane is sometimes called diffusion potential. It exists until the ion concentration equalizes. The potential value is small (0.1 V). Facilitated diffusion occurs through ion channels. Ionic asymmetry is used to generate excitation in nerve and muscle cells. However, the presence of ionic asymmetry on both sides of the membrane is also important for those cells that are not able to generate an excitatory potential.

8.9. QUESTIONS AND TASKS FOR SELF-TEST

PREPARATION FOR CLASSES

AND EXAMINATIONS

1.Give the concept of electrode and redox potentials.

2.Note the main patterns observed in the series of redox potentials.

3.What is a measure of the reducing ability of substances? Give examples of the most common reducing agents.

4.What is a measure of the oxidizing ability of a substance? Give examples of the most common oxidizing agents.

5. How can you experimentally determine the value of the redox potential?

6. How will the potential of the Co 3+ /Co 2+ system change when cyanide ions are introduced into it? Explain your answer.

7.Give an example of reactions in which hydrogen peroxide plays the role of an oxidizing agent (reducing agent) in acidic and alkaline environments.

8.What is the significance of the phenomenon of identifying the ligand environment of the central atom on the redox potential for the functioning of living systems?

9. The Krebs cycle in the biological oxidation of glucose is immediately preceded by the reaction:

where NADH and NAD + are the reduced and oxidized form of nicotinamide dinucleotide. In what direction does this redox reaction proceed under standard conditions?

10.What are the names of substances that react reversibly with oxidizing agents and protect substrates?

11.Give examples of the action of bactericidal substances based on oxidative properties.

12. Reactions underlying the methods of permanganatometry and iodometry. Working solutions and methods for their preparation.

13.What is the biological role of reactions in which the oxidation state of manganese and molybdenum changes?

14.What is the mechanism of the toxic effect of nitrogen (III), nitrogen (IV), nitrogen (V) compounds?

15.How is superoxide ion neutralized in the body? Give the reaction equation. What is the role of metal ions in this process?

16.What is the biological role of half-reactions: Fe 3+ + ē ↔ Fe 2+ ; Cu 2+ + ē ↔ Cu + ; Co 3+ + ē ↔ Co 2+ ? Give examples.

17. How is the standard EMF related to the change in the Gibbs energy of the redox process?

18.Compare the oxidizing ability of ozone, oxygen and hydrogen peroxide with respect to an aqueous solution of potassium iodide. Support your answer with tabular data.

19.What chemical processes underlie the neutralization of superoxide anion radicals and hydrogen peroxide in the body? Give the half-reaction equations.

20. Give examples of redox processes in living systems, accompanied by changes in the oxidation states of d-elements.

21.Give examples of the use of redox reactions for detoxification.

22.Give examples of the toxic effects of oxidizing agents.

23. The solution contains particles Cr 3+, Cr 2 O 7 2-, I 2, I -. Determine which of them interact spontaneously under standard conditions?

24.Which of these particles is a stronger oxidizing agent in an acidic environment, KMnO 4 or K 2 Cr 2 O 7?

25.How to determine the dissociation constant of a weak electrolyte using the potentiometric method? Draw a diagram of the chemical circuit of a galvanic cell.

26. Is it acceptable to simultaneously introduce solutions of RMnO 4 and NaNO 2 into the body?

8.10. TEST TASKS

1. Which halogen molecules (simple substances) exhibit redox duality?

a) none, they are all only oxidizing agents;

b) everything except fluorine;

c) everything except iodine;

d) all halogens.

2. Which halide ion has the greatest reducing activity?

a)F - ;

b)C1 - ;

c)I - ;

d)Br - .

3. Which halogens undergo disproportionation reactions?

a) everything except fluorine;

b) everything except fluorine, chlorine, bromine;

c) everything except chlorine;

d) none of the halogens are involved.

4. Two test tubes contain solutions of KBr and KI. FeCl 3 solution was added to both test tubes. In what case is the halide ion oxidized to a free halogen if E o (Fe 3+ / Fe 2+) = 0.77 V; E°(Br 2 /2Br -) = 1.06 V; E o (I2/2I -) = 0.54 V?

a) KBr and KI;

b)KI;

c) KBr;

d) not in any case.

5. The most powerful reducing agent:

6. In which of the reactions involving hydrogen peroxide will gaseous oxygen be one of the reaction products?

7. Which of the following elements has the highest relative electronegativity?

a)O;

b)C1;

c)N;

d)S.

8. Carbon in organic compounds exhibits the following properties:

a) oxidizing agent;

b) reducing agent;

Reactions differ intermolecular, intramolecular and auto-oxidation-self-healing (or disproportionation):

If the oxidizing and reducing agents are the elements included in the composition different compounds, then the reaction is called intermolecular.

Example: Na2 S O3+ O 2  Na 2 SO 4

ok ok

If the oxidizing agent and the reducing agent are elements that are part of the same compound, then the reaction is called intramolecular.

Example: ( N H 4) 2 Cr 2 O 7  N 2 + Cr 2 O 3 + H 2 O.

v–l o–l

If the oxidizing and reducing agent is the same element in this case, part of its atoms is oxidized, and the other is reduced, then the reaction is called autoxidation–self-healing.

Example: H 3 P O 3  H 3 P O4+ P H 3

in–l/o–l

This classification of reactions turns out to be convenient in determining potential oxidizing and reducing agents among given substances.

4 Determination of the possibility of redox

reactionsby oxidation states of elements

A necessary condition for the interaction of substances according to the redox type is the presence of a potential oxidizing agent and reducing agent. Their definition was discussed above; now we will show how to apply these properties to analyze the possibility of a redox reaction (for aqueous solutions).

Examples

1) HNO 3 + PbO 2  ... - the reaction does not occur, because No

o–l o–l potential reducing agent;

2) Zn + KI ... - the reaction does not occur, because No

v–l v–l potential oxidizing agent;

3) KNO 2 +KBiO 3 +H 2 SO 4  ...- the reaction is possible if

some KNO 2 will be a reducing agent;

4) KNO 2 + KI +H 2 SO 4  ... - the reaction is possible if

o – l v – l KNO 2 will be an oxidizing agent;

5) KNO 2 + H 2 O 2  ... - the reaction is possible if

c – l o – l H 2 O 2 will be an oxidizing agent, and KNO 2

Reducer (or vice versa);

6) KNO 2  ... - reaction possible

o – l / v – l disproportionation

The presence of potential oxidizing and reducing agents is a necessary but not sufficient condition for the reaction to occur. Thus, in the examples discussed above, only in the fifth can we say that one of two possible reactions will occur; in other cases, additional information is needed: will this reaction energetically favorable.

5 Selecting an oxidizing agent (reducing agent) using tables of electrode potentials. Determination of the preferential direction of redox reactions

Reactions occur spontaneously, as a result of which the Gibbs energy decreases (G ch.r.< 0). Для окислительно–восстановительных реакций G х.р. = - nFE 0 , где Е 0 - разность стандартных электродных потенциалов окислительной и восстановительной систем (E 0 = E 0 ок. – E 0 восст.) , F - число Фарадея (96500 Кулон/моль), n - число электронов, участвующих в элементарной реакции; E часто называют ЭДС реакции. Очевидно, что G 0 х.р. < 0, если E 0 х.р. >0.

is it a combination of two

half-reactions:

Zn → Zn 2+ and Cu 2+ → Cu;

the first of them, including reducing agent(Zn) and its oxidized form (Zn 2+) is called restorative system, the second, including oxidizer(Cu 2+) and its reduced form (Cu), - oxidative system.

Each of these half-reactions is characterized by the value of the electrode potential, which are denoted, respectively,

E restore = E 0 Zn 2+ / Zn and E approx. = E 0 Cu 2+ / Cu .

Standard values ​​of E 0 are given in reference books:

E 0 Zn 2+ / Zn = - 0.77 V, E 0 Cu 2+ / Cu = + 0.34 V.

EMF =.E 0 = E 0 approx. – E 0 restore = E 0 Cu 2+ / Cu - E 0 Zn 2+ / Zn = 0.34 – (–0.77) = 1.1 V.

It is obvious that E 0 > 0 (and, accordingly, G 0< 0), если E 0 ок. >E 0 restore , i.e. The redox reaction proceeds in the direction for which the electrode potential of the oxidizing system is greater than the electrode potential of the reducing system.

Using this criterion, you can determine which reaction, direct or reverse, occurs predominantly, as well as choose an oxidizing agent (or reducing agent) for a given substance.

In the example discussed above, E 0 is approx. > E 0 restore Therefore, under standard conditions, copper ions can be reduced by metallic zinc (which corresponds to the position of these metals in the electrochemical series)

Examples

1. Determine whether it is possible to oxidize iodide ions with Fe 3+ ions.

Solution:

a) write a diagram of a possible reaction: I – + Fe 3+  I 2 + Fe 2+,

v–l o–l

b) let’s write the half-reactions for the oxidation and reduction systems and the corresponding electrode potentials:

Fe 3+ + 2e –  Fe 2+ E 0 = + 0.77 B - oxidizing system,

2I –  I 2 + 2e – E 0 = + 0.54 B - reduction system;

c) having compared the potentials of these systems, we will conclude that the given reaction is possible (under standard conditions).

2. Select oxidizing agents (at least three) for a given transformation of a substance and choose from them the one in which the reaction proceeds most completely: Cr(OH) 3  CrO 4 2 – .

Solution:

a) find in the reference book E 0 CrO 4 2 – / Cr (OH)3 = - 0.13 V,

b) using a reference book, we will select suitable oxidizing agents (their potentials should be greater than - 0.13 V), while focusing on the most typical, “non-deficient” oxidizing agents (halogens - simple substances, hydrogen peroxide, potassium permanganate, etc. ).

It turns out that if the transformation Br 2 → 2Br – corresponds to one potential E 0 = +1.1 V, then for permanganate ions and hydrogen peroxide the following options are possible: E 0 MnO 4 – / Mn 2+ = + 1.51 V - V sour environment,

E 0 MnO 4 – / MnO 2 = + 0.60 V - in neutral environment,

E 0 MnO 4 – / MnO 4 2 – = + 0.56 V - in alkaline environment,

E 0 H 2 O 2 / H 2 O = + 1.77 B - in sour environment,

E 0 H 2 O 2/ OH – = + 0.88 B - in alkaline environment.

Considering that the chromium hydroxide specified by the condition is amphoteric and therefore exists only in a weakly alkaline or neutral environment, the following oxidants are suitable:

E 0 MnO4 – /MnO2 = + 0.60 V and. E 0 Br2 /Br – = + 1.1 B..

c) the last condition, the choice of the optimal oxidizer from several, is solved on the basis that the reaction proceeds more completely, the more negative G 0 is for it, which in turn is determined by the value E 0:

The larger algebraically the quantity E 0 , especially the redox reaction proceeds fully, the greater the yield of products.

Of the oxidizing agents discussed above, E 0 will be the largest for bromine (Br 2).

(OB) AND OB – ELECTRODES.

Depending on the oxidation-reduction mechanism, various OM systems can be divided into two types:

1st type: OM – systems in which the redox process is associated with the transfer of only electrons, for example: Fe³ + +ē ↔ Fe² +

2nd type: OB systems in which the redox process is associated not only with the transfer of electrons, but also of protons, for example:

C 6 H 4 O 2 + 2H + +2ē ↔ C 6 H 4 (OH) 2

quinone hydroquinone

MnO 4 - + 8H + + 5ē ↔ Mn² + + 4H 2 O

An inert metal in combination with an OM system is called an oxidation-reduction or redox electrode, and the potential that arises at this electrode is called oxidation-reduction (OR) or redox potential.

The inert metal takes only an indirect part in the potential-determining reaction, being an intermediary in the transfer of electrons from the reduced form of the substance Red to the oxidized OX.

When an inert metal is immersed in a solution containing an excess of the oxidized form of iron, the metal plate becomes positively charged (Fig. 10 a)

With an excess of the reduced form of iron, the platinum surface becomes negatively charged (Fig. 10 b).

Rice. 10. Emergence of OB potential

The transfer of electrons from one ion to another through the metal leads to the formation of an DES on the metal surface.

Inter-electron exchange is possible without metal. But Fe²+ and Fe³+ ions are solvated in different ways and for electron transfer it is necessary to overcome an energy barrier. The transition of electrons from Fe²+ ions to the metal and from the metal surface to the Fe³+ ion is characterized by a lower activation energy.

If the activities of Fe²+ and Fe³+ ions are equal, the platinum plate is charged positively, because The electron-acceptor capacity of Fe³+ ions is greater than the electron-donor capacity of Fe²+.

Peters equation.

The quantitative dependence of the OM - potential on the nature of the OM - system (φ°r), the ratio of activities of the oxidized and reduced forms, temperature, and on the activity of hydrogen ions is established by the Peters equation.

1st type: φr = φ°r + ∙ ln

2nd type: φr = φ°r + ∙ ln

where φr - OB - potential, V;

φ°r - standard OB - potential, V;

z is the number of electrons participating in the OB process;

a (Ox) – activity of the oxidized form, mol/l;

a (Red) – activity of the reducing form, mol/l;

m is the number of protons;

a(n +) – activity of hydrogen ions, mol/l.

The standard OB potential is the potential that arises at the interface between an inert metal and a solution, in which the activity of the oxidized form is equal to the activity of the reduced form, and for a system of the second type, in addition, the activity of hydrogen ions is equal to unity.

Classification of reversible electrodes.

Having examined the principle of operation of electrodes, we can conclude that according to the properties of the substances involved in potential-determining processes, as well as according to their design, all reversible electrodes are divided into the following groups:

Electrodes of the first kind;

Electrodes of the second kind;

Ion selective electrodes;

Redox electrodes.

1. A galvanic cell is a system that produces work rather than consuming it, therefore it is advisable to consider the EMF of the element as a positive value.

2. The EMF of the element is calculated by subtracting the numerical value of the potential of the left electrode from the numerical value of the potential of the right electrode - the “right plus” rule. Therefore, the element circuit is written so that the left electrode is negative and the right electrode is positive.

3. The interface between the conductors of the first and second rows is indicated by one line: Zn׀ZnSO4; Cu׀CuSO4

4. The interface between conductors of the second type is depicted with a dotted line: ZnSO4 (p) ׃ CuSO4 (p)

5. If an electrolyte bridge is used at the interface between two conductors of the second type, it is designated by two lines: ZnSO4 (р) ׀׀ CuSO4 (р).

6. Components of one phase are written separated by commas:

Pt|Fe³+, Fe²+ ; Pt, H2 |HCl(p)

7. The electrode reaction equation is written so that substances in the oxidizing form are located on the left, and in the reducing form on the right.