Redox systems. Redox potential Redox systems
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).
Page 4 of 8
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 valency 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 enzymes of the cell, 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. Therefore, 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".
There is a lot of data on the existence of a close relationship between the process of oxidation of D - lactate or an artificial substrate of ascorbate phenazine methasulfate and the transport of sugars, amino acids and some ions in vesicles artificially obtained from cell membranes. E. coli, Salmonella typhimurium, Pseudomonas putida, Proteus mirabilis, Bacillus megaterium, Bacillus subtilis, Micrococcus denitrificans, Mycobacterium phlei, Staphylococcus aureus.
Substrates that can be used with varying efficiency in redox systems also include α-glycerophosphate and much less often L-lactate, DL-α-hydroxybutyrate and even formate.
Such sugars as β - galactosides, galactose, arabinose, glucose - 6 - phosphate, gluconate and glucuronate, all natural amino acids, with the exception of glutamine (and, possibly, aspargine), arginine, methionine and ornithine, as well as cations are transported by this mechanism. K + and Rb + .
Although the mechanisms of such transport have not yet been fully resolved, it is most likely that protons are generated during the operation of the oxidative system. A membrane potential arises, most likely it serves as a driving force in the transfer of non-electrolytes.
Iron transport
E . coli K 12 has three specific systems for iron transport, and in all cases, outer membrane proteins play a central role in transport.
The Fe–citrate transport system is induced in the presence of citrate, and a new FecA protein receptor for Fe citrate appears in the outer membrane. More effective are systems that include microorganism-synthesized compounds that chelate iron. They secrete substances that convert iron into a soluble form. These substances are called siderophores. They bind iron ions into a complex and transport it in this form; we are talking mainly about low molecular weight water-soluble substances (with a molecular weight less than 1500), binding iron coordination bonds with high specificity and high affinity (stability constant of the order of 10 30). By their chemical nature, these can be phenolates or hydroxamates. Enterochelin belongs to the first; it has six phenolic hydroxy groups and is secreted by some enterobacteria. Once released into the environment, it binds iron, and the formed ferri enterochelin binds to a specific protein of the outer membrane, FepA, and then is absorbed by the cell. In the cell, iron is released as a result of enzymatic hydrolysis of ferri-enterochilin. In addition, this compound is able to cleave Fe 2+ even from the iron-containing proteins transferrin and lactoferrin. The synthesis of the FepA protein, as well as enterochelin, is repressed at a high content of dissolved iron in the medium.
outer membrane E . coli it also has a ferrichrome transport system. Mushrooms have the same transport system. Ferrichrome is classified as a hydroxamate siderophore. It is a cyclic hexapeptide that is formed by three glycine residues and three β-N-acetyl-L-β-hydroxyornithine residues. Ferrichrome forms a stable complex with ferric ions. E . coli , although it does not form ferrichrome itself, it has a very specific system of its transport, in which the outer membrane protein FhuA takes part. In the process of transport, iron is reduced and ferrichrome is modified (acetylated), as a result of which it loses its affinity for iron, and it is released into the cytoplasm.
A similar function is performed by ferrioxamines (in actinomycetes), mycobactins (in mycobacteria) and exochelins (also in mycobacteria).
A distinctive feature of redox reactions is the transfer of electrons between the reacting particles - ions, atoms, molecules and complexes, as a result of which the oxidation state of these particles changes, for example
Fe2+? e? = Fe3+.
Since electrons cannot accumulate in a solution, two processes must take place simultaneously - losses and gains, i.e., the process of oxidation of some and reduction of other particles. Thus, any redox reaction can always be represented as two half-reactions:
aOx1 + bRed2 = aRed1 + bOx2
The starting particle and the product of each half-reaction constitute a redox pair or system. In the above half-reactions, Red1 is conjugated to Ox1 and Ox2 is conjugated to Red1.
Not only particles in solution, but also electrodes can act as electron donors or acceptors. In this case, the redox reaction occurs at the electrode-solution interface and is called electrochemical.
Redox reactions, like all chemical reactions, are reversible to some extent. The direction of reactions is determined by the ratio of the electron-donor properties of the components of the system of one redox half-reaction and the electron-acceptor properties of the second (provided that the factors affecting the equilibrium shift are constant). The movement of electrons during a redox reaction leads to a potential. Thus, the potential, measured in volts, serves as a measure of the redox ability of a compound.
To quantify the oxidative (reductive) properties of the system, an electrode made of a chemically inert material is immersed in the solution. At the phase boundary, an electron exchange process occurs, leading to the emergence of a potential that is a function of the electron activity in the solution. The value of the potential is greater, the higher the oxidizing ability of the solution.
The absolute value of the potential of the system cannot be measured. However, if one of the redox systems is chosen as standard, then it becomes possible to measure the potential of any other redox system relative to it, regardless of the selected indifferent electrode. The H+/H2 system is chosen as standard, the potential of which is assumed to be zero.
Rice. one.
1. Platinum electrode.
2. Hydrogen gas supplied.
3. An acid solution (usually HCl) in which the concentration of H+ = 1 mol/l.
4. A water seal that prevents the ingress of oxygen from the air.
5. An electrolytic bridge (consisting of a concentrated solution of KCl) that allows you to connect the second half of the galvanic cell.
The potential of any redox system, measured under standard conditions against a hydrogen electrode, is called the standard potential (E0) of this system. The standard potential is considered to be positive if the system acts as an oxidizing agent and an oxidation half-reaction occurs on the hydrogen electrode:
or negative if the system plays the role of a reducing agent, and a reduction half-reaction occurs on the hydrogen electrode:
The absolute value of the standard potential characterizes the "strength" of the oxidizing agent or reducing agent.
The standard potential - a thermodynamic standardized value - is a very important physicochemical and analytical parameter that makes it possible to evaluate the direction of the corresponding reaction and calculate the activities of the reacting particles under equilibrium conditions.
To characterize the redox system under specific conditions, the concept of the real (formal) potential E0 "is used, which corresponds to the potential established at the electrode in this particular solution when the initial concentrations of the oxidized and reduced forms of potential-determining ions are equal to 1 mol / l and the fixed concentration of all other components solution.
From an analytical point of view, real potentials are more valuable than standard potentials, since the true behavior of the system is determined not by the standard, but by the real potential, and it is the latter that makes it possible to predict the occurrence of a redox reaction under specific conditions. The real potential of the system depends on the acidity, the presence of foreign ions in the solution, and can vary over a wide range.
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 the 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 + .
