EntranceRedox systems and their types. Redox processes and redox systems in wine - oxidation and reduction processes in wines
A distinctive feature of redox reactions is the transfer of electrons between reacting particles - ions, atoms, molecules and complexes, as a result of which the oxidation state of these particles changes, for example
Fe2+ ? e? = Fe3+.
Since electrons cannot accumulate in a solution, two processes must occur simultaneously - losses and acquisitions, that is, the process of oxidation of some particles and reduction of other particles. Thus, any redox reaction can always be represented in the form of two half-reactions:
aOx1 + bRed2 = aRed1 + bOx2
The parent particle and the product of each half-reaction constitute a redox couple or system. In the above half-reactions, Red1 is conjugated to Ox1 and Ox2 is conjugated to Red1.
Not only particles in solution, but also electrodes can act as electron donors or acceptors. In this case, the redox reaction occurs at the electrode-solution interface and is called electrochemical.
Redox reactions, like everyone else chemical reactions, are reversible to one degree or another. The direction of reactions is determined by the ratio of the electron-donor 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 equilibrium shift are constant). The movement of electrons during a redox reaction creates a potential. Thus, potential, measured in volts, serves as a measure of the redox ability of a compound.
To quantitatively assess the oxidative (reductive) properties of the system, an electrode made of a chemically inert material is immersed in the solution. At the phase interface, an electron exchange process occurs, leading to the emergence of a potential, which is a function of the activity of electrons in the solution. The higher the oxidizing capacity of the solution, the greater the potential value.
The absolute value of a system's potential cannot be measured. However, if you choose one of the redox systems as a standard one, then relative to it it becomes possible to measure the potential of any other redox system. restoration systems s regardless of the selected indifferent electrode. The H+/H2 system, the potential of which is assumed to be zero, is chosen as the standard one.
Rice. 1.
1. Platinum electrode.
2. Supply of hydrogen gas.
3. An acid solution (usually HCl) in which the concentration of H+ = 1 mol/l.
4. A water seal that prevents the ingress of oxygen from the air.
5. Electrolytic bridge (consisting of concentrated solution KCl), allowing you to attach the second half of the galvanic cell.
The potential of any redox system, measured under standard conditions relative to a hydrogen electrode, is called the standard potential (E0) of this system. The standard potential is considered positive if the system acts as an oxidizing agent and an oxidation half-reaction occurs at the hydrogen electrode:
or negative if the system plays the role of a reducing agent, and a reduction half-reaction occurs at the hydrogen electrode:
The absolute value of the standard potential characterizes the “strength” of the oxidizing agent or reducing agent.
The standard potential - a thermodynamic standardized value - is a very important physicochemical and analytical parameter that allows one to evaluate the direction of the corresponding reaction and calculate the activities of reacting particles under equilibrium conditions.
To characterize a redox system under specific conditions, the concept of real (formal) potential E0 is used, which corresponds to the potential established at the electrode in a given specific solution when the initial concentrations of the oxidized and reduced forms of potential-determining ions are equal to 1 mol/l and the fixed concentration of all other components solution.
From an analytical point of view, real potentials are more valuable than standard potentials, since the true behavior of the system is determined not by the standard, but by the real potential, and it is the latter that allows one to predict the occurrence of a redox reaction under specific conditions. The actual potential of the system depends on the acidity, the presence of foreign ions in the solution and can vary over a wide range.
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 matter 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 having big number valence electrons having 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 immersed chemically active metal(Zn, Al) into a solution of its salt, for example Zn into 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 electric 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 cause of occurrence and progression 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 chemical energy redox process turns
to electric. The chemical chain of a galvanic cell is usually written as brief outline, where a more negative electrode is placed on the left, indicate the pair formed on this electrode with a vertical line, showing the potential jump. 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. Quantitative characteristics redox system is the redox potential that arises at the interface between the platinum and aqueous solution phases. 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. - the product of the equilibrium concentrations of the 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 potential of a pair 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 installation, 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 only rely 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) given o-v systems. 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 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. Kinds enzymatic oxidation: dehydrogenase, oxidase (cytochromes, free radical oxidation-reduction). The redox processes occurring in the body can be divided into 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 a unit cell biological organization, providing all body functions. 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. Big role in membrane transport belongs to the passive transport of substances: osmosis,
filtration and bioelectrogenesis. These phenomena are determined by the barrier properties 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 it? biological role 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 a 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 proposed elements has highest value relative electronegativity?
a)O;
b)C1;
c)N;
d)S.
8. Carbon in organic compounds exhibits 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. = - nFE 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 given by condition Chromium hydroxide is amphoteric and therefore exists only in a weakly alkaline or neutral environment; of the selected oxidizing agents, the following 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).
Rental block
Redox reactions are reactions that occur with a change in the oxidation state of two or more substances.
Oxidation state- this is the conventional charge on the atom, if we assume that the molecule is created according to the ionic mechanism (or - this is the number of electrons received or given away).
Restorers– atoms, molecules, ions – donating electrons.
Oxidizing agents- atoms, molecules, ions - accepting electrons.
Reducing agents participate in the oxidation process, increasing their oxidation state.
Oxidizing agents - participate in the reduction process, lowering their oxidation state.
Types of redox reactions
1. Intermolecular - reactions in which oxidizing and reducing atoms are found in molecules of different substances, for example:
H2S + Cl2 → S + 2HCl
2. Intramolecular- reactions in which oxidizing and reducing atoms are found in molecules of the same substance, for example:
2H2O → 2H2 + O2
3. Disproportionation(auto-oxidation-self-healing) - reactions in which the same element acts both as an oxidizing agent and as a reducing agent, for example:
Cl2 + H2O → HClO + HCl
4. Reproportionation (proportionation, counter-disproportionation) - reactions in which one oxidation state is obtained from two different oxidation states of the same element:
Types of redox reactions in the human body.
Dehydrogenation reaction: SH2 + HAD+= S + HADH+H+
Electron loss: O20 + 1eO2-
Transfer of 2H+ from the reduced substrate to molecular oxygen: SH2 + O20 +2e= S + H2O
Addition of oxygen to the substrate: SH2 + 1/2O20 +2e= HO - S -H
The mechanism of occurrence of electrode and redox potentials. Nernst-Peters equations.
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 ZnSO4, 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 Zn2+/Zn° pair potential.
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 the oxidation reaction (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 Cu2+/Cu° pair arises. The potential is formed as a result of the reduction process (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 that arises at the interface between a metal and a 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 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 (tube with a K2SO4 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 into electrical energy. 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°|Zn2+||Cu2+|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 redox potential, arising at the interface between the platinum and aqueous solution phases. 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+/H2 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; so.fyu and sv.f. - the product of the equilibrium concentrations of the 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 [Mn+], 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 [Mn+] 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 Fe3+ + ē = = Fe2+, Sn2+ - 2ē = Sn4+. This isolated redox equilibrium;
2) systems when electron transfer is complemented by proton transfer, 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 utilization of hydrogen peroxide in the body: H2O2 + 2H+ + 2ē ↔ 2H2O, as well as the reduction in an acidic environment of many oxidizing agents containing oxygen: CrO42-, Cr2O72-, MnO4-. For example, MnО4- + 8Н+ + 5ē = = Mn2+ + 4Н2О. This half-reaction involves electrons and protons. The potential of a pair 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 (MnO4-/Mn2+). 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 onredox electrode is an inert metal found in a solution containing oxidized and reduced forms of the pair. For example, when measuring Eo Fe3+/Fe2+ couples 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|(H2°), H+||Fe3+, Fe2+|Pt(+).
So, oxidation-reduction potential (ORP) is the potential of a system in which the activities of the oxidative and reducing forms of a given substance are equal to one. ORP is measured using redox electrodes in combination with standard reference electrodes.
Each redox reaction has its own redox couple– this pair has the substance in oxidized and reduced form (Fe+3/Fe+2).
A quantitative measure of the activity of a redox pair is the value of its ORP.
ORP vapor>>>oxidizer
ORP pairs<<<восстановитель
ORP depends on:
The nature of the redox couple,
Concentrations
Temperatures
Comparative strength of oxidizing agents and reducing agents. Predicting the direction of redox processes based on the values of redox potentials.
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 standard potentials of electrodes (E°), acting as reducing agents in relation to hydrogen, have a “-” sign, and the “+” sign have the standard potentials of electrodes that are oxidizing agents.
Metals arranged in increasing order of their standard electrode potentials form the so-called electrochemical series of metal voltages: Li, Rb, K, Ba, Sr, Ca, Na, Mg, Al, Mn, Zn, Cr, Fe, Cd, Co, Ni, Sn, Pb, H, Sb, Bi, Cu, Hg, Ag, Pd, Pt, Au.
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°(Zn2+(p)/Zn°(t)) = -0.76 V, then in relation to a hydrogen pair, whose potential is higher, this pair acts as a reducing agent. The potential is formed by the first mechanism (oxidation reaction).
2. If the potential of the pair is positive, for example φ°(Cu2+(p)/ Cu(t)) = +0.345 V relative to a hydrogen or other conjugate pair, the potential of which 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 a pair, the higher the oxidizing ability of the oxidized form and the lower the reducing ability of the reduced form of this pair. 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:
Comparison of the values of standard redox potentials allows us to answer the question: does this or that redox reaction occur?
The difference between the standard oxidation potentials of the oxidized and reduced half-pairs is called electromotive force (EMF).
E0 = Eok-Evost
A quantitative criterion for assessing the possibility of a particular redox reaction occurring is the positive value of the difference between the standard redox potentials of the oxidation and reduction half-reactions.
To establish the possibility of spontaneous occurrence of OVR under standard conditions, it is necessary:
G0298= - p F E0
E>0 G< 0 - самопроизвольно
E< 0 G>0 - back
E = 0 G = 0 - chemical equilibrium
Physicochemical principles of electron transport in the electron transport chain of mitochondria.
All types of redox processes occur during the oxidation of substrates in mitochondria, on the inner membranes of which are located ensembles of enzymes - dehydrogenases, coenzymes (NAD +, FAD, UBC), a series of cytochromes b, c1, c and the enzyme - cytochrome oxidase. They form the cellular respiratory chain system, through which protons and electrons are relayed from the substrate to oxygen molecules delivered by hemoglobin to the cell.
Each component of the respiratory chain is characterized by a certain value of redox potential. The movement of electrons along the respiratory chain occurs stepwise from substances with a low potential (-0.32 V) to substances with a higher potential (+0.82 V), since any compound can donate electrons only to a compound with a higher redox potential (table 1).
Table 1
Standard redox potentials of biomolecules of the respiratory chain
SYSTEM
HALF-REACTION
REDOX POTENTIAL, V
NAD+/NAD×H
NAD+ + H+ + 2 ē → NAD×H
FAD/FAD×H2
FAD+ + 2H+ + 2 ē → FAD×N2
UBH/ UBH×N2
UBH+ 2H+ + 2 ē → UBH×N2
cytochrome b
cytochrome c1
cytochrome c
cytochrome a + a3
О2 + 4 Н+ + 4 ē → 2 Н2О
The tissue respiration chain can be represented as a diagram:
As a result of biological oxidation (dehydrogenation), two hydrogen atoms (in the form of two protons and two electrons) from the substrate enter the respiratory chain. First, a relay race of a proton and a pair of electrons occurs to the NAD+ molecule, which turns into the reduced form of NAD × H, then the flavin base system (FAD/FAD × H2 or FMN/FMN × H2), the next acceptor of two protons and two electrons is ubiquinone (UBQ). Next, only electrons are transferred: two electrons from UBH × H2 is sequentially taken over by cytochromes in accordance with the values of their redox potentials (Table 1). The last of the components, cytochrome oxidase, transfers electrons directly to the oxygen molecule. Reduced oxygen with two protons derived from UBH × H2 forms a water molecule.
1/2 O2 + 2H+ + 2 ē → H2O
It should be noted that each oxygen molecule interacts with two electron transport chains, since in the structure of cytochromes only one-electron transfer Fe3+ → Fe2+ is possible.
Chemistry of complex compounds Types of redox (redox) reactions in the human body. Redox reactions are reactions that occur with a change in the oxidation state of two or more substances.
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(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.