In the formation of the chemical properties of soils, redox processes occupy one of the leading places. The most important factors, which determine the redox state of soil horizons, is the oxygen of the soil air and soil solutions, oxide and ferrous compounds of iron, manganese, nitrogen, sulfur, organic matter, and microorganisms.
Oxidation and reduction reactions always occur simultaneously. The oxidation of one substance participating in the reaction is accompanied by the reduction of another substance.
Redox processes mean processes in which, as possible stage involves the transfer of electrons from one particle of matter to another. Oxidation is a reaction in which oxygen is added to a substance or the substance loses hydrogen or electrons. Reduction is the loss of oxygen by a substance and the addition of hydrogen or electrons to the substance.
The ability of a soil to undergo redox reactions is measured using the oxidation reduction potential (ORP).
The redox potential relative to hydrogen is called Eh. This value depends on the concentration and ratio of oxidizing agents and reducing agents formed during 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. Eh values in various types soils and soil horizons vary within 100-800 mV, sometimes having negative values. The value of Eh significantly depends on the acid-base conditions of the environment, vegetation and microorganisms.
Under soil conditions, a significant portion of the components involved in redox reactions are represented by solid phases. In reactions involving solid phases, the soil will be highly buffering until these components react. Buffer capacity is the ability of soil to withstand changes in redox potential at any external influences. This concept characterizes the stability of soil redox systems under natural dynamic conditions and can be called dynamic buffering. In a natural environment, humic substances and iron hydroxide minerals react at low speeds.
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 systems are those that, in the process of changing the redox regime, do not change the total supply of components. Irreversible systems in the process of changing the redox regime, some substances are lost. These substances become gaseous or precipitate. As a rule, irreversible systems predominate in soils.
Reversible redox systems include:
System Fe3+ ⇆Fe2+. This system takes special place among reversible systems. It reacts sensitively to the slightest changes in the redox environment. The solubility of ferric iron compounds is extremely low. Migration of iron compounds is possible mainly in the form of divalent iron compounds under conditions increased acidity and decreased Eh.
System Mn2+ ⇆ Mn4+. This system is extremely sensitive to changes in ORP. Tetravalent manganese compounds are insoluble under conditions characteristic of soil horizons. Exchangeable manganese is divalent. The concentration of divalent manganese ions increases tens of thousands of times with increasing acidity and decreasing Eh. The migration of manganese compounds during soil-forming processes in the vertical and horizontal directions is similar to the migration of iron compounds.
Irreversible redox systems include:
System NO3 → NO2 → NO → N. The process of nitrification and accumulation of nitrates occurs under oxidative conditions 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 sulfuric acid salts are present. With the participation of microorganisms, the sulfate-sulfide system in the presence of organic matter and lack of oxygen shifts towards sulfides. The process of reduction of sulfates to sulfur metals occurs:
Na2SO4 + 2C = Na2S + CO2
Under the influence of carbon dioxide present in the soil, sulfur metals easily decompose and form bicarbonates and carbonates of alkali and alkaline earth metals. In this case, the process of sulfate reduction occurs:
Na2S + H2CO3 = Na2CO3 + H2S
However, in the soil solution the content of elements with variable valence is quite small. Therefore, the soil solution has low OM capacity and buffer capacity, and the Eh value is unstable.
A more significant influence on OM processes in soils is exerted by oxygen dissolved in the soil solution, soil microflora and water.
Almost all soil reactions occur in an aqueous environment, and water itself can act as both an oxidizing agent and a reducing agent.
Based on the characteristics 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.
OM processes are closely related to the transformation of plant residues, the accumulation and composition of the resulting organic matter, and as a result, the formation of the soil profile.
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, relative 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 are called. normal.
O. o.-v. With. According to the magnitude of normal potentials, they can be arranged in a series, with each system being 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 substances in sedimentary rocks, etc.
Substance equivalent or Equivalent is a real or fictitious 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 = NaCl + H2O
the equivalent will be a real particle - the Na + ion, in the reaction
the equivalent will be the imaginary particle ½Zn(OH) 2.
The equivalent of a substance is also often meant 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 a 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 actual molar mass called equivalence factor(usually designated as ).
Molar mass of 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 designated as ).
[edit]Equivalence number
Equivalence number z is a small positive integer equal to the number of equivalents of a substance contained in 1 mole of that substance. The equivalence factor is related to the equivalence number z the following relationship: =1/z.
For example, in react:
Zn(OH) 2 + 2HCl = ZnCl 2 + 2H 2 O
The equivalent is the ½Zn(OH) 2 particle. The number ½ is equivalence factor, z V in this case equals 2
* - for inert gases Z = 1
The equivalence factor helps formulate the law of equivalence.
[edit]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 =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 amount of electricity passes through the electrolyte:
§ where is Faraday's constant.
§ Faraday's constant, is a physical constant that determines the relationship between the electrochemical and physical properties of a substance.
§ Faraday's constant is equal to C mol −1.
§ Faraday's constant is included as a constant in Faraday's second law(law of electrolysis).
§ Numerically, Faraday’s constant is equal to the electric charge, when passing through the electrolyte on the electrode, (1/z) mole of substance A is released in the formula:
Where: 
- the number of electrons participating in the reaction.
§ For Faraday’s constant, the following relation holds:
§ where is the elementary charge, and is Avogadro’s number.
Isotopes(from ancient Greek ισος - "equal", "same", and τόπος - "place") - varieties of atoms (and nuclei) of one chemical element With different quantities neutrons in the nucleus. The name is due to the fact that isotopes are located in the same place (in the same cell) of the periodic table. The chemical properties of an atom depend almost exclusively 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. Typically, an isotope is designated by the symbol of the chemical element to which it belongs, with the addition of an upper-left subscript indicating the mass number (for example, 12 C, 222 Rn). You can also write the name of the element followed by a hyphenated mass number (for example, carbon-12, radon-222). Some isotopes have traditional proper names (for example, deuterium, actinone).
Example of isotopes: 16 8 O, 17 8 O, 18 8 O - three stable isotopes of oxygen.
[edit] Terminology
The basic position of IUPAC is that the correct singular term to designate the atoms (or nuclei) of the same chemical element with the same atomic mass is a nuclide, and the term isotopes may be used to designate a set of nuclides of one element. Term isotopes was proposed and used initially in plural, since at least two types of atoms are needed for comparison. Subsequently, the use of the term in the singular also became widespread - isotope. In addition, the plural term is often used to refer to any collection of nuclides, and not just one element, which is also incorrect. Currently, the positions of international scientific organizations are not brought to uniformity and the term isotope continues to be widely used, including in official materials of various divisions of IUPAC and IUPAP. This is one example of how the meaning of a term, originally inherent in it, ceases to correspond to the concept for which this term is used to designate (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 was obtained by studying the radioactive transformations of atoms of heavy elements. In 1906-07, it turned out that the product of radioactive decay of uranium - ionium and the product of radioactive decay - radiothorium, have the same chemical properties as thorium, but differ from it in atomic mass and radioactive decay characteristics. It was later discovered that all three products had identical optical and x-ray spectra. Such substances, identical in chemical properties, but different in 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 disruption of the isotopic composition of elements (natural fractionation isotopes characteristic of light elements, as well as isotope shifts during the decay of natural long-lived isotopes). The gradual accumulation of nuclei in minerals - the decay products of some long-lived nuclides - is used in nuclear geochronology.
[edit]Use of isotopes by humans
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 fission chain reactions with thermal neutrons and can be used as fuel for nuclear reactors or nuclear weapons. However, natural uranium contains only 0.72% of this nuclide, while chain reaction is practically feasible only with a 235 U content of at least 3%. Due to the proximity physical and chemical properties isotopes of heavy elements, the procedure for isotope enrichment of uranium is an extremely complex technological task that is accessible to only a dozen countries in the world. Isotopic tags are used in many branches of science and technology (for example, in radioimmunoassay).
Dissociation constant- a type of equilibrium constant that shows the tendency of a large object to dissociate (separate) reversibly into small objects, such as when a complex breaks down into its constituent molecules, or when a salt separates into aqueous solution to ions. The dissociation constant is usually denoted Kd and the inverse of the association constant. In the case of salts, the dissociation constant is sometimes called the ionization constant.
where complex A x B y breaks down into x units A and y units of B, the dissociation constant is determined as follows:
where [A], [B] and are the concentrations of A, B and the complex A x B y, respectively.
[edit]Definition
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 − is an 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;
§ is the concentration of anions in solution.
The equilibrium constant in relation to a dissociation reaction is called dissociation constant.
[edit]Dissociation of electrolytes with multivalent ions
In the case of dissociation of electrolytes with multivalent ions, dissociation occurs in steps, and for each step there is eigenvalue dissociation constants.
Example: Dissociation of 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 occurs mainly in the first stage.
[edit]Relation between the dissociation constant and the degree of dissociation
Based on the definition of the degree of dissociation, for the electrolyte CA 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 Ostwald's dilution law. For very small α (α<<1) K=cα² и
Thus, when the electrolyte concentration increases, the degree of dissociation decreases, and when it decreases, it increases. The relationship between the dissociation constant and the degree of dissociation is described in more detail in the article Ostwald's Law of Dilution.
[edit]Difference between experimental results and the Arrhenius model, derivation of the dissociation constant through activities
The above calculations are based on the Arrhenius theory, which is too crude and does not take into account the factors of electrostatic interaction of ions. Deviations from the ideal state in electrolyte solutions occur at very low concentrations, since interionic forces are inversely proportional 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 for real solutions simple equations can be preserved (see above) if instead of ion concentrations one introduces its function, the so-called activity. Activity (a) is related to concentration (c) through a correction factor γ, called the activity coefficient:
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 relation:
If there are no other influences that deviate the solution from the ideal state, then undissociated molecules behave like ideal gases and γ KA = 1, and the true expression of the Ostwald dilution law takes the form:
§ - average electrolyte activity coefficient.
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 activity coefficient γ deviates from unity, and the faster the violation of the classical dilution law occurs.
[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 entire expression tends to infinity. Thus, for strong electrolytes the term “dissociation constant” is meaningless.
[edit]Examples of calculations
[edit]Dissociation of water
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 taking into account 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 concentration = 55.346 mol/l. Therefore, the previous equation can be rewritten as
Using an 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 quantity 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 amount n + N, Where N- number of undissociated molecules. α is often 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 0.01 M 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]Determination methods
§ according to the electrical conductivity of the solution
§ by lowering the freezing point
[edit]Imaginary degree of dissociation
Since strong electrolytes dissociate almost completely, one would expect an isotonic coefficient for them 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 solution of NaCl is 1.9 instead of 2.0 (for a solution of magnesium sulfate of the same concentration it is completely 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 a solution is hindered by the formed solvation shell. In addition, ions also interact with each other: oppositely charged ones attract, and similarly charged ones repel; Mutual attractive forces lead to the formation of groups of ions moving through the solution together. Such groups are called ionic associates or ion pairs. Accordingly, the solution behaves as if it contains fewer particles than it actually does, because their freedom of movement is limited. The most obvious example concerns the electrical conductivity of solutions. λ , which increases with dilution of the solution. Using the ratio of the actual electrical conductivity to that at infinite dilution, one determines imaginary degree of dissociation strong electrolytes, also denoted by α :
,
Where n img- imaginary, and n disslv.- 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 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 (a 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 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 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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There are three main types of redox reactions:
1. Intermolecular (intermolecular oxidation - reduction).
This type includes the most numerous reactions in which the atoms of an oxidizing element and a reducing element are located in different molecules of substances. The reactions discussed above belong to this type.
2. Intramolecular (intramolecular oxidation - reduction).
These include reactions in which an oxidizing agent and a reducing agent in the form of atoms of different elements are contained in the same molecule. Thermal decomposition reactions of compounds proceed according to this type, for example:
2KCIO 3 = 2KCI + 3O 2.
3. Disproportionation (auto-oxidation - self-healing).
These are reactions in which the oxidizing and reducing agent is the same element in the same intermediate oxidation state, which as a result of the reaction both decreases and increases. For example:
3CI 0 2 + 6 KOH = 5 KCI + KCIO 3 + 3H 2 O,
3HCIO = HCIO 3 + 2HCI.
Redox reactions play an important role in nature and technology. Examples of OVR occurring in natural biological systems include the photosynthesis reaction in plants and respiration processes in animals and humans. Fuel combustion processes occurring in the furnaces of boilers at thermal power plants and in internal combustion engines are an example of OVR.
ORRs are used in the production of metals, organic and inorganic compounds, and carry out the purification of various substances, natural and waste waters.
9.5. Oxidation-reduction (electrode) potentials
A measure of the oxidation-reduction capacity of substances is their electrode or redox potential j ox / Red (redox potentials).1 The oxidation-reduction potential characterizes the redox system, consisting of the oxidized form of the substance (Ox), the reduced form (Red) and electrons. It is customary to write redox systems in the form of reversible reduction reactions:
Ox + ne - D Red.
Mechanism of occurrence of electrode potential. We will explain the mechanism of occurrence of the electrode or redox potential using the example of a metal immersed in a solution containing its ions. All metals have a crystalline structure. The crystal lattice of a metal consists of positively charged Men + ions and free valence electrons (electron gas). In the absence of an aqueous solution, the release of metal cations from the metal lattice is impossible, because this process requires large energy costs. When a metal is immersed in an aqueous solution of a salt containing metal cations, polar water molecules, correspondingly oriented at the surface of the metal (electrode), interact with the surface cations of the metal (Fig. 9.1).
As a result of the interaction, the metal is oxidized and its hydrated ions go into solution, leaving electrons in the metal:
Me (k) + m H 2 Oxidation Me n+ *m H 2 O (p) + ne-
The metal becomes negatively charged, and the solution becomes positively charged. Positively charged ions from the solution are attracted to the negatively charged metal surface (Me). An electric double layer appears at the metal-solution interface (Fig. 9.2). The potential difference arising between the metal and the solution is called electrode potential or redox potential of the electrode φ Ме n + /Ме(φ Ox / Red in general). A metal immersed in a solution of its own salt is an electrode (Section 10.1). The symbol of the metal electrode Me/Me n + reflects the participants in the electrode process.
As ions pass into solution, the negative charge of the metal surface and the positive charge of the solution increase, which prevents oxidation (ionization) of the metal.
In parallel with the oxidation process, a reverse reaction occurs - the reduction of metal ions from solution to atoms (metal precipitation) with the loss of the hydration shell on the metal surface:
Me n+ * m H 2 O(p) + ne- reduction Me(k) + m H 2 O.
As the potential difference between the electrode and the solution increases, the rate of the forward reaction decreases, and the rate of the reverse reaction increases. At a certain value of the electrode potential, the rate of the oxidation process will be equal to the rate of the reduction process, equilibrium is established:
Me n + * m H 2 O (p) + ne - D Me (k) + m H 2 O.
For simplicity, hydration water is usually not included in the reaction equation and it is written as
Ме n + (р) + ne - D Ме (к)
or in general form for any other redox systems:
Ox + ne - D Red.
The potential established under equilibrium conditions of the electrode reaction is called equilibrium electrode potential. In the case considered, the ionization process in solution is thermodynamically possible, and the metal surface becomes negatively charged. For some metals (less active), the process of reduction of hydrated ions to metal is thermodynamically more probable, then their surface is charged positively, and the layer of adjacent electrolyte is charged negatively.
Hydrogen electrode device. The absolute values of electrode potentials cannot be measured; therefore, their relative values are used to characterize electrode processes. To do this, find the potential difference between the measured electrode and the reference electrode, the potential of which is conventionally assumed to be zero. A standard hydrogen electrode, which is classified as a gas electrode, is often used as a reference electrode. In general, gas electrodes consist of a metal conductor in contact simultaneously with a gas and a solution containing an oxidized or reduced form of an element 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 should not send its own ions into the solution. Platinum and platinum metals satisfy these conditions.
The hydrogen electrode (Fig. 9.3) is a platinum plate coated with a thin layer of loose porous plate (to increase 
surface of the electrode) and immersed in an aqueous solution of sulfuric acid with an activity (concentration) of H + ions equal to unity.
Hydrogen is passed through a solution of sulfuric acid under atmospheric pressure. Platinum (Pt) is an inert metal that practically does not interact with solvents or solutions (does not send its ions into the solution), but it is capable of adsorbing molecules, atoms, and ions of other substances. When platinum comes into contact with molecular hydrogen, hydrogen is adsorbed on the platinum. Adsorbed hydrogen, interacting with water molecules, goes into solution in the form of ions, leaving electrons in the platinum. In this case, the platinum is charged negatively, and the solution – positively. A potential difference arises between the platinum and the solution. Along with the transition of ions into solution, the reverse process occurs - the reduction of H + ions from solution with the formation of hydrogen molecules . Equilibrium at the hydrogen electrode can be represented by the equation
2H + + 2e - D H 2.
Symbol for hydrogen electrode H 2 , Pt│H + . The potential of the hydrogen electrode under standard conditions (T = 298 K, P H2 = 101.3 kPa, [H + ] = 1 mol/l, i.e. pH = 0) is assumed to be conditionally equal to zero: j 0 2H + / H2 = 0 V.
Standard electrode potentials . Electrode potentials measured relative to a standard hydrogen electrode under standard conditions(T = 298 K; for dissolved substances, concentration (activity) C Red = C o x = 1 mol/l or for metals C Me n + = 1 mol/l, and for gaseous substances P = 101.3 kPa), are called standard electrode potentials and are designated j 0 O x / Red. These are reference values.
The higher the algebraic value of their standard electrode (redox) potential, the higher the oxidative capacity of substances. On the contrary, the lower the standard electrode potential of the reactant, the more pronounced its reducing properties. For example, comparison of standard system potentials
F 2 (g.) + 2e - D 2F (p.) j 0 = 2.87 V
H 2 (r.)+ 2e - D 2H (r.) j 0 = -2.25 V
shows that F 2 molecules have a strongly pronounced oxidative tendency, while H ions have a reducing tendency.
A range of metal stresses. By arranging metals in a series as the algebraic value of their standard electrode potentials increases, the so-called “Series of standard electrode potentials” or “Series of voltages”, or “Series of activity of metals” is obtained.
The position of the metal in the “Series of Standard Electrode Potentials” characterizes the reducing ability of metal atoms, as well as the oxidizing properties of metal ions in aqueous solutions under standard conditions. The lower the value of the algebraic value of the standard electrode potential, the greater the reducing properties of a given metal in the form of a simple substance, and the weaker the oxidizing properties of its ions and vice versa. .
For example, lithium (Li), which has the lowest standard potential, is one of the strongest reducing agents, while gold (Au), which has the highest standard potential, is a very weak reducing agent and only oxidizes when interacting with very strong oxidizing agents. From the data of the “Voltage Series” it is clear that the ions of lithium (Li +), potassium (K +), calcium (Ca 2+), etc. - the weakest oxidizing agents, and the strongest oxidizing agents include ions of mercury (Hg 2+), silver (Ag +), palladium (Pd 2+), platinum (Pt 2+), gold (Au 3+, Au +).
Nernst equation. Electrode potentials are not constant. They depend on the ratio of concentrations (activities) of the oxidized and reduced forms of the substance, on temperature, the nature of the dissolved substance and solvent, pH of the medium, etc. This dependence is described Nernst equation:
,
where j 0 О x / Red is the standard electrode potential of the process; R – universal gas constant; T – absolute temperature; n is the number of electrons participating in the electrode process; a oh, a Red – activity (concentration) of the oxidized and reduced forms of the substance in the electrode reaction; x and y are stoichiometric coefficients in the electrode reaction equation; F is Faraday's constant.
For the case when the electrodes are metal and the equilibria established on them are described in general form
Me n + + ne - D Me,
Nernst's equation can be simplified by taking into account that for solids the activity is constant and equal to unity. For 298 K, after substituting a Me =1 mol/l, x=y=1 and constant values R=8.314 J/K*mol; F = 96485 C/mol, replacing the activity of a Me n + with the molar concentration of metal ions in a solution of C Me n + and introducing a factor of 2.303 (transition to decimal logarithms), we obtain the Nernst equation in the form
j Ме n + / Ме = j 0 Ме n + / Ме + lg С Ме n + .
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 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, select suitable oxidizing agents (their potentials must be greater than -0.13 V), while focusing on the most typical, “non-deficient” oxidizing agents (halogens - simple substances, hydrogen peroxide, potassium permanganate, etc. ).
It turns out that if the transformation Br 2 2Br – corresponds to one potential E 0 = +1.1 V, then for permanganate ions and hydrogen peroxide the following options are possible: E 0 MnO 4 – / Mn 2+ = + 1.51 V - V sour environment,
E 0 MnO 4 – / MnO 2 = + 0.60 V - in neutral environment,
E 0 MnO 4 – / MnO 4 2 – = + 0.56 V - in alkaline environment,
E 0 H 2 O 2 / H 2 O = + 1.77 B - in sour environment,
E 0 H 2 O 2/ OH – = + 0.88 B - in alkaline environment.
Considering that the chromium hydroxide specified by the condition is amphoteric and therefore exists only in a weakly alkaline or neutral environment, the following oxidants are suitable:
E 0 MnO4 – /MnO2 = + 0.60 V and. E 0 Br2 /Br – = + 1.1 B..
c) the last condition, the choice of the optimal oxidizer from several, is solved on the basis that the reaction proceeds more completely, the more negative G 0 is for it, which in turn is determined by the value E 0:
The larger algebraically the quantity E 0 , especially the redox reaction proceeds fully, the greater the yield of products.
Of the oxidizing agents discussed above, E 0 will be the largest for bromine (Br 2).
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