Redox processes and redox systems in wine - oxidation and reduction processes in wines. Reversible redox system Reversible redox systems
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 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).
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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 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".
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).
such a process of interaction between two substances in which a reversible oxidation reaction of one substance occurs due to the reduction of another and a mixture of oxidized and reduced ions is formed in the medium, for example. - Fe"" and Fe", Sn" and Sn"", etc. The intensity level of the redox system is determined by the value of the redox potential Eh, which is expressed in volts, in relation to the potential of a normal hydrogen electrode.
The more positive the potential of the system, the more oxidizing properties it has. Potentials that are obtained in systems containing equal concentrations of oxidized and reduced ions, called. normal.
O. o.-v. With. according to the magnitude of normal potentials, they can be arranged in a row, 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 matter in sedimentary rocks, etc.
Substance equivalent or Equivalent is a real or conditional particle that can attach, release, or otherwise be equivalent to a hydrogen cation in ion exchange reactions or an electron in redox reactions.
For example, in react:
NaOH + HCl \u003d NaCl + H 2 O
the equivalent will be a real particle - Na + ion, in the reaction
the imaginary particle ½Zn(OH) 2 will be the equivalent.
Substance equivalent is also often used to mean number of substance equivalents or equivalent amount of substance- the number of moles of a substance equivalent to one mole of hydrogen cations in the reaction under consideration.
[edit] Equivalent mass
Equivalent mass is the mass of one equivalent of the given substance.
[edit] Equivalent molar mass of a substance
Molar mass equivalents are usually denoted as or . The ratio of the equivalent molar mass of a substance to its own molar mass is called equivalence factor(usually denoted as ).
The molar mass of the equivalents of a substance is the mass of one mole of equivalents, equal to the product of the equivalence factor by the molar mass of this substance.
M eq = f eq ×M
[edit] Equivalence factor
The ratio of the equivalent molar mass to its own molar mass is called equivalence factor(usually denoted as ).
[edit] Equivalence number
Equivalence number z is a small positive integer equal to the number of equivalents of some substance contained in 1 mole of this substance. The equivalence factor is related to the equivalence number z the following relation: =1/z.
For example, in react:
Zn(OH) 2 + 2HCl = ZnCl 2 + 2H 2 O
The equivalent is the particle ½Zn(OH) 2 . The number ½ is equivalence factor, z in this case is 2
* - for inert gases Z = 1
The equivalence factor helps to formulate the law of equivalence.
[edit] The law of equivalents
As a result of the work of I. V. Richter (1792-1800), the law of equivalents was discovered:
§ All substances react in equivalent ratios.
§ formula expressing the Law of Equivalents: m 1 E 2 \u003d m 2 E 1
§ Electrochemical equivalent- the amount of substance that should be released on the electrode, according to Faraday's law, when a unit of electricity passes through the electrolyte:
§ where is the Faraday constant.
§ Faraday constant, is a physical constant that determines the relationship between the electrochemical and physical properties of a substance.
§ Faraday's constant is C mol −1 .
§ The Faraday constant is included as a constant in Faraday's second law(the law of electrolysis).
§ Numerically, the Faraday constant is equal to the electric charge, during the passage of which through the electrolyte on the electrode, (1 / z) mol of substance A is released in the formula:
where: 
is the number of electrons involved in the reaction.
§ For the Faraday constant, the following relation is true:
§ where is the elementary charge, and is the Avogadro number.
Isotopes(from other Greek ισος - "equal", "same", and τόπος - "place") - varieties of atoms (and nuclei) of the same chemical element with a different number of neutrons in the nucleus. The name is due to the fact that the isotopes are in the same place (in the same cell) of the periodic table. The chemical properties of an atom depend practically only on the structure of the electron shell, which, in turn, is determined mainly by the charge of the nucleus Z(that is, the number of protons in it) and almost does not depend on its mass number A(that is, the total number of protons Z and neutrons N). All isotopes of the same element have the same nuclear charge, differing only in the number of neutrons. Usually an isotope is denoted by the symbol of the chemical element to which it belongs, with the addition of an upper left index indicating the mass number (for example, 12 C, 222 Rn). You can also write the name of the element with a hyphenated mass number (for example, carbon-12, radon-222). Some isotopes have traditional proper names (for example, deuterium, actinon).
An example of isotopes: 16 8 O, 17 8 O, 18 8 O - three stable isotopes of oxygen.
[edit] Terminology
The main position of IUPAC is that the correct singular term for atoms (or nuclei) of the same chemical element with the same atomic mass is nuclide, and the term isotopes can be used to designate a set of nuclides of one element. Term isotopes was proposed and used initially in the plural, since at least two types of atoms are needed for comparison. In the future, the use of the term in the singular became widely used in practice - isotope. In addition, the term in the plural is often used to refer to any set of nuclides, and not just one element, which is also incorrect. At present, the positions of international scientific organizations have not been brought to uniformity and the term isotope continues to be widely used, including in the official materials of various divisions of IUPAC and IUPAP. This is one of the examples of how the meaning of the term, originally embedded in it, ceases to correspond to the concept for which this term is used (another textbook example is the atom, which, contrary to the name, is not indivisible).
[edit]History of the discovery of isotopes
The first evidence that substances having the same chemical behavior can have different physical properties came from the study of radioactive transformations of atoms of heavy elements. In 1906-07, it became clear that the product of the radioactive decay of uranium, ionium, and the product of the radioactive disintegrator, radiothorium, have the same chemical properties as thorium, but differ from it in atomic mass and the characteristics of radioactive decay. It was later found that all three products have the same optical and X-ray spectra. Such substances, identical in chemical properties, but different in the mass of atoms and some physical properties, at the suggestion of the English scientist F. Soddy, began to be called isotopes.
[edit] Isotopes in nature
It is believed that the isotopic composition of elements on Earth is the same in all materials. Some physical processes in nature lead to a violation of the isotopic composition of elements (natural fractionation isotopes characteristic of light elements, as well as isotopic shifts during the decay of natural long-lived isotopes). Gradual accumulation in minerals of nuclei - decay products of some long-lived nuclides is used in nuclear geochronology.
[edit]Human uses of isotopes
In technological activities, people have learned to change the isotopic composition of elements to obtain any specific properties of materials. For example, 235 U is capable of a thermal neutron fission chain reaction and can be used as fuel for nuclear reactors or nuclear weapons. However, natural uranium contains only 0.72% of this nuclide, while a chain reaction is practically feasible only if the 235 U content is at least 3%. Due to the closeness of the physicochemical properties of isotopes of heavy elements, the procedure for isotope enrichment of uranium is an extremely complex technological task, which is accessible only to a dozen countries in the world. In many branches of science and technology (for example, in radioimmunoassay), isotope labels are used.
Dissociation constant- a kind of equilibrium constant that indicates the tendency of a large object to dissociate (separate) in a reversible way into small objects, such as when a complex breaks down into its constituent molecules, or when a salt separates into ions in an aqueous solution. The dissociation constant is usually denoted Kd and inverse to the association constant. In the case of salts, the dissociation constant is sometimes called the ionization constant.
In a general reaction
where is the complex A x B y breaks down into x units A and y units B, the dissociation constant is defined as follows:
where [A], [B] and are the concentrations of A, B and the complex A x B y, respectively.
[edit] Definition
The electrolytic dissociation of weak electrolytes, according to the Arrhenius theory, is a reversible reaction, that is, it can be schematically represented by the equations (for monovalent ions:):
KA ↔ K + + A - ,
§ KA - undissociated compound;
§ K + - cation;
§ A − - anion.
The equilibrium constant of such a reaction can be expressed by the equation:
 ,
| (1) |
§ - concentration of undissociated compound in solution;
§ - concentration of cations in solution;
§ - concentration of anions in solution.
The equilibrium constant in relation to the dissociation reaction is called dissociation constant.
[edit] Dissociation of electrolytes with polyvalent ions
In the case of dissociation of electrolytes with multivalent ions, dissociation occurs in steps, and each step has its own value of the dissociation constant.
Example: Dissociation of a polybasic (boric) acid [ source not specified 332 days] :
Stage I: H 3 BO 3 ↔ H + + H 2 BO 3 -,
Stage II: H 2 BO 3 - ↔ H + + HBO 3 2 - ,
Stage III: HBO 3 2− ↔ H + + BO 3 3− ,
The first degree of dissociation for such electrolytes is always much greater than the subsequent ones, which means that the dissociation of such compounds proceeds mainly through the first stage.
[edit] Relationship between dissociation constant and degree of dissociation
Based on the definition of the degree of dissociation, for the KA electrolyte in the dissociation reaction = = α·c, = c - α·c = c·(1 - α), where α is the degree of dissociation of the electrolyte.
| , | (2) |
This expression is called the Ostwald dilution law. For very small α (α<<1) K=cα² и
thus, with an increase in the electrolyte concentration, the degree of dissociation decreases, and with a decrease, it increases. The relationship between the dissociation constant and the degree of dissociation is described in more detail in the article Ostwald's Dilution Law.
[edit] The difference between the experimental results and the Arrhenius model, the derivation of the dissociation constant through activities
The above calculations are based on the Arrhenius theory, which is too rough and does not take into account the factors of the electrostatic interaction of ions. Deviations from the ideal state in electrolyte solutions occur at very low concentrations, since the interionic forces are inversely proportional to square distances between ion centers, while intermolecular forces are inversely proportional seventh degree distances, that is, interionic forces, even in dilute solutions, turn out to be much greater than intermolecular ones.
Lewis showed that simple equations can be preserved for real solutions (see above) if instead of ion concentrations we introduce its function, the so-called activity. Activity (a) is related to concentration (c) through a correction factor γ called the activity factor:
a = γ c
Thus, the expression for the equilibrium constant, according to Arrhenius described by equation (1), according to Lewis will look like:
§ 
;
§ 
;
In the Lewis theory, the relationship between the constant and the degree of dissociation (in the Arrhenius theory written by equation (2) is expressed by the relationship:
If there are no other influences that deviate the solution from the ideal state, then the non-dissociated molecules behave like ideal gases and γ KA = 1, and the true expression of the Ostwald dilution law will take the form:
§ is the average activity coefficient of the electrolyte.
For c→0 and γ→1, the above equation of the Ostwald dilution law takes the form (2). The more the electrolyte dissociates, the faster the value of the activity coefficient γ deviates from unity, and the faster the classical dilution law is violated.
[edit] Dissociation constant of strong electrolytes
Strong electrolytes dissociate almost completely (the reaction is irreversible), therefore, the denominator of the expression for the dissociation constant is zero, and the whole expression tends to infinity. Thus, for strong electrolytes, the term "dissociation constant" is meaningless.
[edit] Calculation examples
[edit] Water dissociation
Water is a weak electrolyte that dissociates according to the equation
The dissociation constant of water at 25 °C is
Considering that in most solutions water is in molecular form (the concentration of H + and OH - ions is low), and given that the molar mass of water is 18.0153 g / mol, and the density at a temperature of 25 ° C is 997.07 g / l, pure water corresponds to the concentration = 55.346 mol/l. Therefore, the previous equation can be rewritten as
The application of the approximate formula gives an error of about 15%:
Based on the found value of the degree of dissociation, we find the pH of the solution:
Degree of dissociation- a value characterizing the state of equilibrium in the dissociation reaction in homogeneous (homogeneous) systems.
The degree of dissociation α is equal to the ratio of the number of dissociated molecules n to the sum n + N, where N is the number of undissociated molecules. Often α is expressed as a percentage. The degree of dissociation depends both on the nature of the dissolved electrolyte and on the concentration of the solution.
[edit] Example
For acetic acid CH 3 COOH, the value of α is 4% (in a 0.01M solution). This means that in an aqueous solution of an acid, only 4 out of every 100 molecules are dissociated, that is, they are in the form of H + and CH 3 COO − ions, while the remaining 96 molecules are not dissociated.
[edit] Definition methods
§ according to the electrical conductivity of the solution
§ to lower the freezing point
[edit] Imaginary degree of dissociation
Since strong electrolytes dissociate almost completely, one would expect for them an isotonic coefficient equal to the number of ions (or polarized atoms) in the formula unit (molecule). However, in reality, this coefficient is always less than that determined by the formula. For example, the isotonic coefficient for a 0.05 mol NaCl solution is 1.9 instead of 2.0 (for a magnesium sulfate solution of the same concentration, i= 1.3). This is explained by the theory of strong electrolytes, developed in 1923 by P. Debye and E. Hückel: the movement of ions in solution is hindered by the formed solvation shell. In addition, ions interact with each other: oppositely charged ones attract, and likewise charged ones repel; the forces of mutual attraction lead to the formation of groups of ions moving through the solution together. Such groups are called ion associates or ion pairs. Accordingly, the solution behaves as if it contains fewer particles than it really is, because the freedom of their movement is limited. The most obvious example concerns the electrical conductivity of solutions λ , which increases with dilution of the solution. Through the ratio of real electrical conductivity to that at infinite dilution, determine imaginary degree of dissociation strong electrolytes, also referred to as α :
,
where nimg- imaginary, and n disslv. is the actual number of particles in the solution.
