In the formation of the chemical properties of soils, redox processes occupy one of the leading places. The most important factors determining the redox state of soil horizons are the oxygen of soil air and soil solutions, oxide and ferrous compounds of iron, manganese, nitrogen, sulfur, organic matter, and microorganisms.
Oxidation and reduction reactions always proceed simultaneously. The oxidation of one substance participating in the reaction is accompanied by the reduction of another substance.
Redox processes are understood as processes in which, as a possible stage, the transition of electrons from one particle of a substance to another is included. Oxidation is a reaction in which oxygen is added to a substance or a substance loses hydrogen or electrons. Recovery is the loss of oxygen by a substance, the addition of hydrogen or electrons to a substance.
The ability of a soil to undergo redox reactions is measured by the redox potential (ORP).
The redox potential with respect to hydrogen is called Eh. This value depends on the concentration and ratio of oxidizing agents and reducing agents formed in the process of soil formation. Due to the existence of certain redox systems in soil horizons, it is possible to determine the potential difference (Eh) in millivolts using a pair of electrodes immersed in the soil. The values of Eh in different types of soils and soil horizons vary within 100–800 mV, and sometimes have negative values. The value of Eh significantly depends on the acid-base conditions of the environment, vegetation and microorganisms.
Under soil conditions, a significant part of the components involved in redox reactions is represented by solid phases. In reactions involving solid phases, the soil will exhibit high buffering capacity until these components react. Buffering capacity is the ability of the soil to resist changes in ORP under any external influences. This concept characterizes the stability of the redox systems of the soil under natural dynamic conditions and can be called dynamic buffering. In a natural setting, humic substances and iron hydroxide minerals react at low rates.
Soils contain a large set of redox systems: Fe3+ - Fe2+, Mn2+ - Mn3+ - Mn4+, Cu+ - Cu2+, Co2+ - Co3+, NO3‾ - NO2‾ - NH3‾, S6‾ - S2‾.
There are reversible and irreversible redox systems. Reversible are such systems that in the process of changing the redox regime do not change the total stock of components. Irreversible systems in the process of changing the redox regime lose some of the substances. These substances pass into a gaseous state or precipitate. As a rule, irreversible systems predominate in soils.
Reversible redox systems include:
Fe3+ ⇆Fe2+ system. This system occupies a special place among reversible systems. It is sensitive to the slightest changes in the redox environment. The solubility of ferric compounds is extremely low. The migration of iron compounds is possible mainly in the form of ferrous compounds under conditions of high acidity and low Eh.
Mn2+ ⇆ Mn4+ system. This system is extremely sensitive to changes in ORP. Compounds of tetravalent manganese are insoluble under conditions typical of soil horizons. Exchangeable manganese is divalent. The concentration of divalent manganese ions with increasing acidity and decreasing Eh increases by tens of thousands of times. The migration of manganese compounds in the course of soil-forming processes in the vertical and horizontal directions is similar to the migration of iron compounds.
Irreversible redox systems include:
The system NO3 → NO2 → NO → N. The process of nitrification and accumulation of nitrates occurs under conditions of an oxidizing regime and at high Eh 400-500 mV. Soil moisture reduces Eh and promotes the development of denitrification processes.
System sulfates ⇆ sulfides. This redox system plays an important role in all soils where sulfate salts are present. With the participation of microorganisms, the sulfate-sulfide system in the presence of organic matter and a lack of oxygen shifts towards sulfides. There is a process of reduction of sulfates to sulfurous metals:
Na2SO4 + 2C = Na2S + CO2
Under the influence of carbon dioxide present in the soil, sulfurous metals easily decompose and form bicarbonates and carbonates of alkali and alkaline earth metals. In this case, the process of reduction of sulfates occurs:
Na2S + H2CO3 = Na2CO3 + H2S
However, the content of elements with variable valence in the soil solution is quite low. Therefore, the soil solution has a low OM capacity and buffer capacity, and the value of Eh is unstable.
Oxygen dissolved in the soil solution, soil microflora, and water have a more significant effect on OM processes in soils.
Almost all soil reactions occur in an aquatic environment, and water itself can act both as an oxidizing agent and as a reducing agent.
According to the features of the course of redox processes, three series of soils are distinguished: 1) automorphic soils with a predominance of an oxidizing environment, 2) soils with a reducing gley environment, 3) soils with a reducing hydrogen sulfide environment.
The transformations of plant residues, the accumulation and composition of the formed organic substances, and, as a result, the formation of the soil profile are closely related to the OM processes.
such a process of interaction between two substances, in which a reversible oxidation reaction of one substance occurs due to the reduction of another and a mixture of oxidized and reduced ions is formed in the medium, for example. - Fe"" and Fe", Sn" and Sn"", etc. The intensity level of the redox system is determined by the value of the redox potential Eh, which is expressed in volts, in relation to the potential of a normal hydrogen electrode.
The more positive the potential of the system, the more oxidizing properties it has. Potentials that are obtained in systems containing equal concentrations of oxidized and reduced ions, called. normal.
O. o.-v. from. according to the magnitude of normal potentials, they can be arranged in a row, and each system is an oxidizing agent in relation to a system with a more negative normal potential, and a reducing agent in relation to a system with a more positive normal potential. Redox systems play an important role in mineral formation, transformation of organic matter in sedimentary rocks, etc.
Substance equivalent or Equivalent is a real or conditional particle that can attach, release, or otherwise be equivalent to a hydrogen cation in ion exchange reactions or an electron in redox reactions.
For example, in react:
NaOH + HCl \u003d NaCl + H 2 O
the equivalent will be a real particle - Na + ion, in the reaction
the imaginary particle ½Zn(OH) 2 will be the equivalent.
Substance equivalent is also often used to mean number of substance equivalents or equivalent amount of substance- the number of moles of a substance equivalent to one mole of hydrogen cations in the reaction under consideration.
[edit] Equivalent mass
Equivalent mass is the mass of one equivalent of the given substance.
[edit] Equivalent molar mass of a substance
Molar mass equivalents are usually denoted as or . The ratio of the equivalent molar mass of a substance to its own molar mass is called equivalence factor(usually denoted as ).
The molar mass of the equivalents of a substance is the mass of one mole of equivalents, equal to the product of the equivalence factor by the molar mass of this substance.
M eq = f eq ×M
[edit] Equivalence factor
The ratio of the equivalent molar mass to its own molar mass is called equivalence factor(usually denoted as ).
[edit] Equivalence number
Equivalence number z is a small positive integer equal to the number of equivalents of some substance contained in 1 mole of this substance. The equivalence factor is related to the equivalence number z the following relation: =1/z.
For example, in react:
Zn(OH) 2 + 2HCl = ZnCl 2 + 2H 2 O
The equivalent is the particle ½Zn(OH) 2 . The number ½ is equivalence factor, z in this case is 2
* - for inert gases Z = 1
The equivalence factor helps to formulate the law of equivalence.
[edit] The law of equivalents
As a result of the work of I. V. Richter (1792-1800), the law of equivalents was discovered:
§ All substances react in equivalent ratios.
§ formula expressing the Law of Equivalents: m 1 E 2 \u003d m 2 E 1
§ Electrochemical equivalent- the amount of substance that should be released on the electrode, according to Faraday's law, when a unit of electricity passes through the electrolyte:
§ where is the Faraday constant.
§ Faraday constant, is a physical constant that determines the relationship between the electrochemical and physical properties of a substance.
§ Faraday's constant is C mol −1 .
§ The Faraday constant is included as a constant in Faraday's second law(the law of electrolysis).
§ Numerically, the Faraday constant is equal to the electric charge, during the passage of which through the electrolyte on the electrode, (1 / z) mol of substance A is released in the formula:
where: 
is the number of electrons involved in the reaction.
§ For the Faraday constant, the following relation is true:
§ where is the elementary charge, and is the Avogadro number.
Isotopes(from other Greek ισος - "equal", "same", and τόπος - "a place") - varieties of atoms (and nuclei) of the same chemical element with a different number of neutrons in the nucleus. The name is due to the fact that the isotopes are in the same place (in the same cell) of the periodic table. The chemical properties of an atom depend practically only on the structure of the electron shell, which, in turn, is determined mainly by the charge of the nucleus Z(that is, the number of protons in it) and almost does not depend on its mass number A(that is, the total number of protons Z and neutrons N). All isotopes of the same element have the same nuclear charge, differing only in the number of neutrons. Usually an isotope is denoted by the symbol of the chemical element to which it belongs, with the addition of an upper left index indicating the mass number (for example, 12 C, 222 Rn). You can also write the name of the element with a hyphenated mass number (for example, carbon-12, radon-222). Some isotopes have traditional proper names (for example, deuterium, actinon).
An example of isotopes: 16 8 O, 17 8 O, 18 8 O - three stable isotopes of oxygen.
[edit] Terminology
The main position of IUPAC is that the correct singular term for atoms (or nuclei) of the same chemical element with the same atomic mass is nuclide, and the term isotopes can be used to designate a set of nuclides of one element. Term isotopes was proposed and used initially in the plural, since at least two types of atoms are needed for comparison. In the future, the use of the term in the singular became widely used in practice - isotope. In addition, the term in the plural is often used to refer to any set of nuclides, and not just one element, which is also incorrect. At present, the positions of international scientific organizations have not been brought to uniformity and the term isotope continues to be widely used, including in the official materials of various divisions of IUPAC and IUPAP. This is one of the examples of how the meaning of the term, originally embedded in it, ceases to correspond to the concept for which this term is used (another textbook example is the atom, which, contrary to the name, is not indivisible).
[edit]History of the discovery of isotopes
The first evidence that substances having the same chemical behavior can have different physical properties came from the study of radioactive transformations of atoms of heavy elements. In 1906-07, it became clear that the product of the radioactive decay of uranium, ionium, and the product of the radioactive disintegrator, radiothorium, have the same chemical properties as thorium, but differ from it in atomic mass and the characteristics of radioactive decay. It was later found that all three products have the same optical and X-ray spectra. Such substances, identical in chemical properties, but different in the mass of atoms and some physical properties, at the suggestion of the English scientist F. Soddy, began to be called isotopes.
[edit] Isotopes in nature
It is believed that the isotopic composition of elements on Earth is the same in all materials. Some physical processes in nature lead to a violation of the isotopic composition of elements (natural fractionation isotopes characteristic of light elements, as well as isotopic shifts during the decay of natural long-lived isotopes). Gradual accumulation in minerals of nuclei - decay products of some long-lived nuclides is used in nuclear geochronology.
[edit]Human uses of isotopes
In technological activities, people have learned to change the isotopic composition of elements to obtain any specific properties of materials. For example, 235 U is capable of a thermal neutron fission chain reaction and can be used as fuel for nuclear reactors or nuclear weapons. However, natural uranium contains only 0.72% of this nuclide, while a chain reaction is practically feasible only if the 235 U content is at least 3%. Due to the closeness of the physicochemical properties of isotopes of heavy elements, the procedure for isotope enrichment of uranium is an extremely complex technological task, which is accessible only to a dozen countries in the world. In many branches of science and technology (for example, in radioimmunoassay), isotope labels are used.
Dissociation constant- a kind of equilibrium constant that indicates the tendency of a large object to dissociate (separate) in a reversible way into small objects, such as when a complex breaks down into its constituent molecules, or when a salt separates into ions in an aqueous solution. The dissociation constant is usually denoted Kd and inverse to the association constant. In the case of salts, the dissociation constant is sometimes called the ionization constant.
In a general reaction
where is the complex A x B y breaks down into x units A and y units B, the dissociation constant is defined as follows:
where [A], [B] and are the concentrations of A, B and the complex A x B y, respectively.
[edit] Definition
The electrolytic dissociation of weak electrolytes, according to the Arrhenius theory, is a reversible reaction, that is, it can be schematically represented by the equations (for monovalent ions:):
KA ↔ K + + A - ,
§ KA - undissociated compound;
§ K + - cation;
§ A − - anion.
The equilibrium constant of such a reaction can be expressed by the equation:
 ,
| (1) |
§ - concentration of undissociated compound in solution;
§ - concentration of cations in solution;
§ - concentration of anions in solution.
The equilibrium constant in relation to the dissociation reaction is called dissociation constant.
[edit] Dissociation of electrolytes with polyvalent ions
In the case of dissociation of electrolytes with multivalent ions, dissociation occurs in steps, and each step has its own value of the dissociation constant.
Example: Dissociation of a polybasic (boric) acid [ source not specified 332 days] :
Stage I: H 3 BO 3 ↔ H + + H 2 BO 3 -,
Stage II: H 2 BO 3 - ↔ H + + HBO 3 2 - ,
Stage III: HBO 3 2− ↔ H + + BO 3 3− ,
The first degree of dissociation for such electrolytes is always much greater than the subsequent ones, which means that the dissociation of such compounds proceeds mainly through the first stage.
[edit] Relationship between dissociation constant and degree of dissociation
Based on the definition of the degree of dissociation, for the KA electrolyte in the dissociation reaction = = α·c, = c - α·c = c·(1 - α), where α is the degree of dissociation of the electrolyte.
| , | (2) |
This expression is called the Ostwald dilution law. For very small α (α<<1) K=cα² и
thus, with an increase in the electrolyte concentration, the degree of dissociation decreases, and with a decrease, it increases. The relationship between the dissociation constant and the degree of dissociation is described in more detail in the article Ostwald's Dilution Law.
[edit] The difference between the experimental results and the Arrhenius model, the derivation of the dissociation constant through activities
The above calculations are based on the Arrhenius theory, which is too rough and does not take into account the factors of the electrostatic interaction of ions. Deviations from the ideal state in electrolyte solutions occur at very low concentrations, since the interionic forces are inversely proportional to square distances between ion centers, while intermolecular forces are inversely proportional seventh degree distances, that is, interionic forces, even in dilute solutions, turn out to be much greater than intermolecular ones.
Lewis showed that simple equations can be preserved for real solutions (see above) if instead of ion concentrations we introduce its function, the so-called activity. Activity (a) is related to concentration (c) through a correction factor γ called the activity factor:
a = γ c
Thus, the expression for the equilibrium constant, according to Arrhenius described by equation (1), according to Lewis will look like:
§ 
;
§ 
;
In the Lewis theory, the relationship between the constant and the degree of dissociation (in the Arrhenius theory written by equation (2) is expressed by the relationship:
If there are no other influences that deviate the solution from the ideal state, then the non-dissociated molecules behave like ideal gases and γ KA = 1, and the true expression of the Ostwald dilution law will take the form:
§ is the average activity coefficient of the electrolyte.
For c→0 and γ→1, the above equation of the Ostwald dilution law takes the form (2). The more the electrolyte dissociates, the faster the value of the activity coefficient γ deviates from unity, and the faster the classical dilution law is violated.
[edit] Dissociation constant of strong electrolytes
Strong electrolytes dissociate almost completely (the reaction is irreversible), therefore, the denominator of the expression for the dissociation constant is zero, and the whole expression tends to infinity. Thus, for strong electrolytes, the term "dissociation constant" is meaningless.
[edit] Calculation examples
[edit] Water dissociation
Water is a weak electrolyte that dissociates according to the equation
The dissociation constant of water at 25 °C is
Considering that in most solutions water is in molecular form (the concentration of H + and OH - ions is low), and given that the molar mass of water is 18.0153 g / mol, and the density at a temperature of 25 ° C is 997.07 g / l, pure water corresponds to the concentration = 55.346 mol/l. Therefore, the previous equation can be rewritten as
The application of the approximate formula gives an error of about 15%:
Based on the found value of the degree of dissociation, we find the pH of the solution:
Degree of dissociation- a value characterizing the state of equilibrium in the dissociation reaction in homogeneous (homogeneous) systems.
The degree of dissociation α is equal to the ratio of the number of dissociated molecules n to the sum n + N, where N is the number of undissociated molecules. Often α is expressed as a percentage. The degree of dissociation depends both on the nature of the dissolved electrolyte and on the concentration of the solution.
[edit] Example
For acetic acid CH 3 COOH, the value of α is 4% (in a 0.01M solution). This means that in an aqueous solution of an acid, only 4 out of every 100 molecules are dissociated, that is, they are in the form of H + and CH 3 COO − ions, while the remaining 96 molecules are not dissociated.
[edit] Definition methods
§ according to the electrical conductivity of the solution
§ to lower the freezing point
[edit] Imaginary degree of dissociation
Since strong electrolytes dissociate almost completely, one would expect for them an isotonic coefficient equal to the number of ions (or polarized atoms) in the formula unit (molecule). However, in reality, this coefficient is always less than that determined by the formula. For example, the isotonic coefficient for a 0.05 mol NaCl solution is 1.9 instead of 2.0 (for a magnesium sulfate solution of the same concentration, i= 1.3). This is explained by the theory of strong electrolytes, developed in 1923 by P. Debye and E. Hückel: the movement of ions in solution is hindered by the formed solvation shell. In addition, ions interact with each other: oppositely charged ones attract, and likewise charged ones repel; the forces of mutual attraction lead to the formation of groups of ions moving through the solution together. Such groups are called ion associates or ion pairs. Accordingly, the solution behaves as if it contains fewer particles than it really is, because the freedom of their movement is limited. The most obvious example concerns the electrical conductivity of solutions λ , which increases with dilution of the solution. Through the ratio of real electrical conductivity to that at infinite dilution, determine imaginary degree of dissociation strong electrolytes, also referred to as α :
,
where nimg- imaginary, and n disslv. is the actual number of particles in the solution.
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 conditional charge on the atom, if we assume that the molecule was created by the ionic mechanism (or - this is the number of received or given electrons).
Restorers- atoms, molecules, ions - donating electrons.
Oxidizers- atoms, molecules, ions - accepting electrons.
Reducing agents participate in the oxidation process by 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 in the molecules of different substances, for example:
H2S + Cl2 → S + 2HCl
2. Intramolecular- reactions in which oxidizing and reducing atoms are in molecules of the same substance, for example:
2H2O → 2H2 + O2
3. Disproportionation(self-oxidation-self-recovery) - reactions in which the same element acts both as an oxidizing agent and as a reducing agent, for example:
Cl2 + H2O → HClO + HCl
4. Reproportionation (proportionate, counterdisproportionation) - 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 2Н+ from the reduced substrate to molecular oxygen: SH2 + O20 + 2e= S + H2O
Attachment of oxygen to the substrate: SH2 + 1/2O20 + 2e= HO - S -H
Mechanism of occurrence of electrode and redox potentials. Nernst-Peters equations.
A measure of the redox ability of substances are redox potentials. Let us consider the mechanism of the emergence of the potential. 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 electric layer on the metal surface, and the emergence of the potential of the Zn2+/Zn° pair.
A metal immersed in a solution of its salt, such as zinc in a solution of zinc sulfate, is called an electrode of the first kind. This is a two-phase electrode that is negatively charged. 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 salt, the opposite process is observed. At the interface between the metal and the salt solution, metal is deposited as a result of the reduction of an ion that has a high acceptor capacity for an electron, which is due to the high nuclear charge and the small radius of the ion. The electrode is positively charged, excess salt anions form a second layer in the near-electrode space, and the electrode potential of the Cu2+/Cu° pair arises. The potential is formed as a result of the recovery process (Fig. 8.2). The mechanism, magnitude and sign of the electrode potential are determined by the structure of the atoms involved in the electrode process.
So, the potential that arises at the interface between the metal and the solution as a result of the oxidation and reduction processes occurring with the participation of the metal (electrode) and the formation of a double electric layer is called electrode potential.
If electrons are removed from a zinc plate to a copper one, then the equilibrium on the plates is disturbed. To do this, we connect zinc and copper plates immersed in solutions of their salts with a metal conductor, near-electrode solutions with an electrolyte bridge (a tube with a K2SO4 solution) to close the circuit. The oxidation half-reaction proceeds on the zinc electrode:
and on copper - the reduction half-reaction:
The electric current is due to the total redox reaction:
An 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. Electric circuit diagram of a galvanic cell
Galvanic cell- 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 as a short diagram, where a more negative electrode is placed on the left, the pair formed on this electrode is indicated by a vertical line, and the potential jump is shown. Two lines mark the boundary between solutions. The charge of the electrode is indicated in parentheses: (-) Zn°|Zn2+||Cu2+|Cu° (+) - scheme of the chemical circuit of the galvanic cell.
The redox potentials of a 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 solution, the temperature of the solution, and are described by the Nernst equation.
The quantitative characteristic of the redox system is redox potential, arising at the phase boundary platinum - aqueous solution. The potential value in SI units is measured in volts (V) and is calculated from the Nernst-Peters equation:
where a(Ox) and a(Red) are the activities of the oxidized and reduced forms, respectively; R- universal gas constant; T- thermodynamic temperature, K; F- Faraday's constant (96,500 C/mol); n is the number of electrons involved in the elementary redox process; a - activity of hydronium ions; m- stoichiometric coefficient in front of the hydrogen ion in the half-reaction. The value of φ° is the standard redox potential, i.e. potential measured under the conditions a(Oх) = a(Red) = a(H+) = 1 and a given temperature.
The standard potential of the 2H+/H2 system is assumed to be 0 V. The standard potentials are reference values and are tabulated at a temperature of 298K. A strongly acidic environment is not characteristic of biological systems, therefore, to characterize the processes occurring in living systems, the formal potential is more often used, which is determined under the condition a(Ox) = a(Red), pH 7.4, and a temperature of 310 K (physiological level). When writing the potential, the vapor is indicated as a fraction, with the oxidizer being written in the numerator and the reducing agent in the denominator.
For 25 °C (298K) after substitution of 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 couple, V; so.fu 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 with it is determined by the equation:
where [Mn+] is the equilibrium concentration of the metal ion; n- the number of electrons involved in the half-reaction, and corresponds to the oxidation state of the metal ion.
Redox systems are divided into two types:
1) only electron transfer Fe3+ + ē = Fe2+, Sn2+ - 2ē = Sn4+ takes place in the system. This isolated redox equilibrium;
2) systems where electron transfer is supplemented by proton transfer, i.e. observed combined equilibrium of different types: protolytic (acid-base) and redox with possible competition of 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 calculation of the potential of a pair is carried out according to the formula:
In a wider range of conjugated pairs, the oxidized and reduced forms of the pair are in solution in various degrees of oxidation (MnO4-/Mn2+). As measuring electrode
in this case, an electrode made of an 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 formed due to the redox process occurring in solution is called redox potential.
It is measured onredox electrode is an inert metal in solution containing oxidized and reduced forms of a pair. For example, when measuring Eo Fe3+/Fe2+ pairs use a redox electrode - a platinum measuring electrode. The reference electrode is hydrogen, the potential of the pair of which is known.
The reaction taking place in the galvanic cell:
Chemical chain scheme: (-)Pt|(H2°), H+||Fe3+, Fe2+|Pt(+).
So, the redox potential (ORP) is the potential of the system in which the activities of the oxidizing and reducing forms of a given substance are equal to unity. ORP is measured using redox electrodes in combination with standard reference electrodes.
Each redox reaction has its own redox pair- this pair has a substance in the oxidized and reduced form (Fe + 3 / Fe + 2).
A quantitative measure of the activity of a redox pair is its ORP value.
ORPpairs>>>oxidizer
ORPcouples<<<восстановитель
ORP depends on:
The nature of the redox pair,
Concentrations
Temperatures
Comparative strength of oxidizing and reducing agents. Predicting the direction of redox processes by the values of redox potentials.
The redox potential is a measure of the redox ability of substances. The value of the standard pair potentials is indicated in the reference tables.
The standard potentials of the electrodes (E°), acting as reducing agents with respect to hydrogen, have the “-” sign, and the “+” sign has the standard potentials of the electrodes, which are oxidizing agents.
Metals, arranged in ascending order of their standard electrode potentials, form the so-called electrochemical series of voltages of metals: 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.
In the series of redox potentials, the following regularities are noted.
1. If the standard redox potential of a pair is negative, for example φ°(Zn2+(p)/Zn°(t)) = -0.76 V, then with respect to a hydrogen pair whose potential is higher, this pair acts as a reducing agent. The potential is formed by the first mechanism (oxidation reactions).
2. If the potential of the pair is positive, for example φ ° (Cu2 + (p) / Cu (t)) \u003d +0.345 V with respect to a hydrogen or other conjugated pair whose potential is lower, this pair is an oxidizing agent. The potential of this pair is formed according to the second mechanism (reduction reactions).
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 pair. A decrease in the value of the positive potential and an increase in the negative potential corresponds to a decrease in the oxidative and an increase in the reduction activity. For example:
Comparison of the values of standard redox potentials allows us to answer the question: does this or that redox reaction proceed?
The difference between the standard oxidation potentials of the oxidized and reduced semi-pairs is called the electromotive force (EMF).
E0 = Eok-Evost
A quantitative criterion for assessing the possibility of a particular redox reaction 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 flow under standard conditions of OVR, it is necessary:
G0298= - P F E0
E> 0 G< 0 - самопроизвольно
E< 0 G>0 - back
E \u003d 0 G \u003d 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 there are ensembles of enzymes - dehydrogenases, coenzymes (NAD +, FAD, UBX), a series of cytochromes b, c1, c and an enzyme - cytochrome oxidase. They form a system of the cellular respiratory chain, with the help of which the relay-race transfer of protons and electrons from the substrate to the oxygen molecules delivered by hemoglobin to the cell takes place.
Each component of the respiratory chain is characterized by a certain value of the redox potential. The movement of electrons along the respiratory chain occurs in steps from substances with a low potential (-0.32 V) to substances with a higher potential (+0.82 V), since any compound can only donate electrons to a compound with a higher redox potential (table one).
Table 1
Standard redox potentials of respiratory chain biomolecules
SYSTEM
HALF REACTION
REDOX POTENTIAL, V
OVER+/OVER×N
OVER+ + H+ + 2 ē → OVER×H
FAD/FAD×H2
FAD+ + 2Н+ + 2 ē → FAD×Н2
UBH/ UBH×H2
UBK+ 2Н+ + 2 ē → UBK×Н2
cytochrome b
cytochrome c1
cytochrome c
cytochrome a + a3
O2 + 4 H+ + 4 ē → 2 H2O
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, there is a relay transfer of a proton and a pair of electrons 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). Then 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 from UBH × H2 forms a water molecule.
1/2 О2 + 2Н+ + 2 ē → Н2О
It should be noted that each oxygen molecule interacts with two electron transport chains, since only one-electron transfer Fe3+ → Fe2+ is possible in the structure of cytochromes.
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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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 the oxidizing element and the reducing element are in the composition of different molecules of substances. The above reactions are of this type.
2. Intramolecular (intramolecular oxidation - reduction).
These include reactions in which the oxidizing agent and reducing agent in the form of atoms of different elements are part of the same molecule. Thermal decomposition reactions of compounds proceed according to this type, for example:
2KCIO 3 = 2KCI + 3O 2 .
3. Disproportionation (self-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 simultaneously. 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 reaction of photosynthesis in plants and the processes of respiration in animals and humans. The processes of fuel combustion occurring in the furnaces of boilers of thermal power plants and in internal combustion engines are an example of RWR.
OVR are used in the production of metals, organic and inorganic compounds, they are used to purify various substances, natural and waste waters.
9.5. Redox (electrode) potentials
A measure of the redox ability of substances is their electrode or redox potentials j ox / Red (redox potentials). electrons. It is customary to write redox systems in the form of reversible reduction reactions:
Oh + ne - D Red.
The mechanism of the occurrence of the electrode potential. Let us explain the mechanism of the occurrence of an 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 Me n + 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 a lot of energy. When a metal is immersed in an aqueous solution of a salt containing metal cations in its composition, polar water molecules, respectively, orienting themselves at the surface of the metal (electrode), interact with surface metal cations (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 of Me n + * m H 2 O (p) + ne-
The metal becomes negatively charged and the solution positively charged. Positively charged ions from the solution are attracted to the negatively charged metal surface (Me). A double electric layer appears at the metal-solution boundary (Fig. 9.2). The potential difference between a metal and a solution is called electrode potential or redox potential of the electrode φ Me n + / Me(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 the ions pass into the solution, the negative charge of the metal surface and the positive charge of the solution increase, which prevents the oxidation (ionization) of the metal.
In parallel with the oxidation process, the reverse reaction proceeds - the reduction of metal ions from the 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.
With an increase in the potential difference between the electrode and the solution, the rate of the forward reaction decreases, while 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, and equilibrium is established:
Me n + * m H 2 O (p) + ne - D Me (k) + m H 2 O.
To simplify, water of hydration is usually not included in the reaction equation and it is written as
Me n + (p) + ne - D Me (k)
or in general terms for any other redox systems:
Oh + ne - D Red.
The potential established under the conditions of equilibrium of the electrode reaction is called equilibrium electrode potential. In the considered case, the ionization process in the solution is thermodynamically possible, and the metal surface is charged negatively. For some metals (less active), thermodynamically more probable is the process of reduction of hydrated ions to metal, then their surface is positively charged, and the adjacent electrolyte layer is negatively charged.
Hydrogen electrode device. 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 conditionally taken equal to zero. As a reference electrode, a standard hydrogen electrode, related to gas electrodes, is often used. In the general case, gas electrodes consist of a metal conductor that is in contact simultaneously with a gas and a solution containing an oxidized or reduced form of an element that is part of the gas. The metal conductor serves to supply and remove electrons and, in addition, is a catalyst for the electrode reaction. The metal conductor must 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 a loose porous plate (to increase 
electrode surface) and immersed in an aqueous solution of sulfuric acid with an activity (concentration) of H + ions equal to one.
Hydrogen is passed through a solution of sulfuric acid under atmospheric pressure. Platinum (Pt) is an inert metal that practically does not interact with a solvent, solutions (does not send its ions into a solution), but it is able to adsorb molecules, atoms, ions of other substances. When platinum comes into contact with molecular hydrogen, hydrogen is adsorbed on platinum. Adsorbed hydrogen, interacting with water molecules, goes into solution in the form of ions, leaving electrons in platinum. In this case, platinum is charged negatively, and the solution is positively charged. There is a potential difference between the platinum and the solution. Along with the transition of ions into the solution, the reverse process occurs - the reduction of H + ions from the solution with the formation of hydrogen molecules . The equilibrium on the hydrogen electrode can be represented by the equation
2Н + + 2е - D Н 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 conventionally assumed to be zero: j 0 2H + / H2 = 0 V.
Standard electrode potentials . Electrode potentials measured with respect to a standard hydrogen electrode under standard conditions(T = 298K; for dissolved substances, the concentration (activity) C Red = C ox = 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 denoted by j 0 O x / Red. These are reference values.
The oxidizing ability of substances is the higher, the greater the algebraic value of their standard electrode (redox) potential. On the contrary, the smaller the value of the standard electrode potential of the reactant, the more pronounced its reducing properties. For example, comparing the standard potentials of systems
F 2 (g.) + 2e - D 2F (p.) j 0 \u003d 2.87 V
H 2 (r.) + 2e - D 2H (r.) j 0 \u003d -2.25 V
shows that the F 2 molecules have a pronounced oxidative tendency, while the H ions have a reduction tendency.
A number of stresses of metals. By arranging the metals in a row as the algebraic value of their standard electrode potentials increases, the so-called “Standard Electrode Potential Series” or “Voltage Series” or “Metal Activity Series” are obtained.
The position of the metal in the "Row 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 reduction properties of the 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 oxidizes only when interacting with very strong oxidizing agents. From the data of the "Series of voltages" it can be seen that the ions of lithium (Li +), potassium (K +), calcium (Ca 2+), etc. - the weakest oxidizing agents, and the strongest oxidizing agents are mercury ions (Hg 2+), silver (Ag +), palladium (Pd 2+), platinum (Pt 2+), gold (Au 3+, Au +).
Nernst equation. The 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 solute and solvent, the 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 is the universal gas constant; T is the absolute temperature; n is the number of electrons involved in the electrode process; and ox, and Red are the activities (concentrations) 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 metallic and the equilibria established on them are described in general form
Me n + + ne - D Me,
the Nernst 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 \u003d 96485 C / mol, replacing the activity a Me n + with the molar concentration of metal ions in the C Me n + solution and introducing a factor of 2.303 (transition to decimal logarithms), we obtain the Nernst equation in the form
j Me n + / Me = j 0 Me n + / Me + lg C Me n + .
Distinguish reactions intermolecular, intramolecular and self-oxidation-self-healing (or disproportionation):
If the oxidizing and reducing agents are the elements that make up the composition different compounds, the reaction is called intermolecular.
Example: Na 2 S O 3 + O 2 Na 2 SO 4
sun-ok-l
If the oxidizing agent and reducing agent are elements that make up the same compound, then the reaction is called intramolecular.
Example: ( N H4) 2 Cr 2 O 7 N 2 + Cr 2 O 3 + H 2 O.
v-l o-l
If the oxidizing agent and reducing agent is the same element while some of its atoms are oxidized, and the other is reduced, then the reaction is called self-oxidation-self-healing.
Example: H 3 P O 3 H 3 P O4+ P H3
v-l / o-l
Such a classification of reactions turns out to be convenient in determining the potential oxidizing and reducing agents among given substances.
4 Determination of the possibility of redox
reactionsaccording to the oxidation states of the elements
A necessary condition for the interaction of substances in 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 go, because No
o–l o–l potential reducing agent;
2) Zn + KI ... - the reaction does not take place, because No
v–l v–l potential oxidizing agent;
3) KNO 2 + KBiO 3 + H 2 SO 4 ...- the reaction is possible if at the same time
v-l o-l KNO 2 will be a reducing agent;
4) KNO 2 + KI + H 2 SO 4 ... - the reaction is possible if at the same time
o - l in - l KNO 2 will be an oxidizing agent;
5) KNO 2 + H 2 O 2 ... - the reaction is possible if at the same time
c - l o - l H 2 O 2 will be an oxidizing agent, and KNO 2
Reducing agent (or vice versa);
6) KNO 2 ... - possible reaction
o - l / in - l disproportionation
The presence of a potential oxidizing agent and reducing agent is a necessary but not sufficient condition for the reaction to proceed. So, in the examples considered above, only in the fifth one can it be said that one of the two possible reactions will occur; in other cases, additional information is needed: whether this reaction will energetically beneficial.
5 The choice of oxidizing agent (reducing agent) using tables of electrode potentials. Determination of the predominant direction of redox reactions
Reactions proceed 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.
v–l o–l combination of two
half reactions:
Zn Zn 2+ and Cu 2+ Cu;
the first one, which includes 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 magnitude of the electrode potential, which denote, respectively,
E restore = E 0 Zn 2+ / Zn and E approx. \u003d 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 \u003d E 0 Cu 2+ / Cu - E 0 Zn 2+ / Zn \u003d 0.34 - (-0.77) \u003d 1.1V.
Obviously, 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, it is possible to determine which reaction, direct or reverse, proceeds predominantly, as well as choose an oxidizing agent (or reducing agent) for a given substance.
In the above example, E 0 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 scheme of a possible reaction: I - + Fe 3+ I 2 + Fe 2+,
v-l o-l
b) write the half-reactions for the oxidizing and reducing systems and the corresponding electrode potentials:
Fe 3+ + 2e - Fe 2+ E 0 \u003d + 0.77 B - oxidizing system,
2I - I 2 + 2e - E 0 \u003d + 0.54 B - recovery system;
c) comparing the potentials of these systems, we conclude that the given reaction is possible (under standard conditions).
2. Choose oxidizing agents (at least three) for a given transformation of a substance and choose from them the one in which the reaction proceeds most fully: Cr (OH) 3 CrO 4 2 -.
Solution:
a) find in the reference book E 0 CrO 4 2 - / Cr (OH) 3 \u003d - 0.13 V,
b) we select suitable oxidizing agents using the reference book (their potentials should be greater than - 0.13 V), while focusing on the most typical, “non-deficient” oxidizing agents (halogens are simple substances, hydrogen peroxide, potassium permanganate, etc. ).
In this case, it turns out that if the transformation Br 2 2Br - corresponds to one potential E 0 \u003d + 1.1 V, then for permanganate ions and hydrogen peroxide, options are possible: E 0 MnO 4 - / Mn 2+ \u003d + 1.51 B - in sour environment,
E 0 MnO 4 - / MnO 2 \u003d + 0.60 B - in neutral environment,
E 0 MnO 4 - / MnO 4 2 - \u003d + 0.56 B - in alkaline environment,
E 0 H 2 O 2 / H 2 O \u003d + 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 slightly alkaline or neutral environment, the following are suitable oxidizing agents:
E 0 MnO4 - / MnO2 \u003d + 0.60 B and. E 0 Br2 /Br - = + 1.1 B..
c) the last condition, the choice of the optimal oxidant from several, is decided on the basis that the reaction proceeds the more completely, the more negative G 0 for it, which in turn is determined by the value E 0:
The larger the algebraic value 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).
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