Redox processes. redox potentials. redox systems redox systems
(OV) AND OV - ELECTRODES.
Depending on the mechanism of oxidation - reduction, various OM - systems can be divided into two types:
1st type: RH - systems in which the redox process is associated with the transfer of only electrons, for example: Fe³ + +ē ↔ Fe² +
2nd type: OB - systems in which the redox process is associated not only with the transfer of electrons, but also of protons, for example:
C 6 H 4 O 2 + 2H + + 2ē ↔ C 6 H 4 (OH) 2
quinone hydroquinone
MnO 4 - + 8H + + 5ē ↔ Mn² + + 4H 2 O
An inert metal in combination with an RH system is called a redox or redox electrode, and the potential arising at this electrode is called a redox (OR) or redox potential.
The inert metal takes only an indirect part in the potential-determining reaction, being an intermediary in the transfer of electrons from the reduced form of the Red substance to the oxidized OX.
When an inert metal is immersed in a solution containing an excess of the oxidized form of iron, the metal plate is positively charged (Fig. 10 a)
With an excess of the reduced form of iron, the surface of platinum becomes negatively charged (Fig. 10b).
Rice. 10. Occurrence of the OB potential
The transfer of electrons from one ion to another through a metal leads to the formation of a DEL on the metal surface.
Interion exchange of electrons is also possible without metal. But the Fe²+ and Fe³+ ions are solvated in different ways, and for electron transfer it is necessary to overcome the energy barrier. The transition of electrons from Fe²+ ions to the metal and from the metal surface to the Fe³+ ion is characterized by a lower activation energy.
When the activities of the Fe²+ and Fe³+ ions are equal, the platinum plate is positively charged, because the electron-acceptor capacity of Fe³+ ions is greater than the electron-donor capacity of Fe²+.
Peters equation.
The quantitative dependence of the OM - potential on the nature of the OM - system (φ°r), the ratio of the activities of the oxidized and reduced forms, temperature, and the activity of hydrogen ions is established by the Peters equation.
1st type: φr = φ°r + ln
2nd type: φr = φ°r + ln
where φr - OB - potential, V;
φ°r - standard RH - potential, V;
z is the number of electrons participating in the OF process;
а (Ох) is the activity of the oxidized form, mol/l;
a (Red) is the activity of the reducing form, mol/l;
m is the number of protons;
a(n+) is the activity of hydrogen ions, mol/l.
The standard OB potential is the potential arising at the inert metal-solution interface, in which the activity of the oxidized form is equal to the activity of the reduced form, and for the system of the second type, in addition, the activity of hydrogen ions is equal to one.
Classification of reversible electrodes.
Having considered the principle of operation of the electrodes, we can conclude that according to the properties of the substances involved in the potential-determining processes, as well as according to the device, all reversible electrodes are divided into the following groups:
Electrodes of the first kind;
Electrodes of the second kind;
Ion-selective electrodes;
Redox electrodes.
1. A galvanic cell is a system that produces work, and does not consume it, so it is advisable to consider the EMF of the cell as a positive value.
2. The EMF of the element is calculated by subtracting from the numerical value of the potential of the right electrode the numerical value of the potential of the left electrode - the "right plus" rule. Therefore, the element circuit is written so that the left electrode is negative and the right electrode is positive.
3. The interface between the conductors of the first and second row is indicated by one line: Zn׀ZnSO4; Cu׀CuSO4
4. The interface between conductors of the second kind is depicted by a dotted line: ZnSO4 (p) ׃ CuSO4 (p)
5. If an electrolyte bridge is used at the interface between two conductors of the second kind, it is denoted by two lines: ZnSO4 (p) ׀׀ CuSO4 (p).
6. The components of one phase are written separated by commas:
Pt|Fe³+, Fe²+ ; Pt, H2 |HCl(p)
7. The equation of the electrode reaction is written so that the substances in the oxidizing form are located on the left, and in the reducing form on the right.
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REDOX PROCESSES AND REDOX SYSTEMS IN WINE
General information about redox processes
A substance is oxidized when it binds oxygen or gives up hydrogen; for example, when sulfur S is burned, sulfur dioxide SO 2 is formed, when sulfurous acid H 2 SO3 is oxidized, sulfuric acid H5SO4 is formed, and when hydrogen sulfide H 2 S is oxidized, sulfur S; when ferrous sulfate is oxidized in the presence of acid, ferric sulfate is formed
4FeSO„ + 2H 2 SO4 + 02 \u003d 2Fe2 (SO4) 3 + 2H20.
or during the decomposition of divalent sulfate into an anion SO ~ h, the Fe ++ cation is obtained
4Fe++ + 6SO "+ 4H+ + 02 = 4Fe+++ + + 6SO~~ + 2H 2 0,
or, reducing the anions not participating in the reaction, find
4Fe++ + 4H+ + 02 = 4Fe+++ + 2H20.
The latter reaction is identical in the case of oxidation of another ferrous salt; it does not depend on the nature of the anion. Therefore, the oxidation of a ferrous ion to a ferric ion is to increase its positive charge at the expense of the hydrogen ion, which loses its charge to form a hydrogen atom, which combines with oxygen to give water. As a result, this oxidation leads to an increase in the positive charge of the cation, or, equivalently, a decrease in the negative charge of the anion. For example, the oxidation of hydrogen sulfide H 2 S consists in the conversion of the sulfur ion S to sulfur (S). In fact, in both cases, there is a loss of negative electric charges or electrons.
In contrast, when x is reduced, the positive charge of the cation decreases or the negative charge of the anion increases. For example, in the previous reaction, one can say that there is a reduction of the H+ ion to atomic hydrogen H and that in the reverse direction of the reaction, the reduction of the Fe+++ ion to the Fe++ ion occurs. Thus, reduction is reduced to an increase in the number of electrons.
However, when it comes to the oxidation of organic molecules, the term "oxidation" retains its meaning of the transformation of one molecule into another or a combination of others richer in oxygen or less rich in hydrogen. Recovery is a reverse process, for example, the oxidation of alcohol CH3-CH2OH to aldehyde CH3-CHO, then to acetic acid CH3-COOH:
-2N +N,0-2N
CH3-CH2OH -> CH3-CHO -->
-> CH3-COOH.
The processes of oxidation of organic molecules in the cell, which are constantly encountered in biological chemistry and microbiology, occur most often by dehydrogenation. They are combined with reduction processes and constitute redox processes, for example, oxidation during alcoholic fermentation between glycerol and acetaldehyde, catalyzed by codehydrase and leading to alcohol:
CH2OH-CHOH-CHO + CH3-CHO + H20 - + CH2OH-CHOH-COOH + CH3-CH2OH.
Here we are talking about an irreversible redox process, which, however, can become reversible in the presence of a catalyst, as will be shown below. An example of an oxidation-reduction via electron exchange and reversible even in the absence of any catalyst is the equilibrium
Fe+++ + Cu+ Fe++ + Cu++.
It is the sum of two elementary reactions supplied by an electron
Fe++++e Fe++ and Cu+ Cu++ + e.
Such elementary reversible reactions constitute redox systems or redox systems.
They are of direct interest to oenology. Indeed, on the one hand, as has been shown, Fe++ and Cu+ ions are auto-oxidizable, i.e., they are oxidized directly, without a catalyst, by dissolved molecular oxygen, and the oxidized forms can re-oxidize other substances, therefore, these systems constitute oxidation catalysts. On the other hand, they are turbidity agents, which are always dangerous from the point of view of winemaking practice, and it is this circumstance that is closely related to their ability to move from one valence to another.
The general view of an ionized redox system, i.e., formed in solution by positively or negatively charged ions, can be expressed as follows:
Red \u003d 5 ± Ox + e (or ne).
A general view of an organic redox system in which the transition of a reduced to oxidized component occurs by releasing hydrogen, not electrons:
Red * Ox + H2.
Here Red and Ox represent molecules that do not have electric charges. But in the presence of a catalyst, for example, one of the redox systems shown above or some cell enzymes, H,2 is in equilibrium with its ions and constitutes a redox system of the first type
H2 *± 2H+ + 2e,
whence, summing the two reactions, we obtain the equilibrium
Red * Ox + 2H+ + 2e.
Thus, we come to a form similar to that of ionized systems that release electrons simultaneously with the exchange of hydrogen. Consequently, these systems, like the previous ones, are electroactive.
It is impossible to determine the absolute potential of the system; one can only measure the potential difference between two redox systems:
Redi + Ox2 * Red2 + Oxj.
The determination and measurement of the redox potential of a solution such as wine is based on this principle.
Classification of redox systems
In order to better consider the redox systems of wine and understand their role, it is advisable to use the Wurmser classification, which divides them into three groups:
1) directly electroactive substances, which in solution, even alone, directly exchange electrons with an inert electrode made of platinum, which accepts a well-defined potential. These isolated substances make up redox systems.
These include: a) heavy metal ions that make up the Cu++/Cu+ and Fe++/Fe+++ systems; b) many dyes, the so-called redox dyes, used for the colorimetric determination of the redox potential; c) riboflavin, or vitamin Bg, and dehydrogenases, in which it is included (yellow enzyme), participating in cellular respiration in grapes or in yeast in aerobiosis. These are auto-oxidizing systems, i.e., in the presence of oxygen, they take an oxidized form. No catalyst is required for their oxidation with oxygen;
2) substances with weak electrical activity that do not react or react weakly to a platinum electrode and do not independently provide conditions for equilibrium, but become electroactive when they are in solution in the presence of substances of the first group in very low concentrations and in this case give a certain potential . Substances of the second group react with the first, which catalyze their redox transformation and make irreversible systems reversible. Consequently, redox dyes make it possible to study the substances of this group, determine the normal potential for them, and classify them. Similarly, the presence of iron and copper ions in wine makes systems electroactive which, when isolated, are not redox systems.
These include: a) substances with an enol function with a double bond (-SON = COH-), in equilibrium with a di-ketone function (-CO-CO-), for example, vitamin C, or ascorbic acid, reductones, dihydroxymaleic-new acid; b) cytochromes, which play a major role in cellular respiration in both plants and animals;
3) electroactive substances in the presence of diastases. Their dehydrogenation is catalyzed by dehydrogenases, whose role is to ensure the transfer of hydrogen from one molecule to another. In general, these systems are given the electroactivity that they potentially possess by adding catalysts to the medium that provide redox transformations; then they create conditions for redox equilibrium and a certain potential.
These are systems lactic acid - pyruvic acid in the presence of an autolysate of lactic bacteria, which bring into redox equilibrium CH3-CHOH-COOH and CH3-CO-COOH - a system involved in lactic acid fermentation; ethanol - ethanal, which corresponds to the transition of aldehyde to alcohol in the process of alcoholic fermentation, or the butanediol - acetoin system. The latter systems are not relevant for the wine itself, although it can be assumed that the wine may contain dehydrases in the absence of microbial cells, but they are important for alcoholic or lactic acid fermentation, as well as for the finished wine containing living cells. They explain, for example, the reduction of ethanal in the presence of yeast or bacteria, a fact that has been known for a long time.
For all these oxidizing or reducing substances it is possible to determine the redox potential, normal or possible, for which the system is half oxidized and half reduced. This allows them to be classified in order of oxidizing or reducing strength. It is also possible to foresee in advance what form (oxidized or reduced) a given system is in a solution with a known redox potential; predict changes in dissolved oxygen content; determine the substances that are oxidized or reduced first. This issue is sufficiently covered in the section "The concept of redox potential".
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.
Redox potential (synonymous with redox potential; from Latin reductio - reduction and oxydatio - oxidation) - the potential that occurs on an inert (usually platinum) electrode immersed in a solution containing one or more reversible redox systems.
A reversible redox system (redox system) is a solution containing oxidized and reduced forms of substances, each of which is formed from the other through a reversible redox reaction.
The simplest redox systems include cations of the same metal of different valence, for example
or anions of the same composition, but of different valency, for example
In such systems, the redox process is carried out by the transfer of electrons from the reduced form to the oxidized one. Such redox systems include a number of respiratory enzymes containing hemin, for example, cytochromes. The redox potential of such systems can be calculated using the Peters formula:
where e- redox potential in volts, T - temperature on an absolute scale, n - the number of electrons lost by one molecule or ion of the reduced form during its transition to the oxidized form; [Ox] and - molar concentrations (more precisely, activities) of the oxidized and reduced forms, respectively; e0 is the normal redox potential of this system, equal to its redox potential, provided that =. The normal redox potentials of many redox systems can be found in physicochemical and biochemical reference books.
In many biological systems, redox reactions are carried out by transferring from the reduced form to the oxidized one not only electrons, but also an equal number of protons, for example 
The value of the redox potential of such systems is determined not only by the ratio [Ox] : = and pH = 0; the other values have the same values as in equation (1). The redox potential of biological systems, as a rule, is determined at pH=7, and the value e0-1.984·10-4·T·pH is denoted by e0. In this case, equation (2) takes the form:
Experimentally, the redox potential is determined potentiometrically (see Potentiometry). The redox potential of isolated cells and other biological objects is often measured colorimetrically using redox indicators (see). The magnitude of the redox potential is a measure of the redox or redox capacity of a given system. A redox system with a higher redox potential oxidizes a system with a lower redox potential. Thus, knowing the values of the redox potential of biological redox systems, it is possible to determine the direction and sequence of redox reactions in them. Knowing the redox potential also makes it possible to calculate the amount of energy that is released at a certain stage of the oxidative processes occurring in biological systems. See also biological oxidation.
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 \u003d C ox \u003d 1 mol / l or for metals C Me n + \u003d 1 mol / l, and for gaseous substances P \u003d 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. 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 + .