The history of the discovery of such an interesting phenomenon in chemistry as electrolytic dissociation began in 1887, when the Swedish chemist Svante Arennius, while studying the electrical conductivity of aqueous solutions, suggested that in such solutions substances could disintegrate into charged particles - ions. These ions are in motion, moving to the electrodes, both the positively charged cathode and the negatively charged anode. This decomposition process is called electrolytic dissociation; it is the reason for the appearance of electric current in solutions.
Electrolytic dissociation theory
The classical theory of electrolytic dissociation, developed by the discoverer S. Ahrennius together with W. Oswald, first of all, assumed that the disintegration of molecules into ions (dissociation itself) occurs under the influence of an electric current. Subsequently, it turned out that this was not entirely true, since the existence of ions in aqueous solutions was revealed, regardless of whether current passed through them or not. Then Svante Ahrennius formed a new theory, its essence is that electrolytes spontaneously disintegrate into ions under the influence of a solvent. And the presence of ions creates ideal conditions for electrical conductivity in the solution.
This is roughly what electrolytic dissociation looks like schematically.
The great importance of electrolytic dissociation in solutions lies in the fact that it allows us to describe the properties of acids, bases and salts, and further we will dwell on this in detail
Electrolytic dissociation of acids
N 3 PO 4 ⇄ N + N 2 PO- 4 (first stage)
N 2 PO 4 ⇄ N + NPO 2 - 4 (second stage)
N 2 PO 4 ⇄ N+ PO Z - 4 (third stage)
This is what the chemical equations for electrolytic dissociation of acids look like. The example shows the electrolytic dissociation of phosphoric acid H 3 PO 4, which decomposes into hydrogen H (cation) and anode ions. Moreover, the dissociation of many basic acids occurs, as a rule, only in the first step.
Electrolytic dissociation of bases
Bases differ from acids in that when they dissociate, hydroxide ions are formed as cations.
Example of a base chemical dissociation equation
KOH ⇄ K + OH-; NH 4 OH ⇄ NH+ 4 + OH-
Bases that dissolve in water are called alkalis, there are not so many of them, mainly alkali and alkaline earth bases, such as LiOH, NaOH, KOH, RbOH, CsOH, FrOH and Ca(OH) 2, Sr(OH) 2 , Ba(OH) 2 , Ra(OH) 2
Electrolytic dissociation of salts
During the electrolytic dissociation of salts, metals are formed as cations, as well as the ammonium cation NH 4, and acid residues become anions.
(NH 4) 2 SO 4 ⇄ 2NH+ 4 + SO 2 - 4; Na 3 PO 4 ⇄ 3Na + PO 3- 4
An example of an equation for the electrolytic dissociation of salts.
Electrolytic dissociation, video
And finally, an educational video on the topic of our article.
MINISTRY OF EDUCATION AND SCIENCE OF THE RF
GOU VPO TVER STATE UNIVERSITY
FACULTY OF CHEMICAL TECHNOLOGY
Abstract on the history of chemistry
Origins of the theory of electrolytic dissociation. The emergence of solution theory.
Completed:
Ilyina N.V.
Selina T.Yu.
TVER 2012
TABLE OF CONTENTS
Introduction 3
1 Origins of the theory of electrolytic dissociation
1.1 The founder of TED is S. Arrhenius 4
1.2 New ideas about salts, acids and bases 6
1.3 Further development of TED 7
2 The struggle for recognition of TED
2.1 TED and D.I.Mendeleev 11
2.2 TED in Russia 13
3 Theory of solutions
3.1 Chemical theory of solutions D.I. Mendeleeva 16
3.2 Van't Hoff's osmotic theory 18
3.3 Non-aqueous solutions 19
3.4 Further development of solution theory 24
Conclusion 25
References 26
INTRODUCTION
A load of gifts for two guys
Ion put it on his back:
For Katya, He brings his own plus,
For Anya, He carries his own minus.
Even at the dawn of the study of electrical phenomena, scientists noticed that not only metals, but also solutions can conduct current. But not all of them. Thus, aqueous solutions of table salt and other salts, solutions of strong acids and alkalis conduct current well. Solutions of acetic acid, carbon dioxide and sulfur dioxide conduct it much worse. But solutions of alcohol, sugar and most other organic compounds do not conduct electricity at all.
The English physicist Michael Faraday, back in the 30s of the 19th century, while studying the laws of the passage of electric current through solutions, introduced the terms “electrolyte”, “electrolysis”, “ion”, “cation”, “anion”. An electrolyte is a substance whose solution conducts electric current. This happens as a result of the movement of charged particles – ions – in the solution. However, the reason for the appearance of charged particles in solutions was completely unclear.
In 1887, the Swedish physicist and chemist Svante Arrhenius, studying the electrical conductivity of aqueous solutions, suggested that in such solutions substances disintegrate into charged particles - ions, which can move to the electrodes - a negatively charged cathode and a positively charged anode. This is the reason for the electric current in solutions. This process is called electrolytic dissociation (literal translation - splitting, decomposition under the influence of electricity). This name also suggests that dissociation occurs under the influence of an electric current. Further research showed that this is not so: ions are only charge carriers in a solution and exist in it regardless of whether current passes through the solution or not.
Many scientists - contemporaries of Arrhenius, initially did not accept his theory. Many of them at that time did not yet have a clear understanding of how ions differ from neutral atoms. As a result, Arrhenius' dissertation received a number of negative reviews. Among the most irreconcilable opponents of Arrhenius was D.I. Mendeleev, who created the “chemical” theory of solutions, in contrast to the “physical” theory of Arrhenius. However, the subsequent successes of the new theory were so impressive, and its recognition (even if not universal) was so enthusiastic that against this background all doubts were forgotten. The theory of electrolytic dissociation convincingly explained many facts that had been known for a long time and yet remained unclear.
By the mid-19th century, scientists' interest in the nature of solutions increased. It is becoming more and more obvious that without knowledge of the nature of solutions it is impossible to study many phenomena and penetrate into the essence of various production processes. In connection with the development of chemical production, there is an urgent need to study the properties and composition of various solutions. Many scientists of that period viewed solutions as mechanical mixtures of certain solute compounds with solvent molecules. The theory of solutions received the greatest development in the fundamental works of D.I. Mendeleev.
- ORIGINS OF THE THEORY OF ELECTROLYTIC DISSOCIATION
1.1 THE FOUNDER OF THE THEORY OF ELECTROLYTIC DISSOCIATION - S. ARRENIUS
The history of the emergence of the theory of electrolytic dissociation is associated with the name of the Swedish physicist and chemist Svante Arrhenius (1859-1927). He has written 200 scientific papers in the fields of chemistry, physics, geophysics, meteorology, biology, and physiology.
ARRHENIUS, Svante August. Swedish physical chemist Svante August Arrhenius was born on the Wijk estate, near Uppsala. He was the second son of Caroline Christina (Thunberg) and Svante Gustav Arrhenius, the estate manager. Arrhenius's ancestors were farmers. A year after the birth of their son, the family moved to Uppsala, where S.G. Arrhenius joined the board of inspectors of Uppsala University. While attending the cathedral school in Uppsala, Arrhenius showed exceptional abilities in biology, physics and mathematics.
In 1876 Arrhenius entered Uppsala University, where he studied physics, chemistry and mathematics. In 1878 He was awarded a Bachelor of Science degree. However, he continued to study physics at Uppsala University for the next three years, and in 1881. went to Stockholm, to the Royal Swedish Academy of Sciences, to continue research in the field of electricity under the leadership of Erik Edlund.
In 1884 Arrhenius is pursuing his doctorate at Uppsala University. And in 1886 Arrhenius becomes a fellow of the Royal Swedish Academy of Sciences, which allowed him to work and conduct research abroad. Over the next five years he worked in Riga with Ostwald, in Würzburg with Friedrich Kohlrausch (here he met Walter Nernst), at the University of Graz with Ludwig Boltzmann and in Amsterdam with Jacob van't Hoff. Returning to Stockholm in 1891, Arrhenius began lecturing on physics at Stockholm University, and in 1895. receives a professorship there. In 1897 he holds the post of rector of the university.
In 1903 Arrhenius was awarded the Nobel Prize in Chemistry "in recognition of the special significance of his theory of electrolytic dissociation for the development of chemistry." Speaking on behalf of the Royal Swedish Academy of Sciences, H.R. Terneblad emphasized that Arrhenius’s ion theory laid the qualitative foundation for electrochemistry, “allowing a mathematical approach to be applied to it.” “One of the most important results of Arrhenius’s theory,” said Terneblad, “is the completion of the colossal generalization for which the first Nobel Prize in chemistry was awarded to van’t Hoff.”
Arrhenius received many awards and titles. Among them: the Davy Medal of the Royal Society of London (1902), the first Willard Gibbs Medal of the American Chemical Society (1911), the Faraday Medal of the British Chemical Society (1914). He was a member of the Royal Swedish Academy of Sciences, a foreign member of the Royal Society of London and the German Chemical Society. Arrhenius was awarded honorary degrees from many universities, including Birmingham, Edinburgh, Heidelberg, Leipzig, Oxford and Cambridge.
It is interesting that the idea that became the basis of this theory arose on the basis of experiments carried out to solve a completely different problem.
According to Yu.I. Soloviev, “while still a student at Uppsala University, S. Arrhenius, listening to lectures by his teacher Professor P.T. Kleve, learned that it is impossible to determine the molecular weight of substances that, like cane sugar, do not go into a gaseous state. In order to bring “great benefit” to chemistry, the young scientist decides to determine the electrical conductivity of salts in solutions containing, along with water, a large number of non-electrolytes. In doing so, he proceeded from the principle that the greater the molecular weight of the solvent, the greater the resistance of the electrolyte solution. This was the original work plan.
But as a result of his first observations, S. Arrhenius loses interest in the intended topic. He is carried away by a new thought. What happens to an electrolyte molecule in solution? The young scientist realized that a successful solution to this issue would shed bright light on the dark area of \u200b\u200bsolutions. So, instead of determining the molecular mass of a dissolved non-electrolyte, S. Arrhenius begins to intensively study the state of the electrolyte molecule in solution.
Work in a new direction soon yielded excellent results. The data obtained by measuring the electrical conductivity of aqueous solutions of electrolytes of various concentrations allowed S. Arrhenius to make a bold conclusion: electrolyte molecules dissociate into ions without the influence of current, and the degree of dissociation increases with dilution. As it now seems to us, this was a seemingly obvious and simple conclusion from the experimental data. But it was not at all simple for S. Arrhenius, for this conclusion destroyed the solid, “granite-like” traditional ideas about the state of the molecules of salts, acids and bases in solution.”
S. Arrhenius was afraid to even express his thoughts in a categorical form, complicating the main provisions of his theory with not very clear terms. The fact is that he had no idea where the energy needed to break electrolyte molecules (or crystals) into ions comes from. And this energy is considerable. If, for example, you take one mole of table salt and “scatter” it into ions, you will need 800 kJ of energy. Such a large energy is needed in order to overcome the Coulomb attraction between oppositely charged ions in crystals or electrolyte molecules.
And yet, the experiments of S. Arrhenius, and other scientists, persistently led to the conclusion that ions do exist in electrolyte solutions, regardless of whether an electric field acts on them or not. And this gave S. Arrhenius confidence that he was right. Arrhenius' theory was greeted differently: some scientists were enthusiastic, others were hostile. This is understandable. It had many advantages, but no less disadvantages.
1.2 NEW INSIGHTS ABOUT SALTS, ACIDS AND BASES
S. Arrhenius created virtually new ideas about acids, bases and salts. He considered an acid to be a compound that dissociates in an aqueous solution with the abstraction of hydrogen ions. For example:
HCl « H + + Cl -
H 2 SO 4 « 2H + + SO 4 2-
From here it became clear why acids have a number of common properties. The sour taste, the same color of the indicator, the release of hydrogen under the action of active metals - hydrogen ions formed during the dissociation of acids were responsible for all these properties.
He considered the base to be a compound that dissociates in an aqueous solution to form hydroxide ions:
KOH « K + + OH -
Ca(OH) 2 « Ca 2+ + 2OH -
Then the general properties of bases became clear. Bitter taste, soapy feeling, identical reaction to indicators - this is all the “handiwork” of OH - ions.
S. Arrhenius explained the different electrical conductivities of acids and bases by their different abilities to dissociate. He called well-dissociating acids or bases, which produce many ions in solutions, strong, and poorly dissociating ones, which form few ions, weak. To characterize the “strength” of electrolytes, S. Arrhenius introduced a new concept - the degree of electrolytic dissociation. Now it has become clear why a solution of acetic acid conducts electric current worse than a solution of sulfuric acid of the same concentration.
1.3 FURTHER DEVELOPMENT OF THE THEORY OF ELECTROLYTIC DISSOCIATION
Of great importance for the further development of the theory of dissociation was the famous work of Van’t Hoff “Chemical equilibrium in systems of gases and dilute solutions” (1885), in which it was established that the actual decrease in the melting point, vapor pressure and osmotic pressure of salts, acids and bases is less than than calculated theoretically according to Raoult's law. These inconsistencies confirmed the provisions of the dissociation theory, according to which the electrolyte in an aqueous solution disintegrates into freely moving ions.
VANT-HOFF (van"t Hoff), Jakob Henrik. Dutch chemist Jakob Henrik Van't Hoff was born in Rotterdam, in the family of the doctor Jakob Henrik Van't Hoff. At the insistence of his parents, Van't Hoff began studying engineering at the Polytechnic School in Delft. Van't Hoff completed a three-year training program in two years and passed the final exam better than anyone else.
In 1871 Van't Hoff became a student at the Faculty of Science and Mathematics at Leiden University. The following year he moved to the University of Bonn to study chemistry under Friedrich August Kekule. Two years later Van't Hoff continued his studies at the University of Paris, where he completed his dissertation. Returning to the Netherlands, he presented her for defense at the University of Utrecht.
Van't Hoff extended the concept of the tetrahedral carbon atom to compounds containing carbon-carbon double bonds (two tetrahedra sharing an edge) and triple bonds (two tetrahedra sharing a common edge). Van't Hoff did not dare to present his theory as a doctoral dissertation. Instead, he wrote a dissertation on cyanoacetic and malonic acids and received a doctorate in chemistry in 1874.
Van't Hoff's scientific career progressed slowly. At first he had to give private lessons in chemistry and physics, and only in 1876 did he receive a position as lecturer in physics at the Royal Veterinary School in Utrecht. The following year he becomes lecturer (and later professor) of theoretical and physical chemistry at the University of Amsterdam. Here, over the next 18 years, he gave five lectures every week on organic chemistry and one lecture on mineralogy, crystallography, geology and paleontology, and also directed a chemical laboratory.
In 1901, Van't Hoff became the first winner of the Nobel Prize in Chemistry, which was awarded to him "in recognition of the enormous importance of his discovery of the laws of chemical dynamics and osmotic pressure in solutions." Representing Van't Hoff on behalf of the Royal Swedish Academy of Sciences, S.T. Odner called the scientist the founder of stereochemistry and one of the creators of the doctrine of chemical dynamics, and also emphasized that van’t Hoff’s research “made a significant contribution to the remarkable achievements of physical chemistry.”
In addition to the Nobel Prize, Van't Hoff was awarded the Davy Medal of the Royal Society of London (1893) and the Helmholtz Medal of the Prussian Academy of Sciences (1911). He was a member of the Royal Netherlands and Prussian Academies of Sciences, the British and American Chemical Societies, the American National Academy of Sciences and the French Academy of Sciences. Van't Hoff was awarded honorary degrees from the University of Chicago, Harvard and Yale.
In the spring of 1887, Arrhenius worked in Würzburg with F. Kohlrausch. “Shortly before I left Würzburg (March 1887),” Arrhenius recalled, “I received Van’t Hoff’s work published by the Swedish Academy of Sciences. I looked through it one evening after finishing my daily work at the institute. It immediately became clear to me that the deviation of electrolytes in an aqueous solution from the van’t Hoff-Raoult laws on lowering the freezing point is the most compelling evidence of their disintegration into ions. Now I had before me two ways to calculate the degree of dissociation: on the one hand, by lowering the freezing point, on the other, from conductivity. Both of them gave the same result in the vast majority of cases, and I could talk openly about the dissociation of electrolytes."
In a letter to Van't Hoff in March 1887, the Swedish scientist wrote: “Both theories are still at the very beginning of their development, and I most earnestly hope that in the near future not one, but several bridges will be thrown between both areas.” And so it happened. In 1887, the famous article by Arrhenius “On the dissociation of substances dissolved in water” appeared in the first volume of the “Journal of Physical Chemistry” organized by W. Ostwald. Here the author has already boldly and openly stated that molecules of electrolytes (salts, acids, bases) disintegrate in solution into electrically charged ions.
After 1887, the studies of S. Arrhenius, W. Ostwald, N. Nernst, M. Leblanc and other scientists not only confirmed the validity of the basic provisions of the theory of electrolytic dissociation, but also significantly expanded the number of individual facts that can be substantiated by the theory.
In 1888, Walter Friedrich Nernst (1864-1941), professor of physical chemistry in Göttingen and Berlin, winner of the 1920 Nobel Prize in Chemistry for the discovery of the third law of thermodynamics, comparing the rate of ion diffusion with the speed of ion movement during electrolysis, showed that these numbers coincide . In 1889, based on the theory of osmotic pressure and the theory of electrolytic dissociation, Nernst developed the osmotic theory of the generation of galvanic current.
In 1884-1886, W. Ostwald managed to find a lot of data confirming the parallelism between the chemical activity of substances and their electrical conductivity. In 1888, he proposed a method for determining the basicity of acids by the electrical conductivity of their solutions and showed that the rate of a chemical reaction in solutions depends only on the dissociated part of the solute.
OSTWALD, Friedrich Wilhelm. German chemist Friedrich Wilhelm Ostwald was born in Riga (Latvia). He was the second son of Gottfried Ostwald, a skilled cooper, and Elisabeth (Leukel) Ostwald. While studying at the Riga real gymnasium, Ostwald proved himself to be a good student with an unusually wide range of interests. He was interested in physics, chemistry, literature and drawing, and also played the viola and piano. Despite the fact that his father advised him to study engineering, Ostwald became interested in chemistry and in 1872 became a student at the Faculty of Chemistry at Dorpat (now Tartu) University. Four years later, he received his bachelor's degree and remained in Dorpat for graduate school, simultaneously holding the position of privat-docent (freelance teacher).
In 1878 he was awarded a doctorate for his thesis on the optical refractive index of acid-base reactions. Working as an assistant to the physicist Arthur von Oettingen and teaching physics and chemistry at a local school, Ostwald continued to study the application of physical characteristics to the analysis of chemical reactions. In 1881 he was elected professor of chemistry at the Riga Polytechnic Institute. In subsequent years, he wrote several textbooks that played an important role in establishing physical chemistry as an independent discipline.
In 1884, Ostwald received the text of Svante Arrhenius's controversial doctoral dissertation, which was presented for defense at Uppsala University. In his dissertation, Arrhenius proposed a theory explaining the dissociation of acids and bases in aqueous solutions into electrically charged ions. Since the prevailing belief at the time was that oppositely charged particles could not coexist in a solution, Arrhenius's work received a low rating at Uppsala University. Ostwald, however, found his ideas worthy of attention and immediately applied them to verify the results of his own studies of acid affinities. “Using a resistance store borrowed for a few days at the telegraph (they couldn’t do without it for much longer) ... I soon carried out experiments with all the acids at hand, which other researchers provided me with,” Ostwald later recalled, “with increasing excitement I discovered that that the results one after another confirmed the predictions and expectations.”
In 1909, Ostwald was awarded the Nobel Prize in Chemistry “in recognition of his work in catalysis and for his studies of the fundamental principles of the control of chemical equilibria and reaction rates.” Presenting on behalf of the Royal Swedish Academy of Sciences, Hans Hildebrand pointed out the value of Ostwald's discoveries not only for the development of theory, but also for their practical applications, such as the production of sulfuric acid and the synthesis of indigo dyes. Hildebrand also predicted that the chemistry of catalysis would contribute greatly to understanding enzyme function. In the last years of his life, Ostwald became involved in various educational, cultural and reformist movements, including the Internationalist, pacifist and natural resource conservation movements. He was active in numerous international scientific societies, including the International Atomic Weights Commission and the International Association of Chemical Societies. Ostwald was also involved in issues of public education and training of scientists.
Also in 1888, W. Ostwald found a pattern connecting the degree of dissociation of an electrolyte with its concentration. In 1884-1886, he established that the electrical conductivity of acids increases with dilution - asymptotically approaching a certain limiting value. He found that for solutions of weak acids (succinic, etc.) and bases, the increase in molecular electrical conductivity with dilution is much more noticeable than for strong acids, such as sulfuric, etc. In one of the works written in 1888, W. Ostwald gave a mathematical formulation of the law of dilution. He compared the electrical conductivity of the electrolyte with the limit for an infinitely large dilution.
The new law became fundamental for the chemistry of aqueous solutions. “W. Ostwald’s law of dilution,” writes Yu.I. Soloviev, confirmed the theory of electrolytic dissociation and made it possible to determine the dependence of the degree of dissociation of electrolyte molecules on the concentration of the solution. Subsequently, this law was subjected to repeated testing. It was found that it is not applicable for strong electrolytes and concentrated solutions. It took numerous studies by scientists of the late 19th and early 20th centuries to explain the reason why strong electrolytes did not obey the law of dilution. The fruitfulness of the theory of electrolytic dissociation was especially clearly manifested in the fact that it was successfully used to explain the mechanism of many chemical reactions and the nature of various compounds, for example complex ones.”
The theory of electrolytic dissociation was able to combine both the theory of solutions and the electrochemical theory. As Arrhenius predicted, both streams merged into a single one. “After the founding of the mechanical theory of heat,” Ostwald wrote in 1889, “in the physical sciences there was not a single series of ideas as comprehensive as the solution theory of Van’t Hoff and Arrhenius.”
The theory of electrolytic dissociation was subsequently improved thanks to the work of, first of all, N. Bjerrum, P. Debye and E. Hückel. They developed the ideas previously expressed by I. Van Laar that the unusual behavior of strong electrolytes can be explained by the action of Coulomb forces.
- THE STRUGGLE FOR RECOGNITION OF THE THEORY OF ELECTROLYTIC DISSOCIATION
Almost all the major physical chemists of that time were involved in the struggle. This led to the study of a number of new issues put forward in the process of discussing controversial issues. The discussion was strictly scientific, principled, and creative in nature. Her motto was - no unfounded statements and empty declarations. Scientists acted armed with experimental facts and new hypotheses. This explains the fruitfulness of such a struggle, which gave impetus to the movement of science forward.
In the course of a long debate, both the chemical and physical theories of solutions were posed with the most complex, confusing questions that required a radical solution. This forced the participants in the debate to think deeply about individual provisions and conduct new experiments.
What is the reason that the theory of electrolytic dissociation has caused almost total opposition from chemists? The main reason was that the new theory was in deep contradiction with the then dominant theoretical concepts and experimental data. Most chemists have an “unshakable” belief that the decomposition of an electrolyte in a solution occurs only under the influence of an electric current. Also widely accepted was the unproven fact that, for example, an aqueous solution of sodium chloride contains only its molecules. They thought so because when the solution evaporates, the same sodium chloride that was taken before dissolution is obtained.
V. Ostwald recalled with what surprise P. Kleve, Arrhenius’s teacher, a famous chemist, asked him, pointing to a glass with an aqueous solution of potassium chloride: “But it is nonsense to admit, together with Arrhenius, that in dissolved potassium chloride chlorine and potassium are separated from each other?
The English scientist T. Fitzpatrick, for example, in 1888 could not assume the existence of “free” atoms in a solution, since if there were free chlorine atoms in the solution, the solution would have to have some of the properties of a chlorine solution. Arrhenius's theory answered this question simply. In the process of electrolytic dissociation, for example, table salt, not sodium and chlorine atoms are formed, but ions, which, due to the electric charge, have special properties that are sharply different from the properties of electrically neutral atoms.
It remained unclear, however, what reasons determine the appearance of free charged ions, what are the conditions for their existence in solutions? Where does the energy come from for the decomposition of strong compounds during dissolution? S. Arrhenius could not give an answer to these questions. The fact is that he considered the solvent – water – as an inert medium that does not interact with ions. But this is completely false. And the first who guessed about the chemical interaction between the soluble substance and the solvent was Dmitry Ivanovich Mendeleev.
2.1 THEORY OF ELECTROLYTIC DISSOCIATION AND D.I. MENDELEEV
D.I. Mendeleev closely monitored the development of this new theory, but refrained from any categorical assessment of it. He examines in detail some of the arguments that supporters of the theory of electrolytic dissociation appeal to when proving the very fact of the decomposition of salts into ions, including a decrease in the freezing point and other factors determined by the properties of solutions. His “Note on the dissociation of dissolved substances” is devoted to these and other questions related to the understanding of this theory.
He talks about the possibility of solvents combining with dissolved substances and their influence on the properties of solutions. Without making a categorical statement, D.I. Mendeleev, at the same time, points out the need not to discount the possibility of a multilateral consideration of processes: “Before recognizing the dissociation into ions M + X in a salt solution MX, it follows, in the spirit of all information about solutions, look for aqueous solutions of MX salts for the effect of H2O producing MOH + HX particles, or for the dissociation of MX(n+1)H2O hydrates into MOHmH2O + HX(n-m)H2O hydrates, or even direct MXnH2O hydrates into individual molecules.” It follows from this that D.I. Mendeleev did not deny the theory itself, but rather pointed out the need for its development and understanding, taking into account the consistently developed theory of interaction between the solvent and the solute.
After numerous experiments with solutions of sulfuric acid and certain salts, he developed a chemical theory of solutions. Its main idea is approximately this: the dissolution of a substance in water is accompanied by a chemical interaction between the dissolved substance and water. D.I. Mendeleev called the compounds formed in this case hydrates, and the theory itself - hydrate. He managed to quite reliably detect some hydrates in his experiments.
Mendeleev's hydration theory helped to draw many important conclusions, including explaining where the energy needed to separate ions comes from. The very strong chemical interaction between ions and solvent molecules provides the energy needed to destroy the crystal lattice or electrolyte molecules. In the case of aqueous solutions, this energy is called the energy of hydration (hydor in Greek water) and it can reach very large values; Thus, the hydration energy of Na + cations is almost twice as high as the bond breaking energy in the Cl 2 molecule. To separate cations and anions in electrolyte crystals, a lot of energy is also required (it is called the energy of the crystal lattice). As a result, if the total energy of hydration of cations and anions during the formation of a solution is greater than the energy of the crystal lattice (or the binding energy between atoms in electrolytes such as HCl, H2SO4), dissolution will be accompanied by heating, and if less, by cooling of the solution. That is why, when substances such as LiCl, anhydrous CaCl 2 and many others are dissolved in water, the solution heats up, and when KCl, KNO 3, NH 4 NO 3 and some others are dissolved, it cools. The cooling can be so strong that the glass in which the solution is prepared becomes covered with dew on the outside and may even freeze to the wet stand.
The mechanism of electrolytic dissociation can be considered using the example of hydrogen chloride. The H–Cl bond is covalent, polar, HCl molecules are dipoles with a negative pole on the Cl atom and a positive pole on the H atom. Water molecules are also polar. In an aqueous solution, HCl molecules are surrounded on all sides by water molecules so that the positive poles of the H2O molecules are attracted to the negative poles of the HCl molecules, and the negative poles are attracted to the positive poles of the HCl molecules. As a result, the H–Cl bond becomes strongly polarized and breaks with the formation of hydrated H + cations and Cl – anions: H 2 O dipoles seem to pull the HCl molecules apart into separate ions. Each H + cation in solution is surrounded on all sides by H 2 O dipoles, with their negative poles directed towards it, and each Cl – anion is surrounded by oppositely oriented H 2 O dipoles. Similar processes occur in water with H 2 SO 4 molecules, other molecules with polar covalent bonds, as well as with ionic crystals. They already contain “ready” ions, and the role of water dipoles is reduced to separating cations from anions.
However, S. Arrhenius, carried away by the successes of the theory of electrolytic dissociation, did not understand and did not want to listen to objections to some significant shortcomings of his theory. It is possible that the desire to protect his theory from attacks and the need for serious changes in some of its basic premises prevented Arrhenius from correctly assessing the rational aspects of Mendeleev’s teaching. One way or another, Arrhenius himself thought for a long time that anomalies and contradictions would be overcome within the framework of the theory of electrolytic dissociation itself as accurate experimental material accumulated.
In the responses of Arrhenius, Ostwald, van't Hoff to criticism, there were no attempts at reconciliation or compromise with opponents. On the contrary, they quite confidently and sharply criticized certain provisions of their opponents and proved the correctness of their views. Over time, the number of supporters of electrolytic dissociation increased so much that its opponents began to talk about “a wild horde of ionists” taking part in the struggle.
2.2 THEORY OF ELECTROLYTIC DISSOCIATION IN RUSSIA
Heated debates between supporters of the chemical theory of solutions and the theory of electrolytic dissociation especially flared up in Russia in the 90s of the last century. In those years, every congress of chemists or meeting of a scientific chemical society was an arena of heated debates and brilliant speeches in defense of one theory or another. This was the case at the IX (1894) and XI (1901) congresses of Russian naturalists and doctors. Before the XI Congress, on the initiative of D.P. Konovalov, a staunch opponent of the theory of electrolytic dissociation, the Russian Chemical Society chose a topic for a report at a joint meeting of the section of physics and chemistry “Analysis of objections to the theory of electrolytic dissociation” and instructed V.A. Kistyakovsky to make this report as a supporter this theory.
Kistyakovsky, Vladimir Aleksandrovich - Russian Soviet physical chemist, academician of the USSR Academy of Sciences (1929; corresponding member 1925). Born in Kyiv; son of the famous lawyer Alexander Fedorovich Kistyakovsky (1833-1885). In 1889 he graduated from St. Petersburg University. In 1889-1890 worked at the University of Leipzig in the laboratory of V.F. Ostwald. In 1896-1903. Private Associate Professor at St. Petersburg University. In 1902-1903 assistant, 1903-1934 Professor of the St. Petersburg (Leningrad) Polytechnic Institute. Head of the Colloid-Electrochemical Laboratory of the USSR Academy of Sciences (1930-1935), Director of the Colloid-Electrochemical Institute of the USSR Academy of Sciences (1935-1939), Head of the Metal Corrosion Department of the Institute of Physical Chemistry (1939-1952).
Scientific works are devoted to the study of solutions, chemical thermodynamics, and electrochemistry. He was one of the first to put forward (1888) the idea of combining the hydration theory of solutions by D.I. Mendeleev and Ostwald’s theory of electrolytic dissociation. Simultaneously and independently of I.A. Kablukov, he introduced (1889-1891) the concept of ion solvation. He discovered (1904) a rule relating the dependence of the height of the capillary rise of a liquid at the boiling point on the molecular weight (Kistyakovsky’s rule). He derived a formula connecting the vapor pressure in capillaries with the value of surface tension and the molecular weight of the liquid. He established the relationships between the molar values of the heat of evaporation and the volume of vapor at the boiling point (1916), the compressibility coefficient and internal pressure of the liquid (1918), the heat of evaporation and the boiling point of an unassociated liquid (1922), the heat of fusion and the number of atoms in the molecule (1922). He compiled a theoretically based table of electrode potentials and carried out research in the field of electrochemistry of magnesium, chromium, iron, aluminum and other metals (1910). He developed ideas about the processes of metal corrosion and electrocrystallization of metals with the formation of a thin protective film on their surface, impenetrable to atmospheric oxygen. Investigated (1929-1939) the phenomena of corrosion during polyphase contact. The results of Kistyakovsky's research have found application in the practice of protecting metals from corrosion, in the techniques of electroplating and metal refining.
V.A. Kistyakovsky agreed to make a report, although he clearly understood that it would not be easy for him to speak in front of authoritative Russian chemists, who for the most part were opponents of the theory of electrolytic dissociation.
In his report, V.A. Kistyakovsky noted that the theory of electrolytic dissociation represents a step forward towards the introduction of quantitative methods in chemistry. In his opinion, the hypothesis of free ions not only does not contradict, as his opponents say, the existing basic principles of physics, but, on the contrary, is a direct consequence of the principle of conservation of matter and energy and Faraday’s law. He showed that the theory of electrolytic dissociation is increasingly confirmed by new facts. The data of this theory in the light of the doctrine of the dielectric constant and the association of molecules are associated with chemistry. Therefore, chemistry can serve to qualitatively explain the properties of solutions. According to Kistyakovsky, theoretical chemistry should remain “on the basis of the theory of electrolytic dissociation, as on a path leading chemistry onto the broad path of theoretical knowledge.”
Towards the end of the first decade of the 20th century, disputes between supporters of the chemical theory of solutions and the theory of electrolytic dissociation began to subside. Many issues were clarified during the joint discussion; certain provisions of Arrhenius's theory were changed and supplemented. The convergence of the two theories occurred with the further development of the study of the chemical properties of solutions, the processes of solvation, association and complex formation. It turned out that both theories are right, that they cannot even “live” without each other, because they describe the same phenomenon from different sides - the dissolution of substances.
These theories were combined by the Russian chemist I.A. Kablukov. In 1891, his book “Modern Theories of Solutions in Connection with the Doctrine of Chemical Equilibrium” appeared. In it, he showed that the chemical theory of Mendeleev and the theory of electrolytic dissociation of Arrhenius do not contradict each other, but are mutually complementary, if we assume that electrolytes are dissociated into hydrated ions.
Kablukov, Ivan Alekseevich - Russian Soviet physical chemist. Born in the village. Prussians (now Moscow region) in the family of a dentist (freed serf). In 1880 he graduated from the natural sciences department of the Faculty of Physics and Mathematics of Moscow University, where he studied chemistry under V.V. Markovnikov. In 1881-1882 worked in the chemical laboratory of A.M. Butlerov at St. Petersburg University, after which he continued to work at Moscow University with V.V. Markovnikov. In 1882-1884 taught at the Higher Women's Courses in Moscow, from 1885. – private associate professor at Moscow University. In 1889 he worked at the University of Leipzig in the laboratory of V.F. Ostwald under the direction of S. Arrhenius. Since 1899 - Professor at the Moscow Agricultural Institute, since 1903 - Professor at Moscow University. Honorary Member of the USSR Academy of Sciences (1932; Corresponding Member 1928), Honored Scientist of the RSFSR (1929), Honored Professor of Moscow University (since 1910).
The works relate mainly to the electrochemistry of non-aqueous solutions. Studied (1889-1891) the electrical conductivity of electrolytes in organic solvents; established the anomalous conductivity of non-aqueous solutions and its increase when water is added to alcohol solutions. Based on these observations, he suggested the presence of a chemical interaction between solvents and the solute.
Author of textbooks on inorganic and physical chemistry (“Fundamentals of inorganic chemistry”, “Thermochemistry”, “Physical and colloidal chemistry”), a number of works on the history of chemistry. Known as an outstanding teacher and popularizer of science. He actively participated in the work of scientific societies - Russian Physico-Chemical and others.
Kablukov proved that the dissolution of electrolytes in water is accompanied by their dissociation, but the resulting ions immediately undergo hydration. The first process requires significant energy costs, but the second process is accompanied by the release of a significant amount of energy, which mainly covers and sometimes even exceeds the costs of dissociation. This was how the foundation of the modern concept of solutions was laid.
- SOLUTION THEORY
In 1730, R. Reaumur observed that the thermal expansion of alcohol increases the more it is purified. He found that the boiling points of water and alcohol are constant, that when two different liquids are dissolved, an increase or decrease in volume occurs (1733): when mixing alcohol and water, a compression of volume was observed, and it was greatest in the case of a mixture of two parts water and one part alcohol . Even earlier, in 1713, E. Geoffroy noticed that if you add alcohol to water, the temperature of the solution will increase.
In 1732, G. Boerhaave found that water, dissolving a certain amount of salt, forms a saturated solution, which no longer has the ability to dissolve salt. Further information about solubility was obtained by studying the increase in solubility with increasing temperature. In connection with these observations and measurements, scientists had to solve the question of what the dissolution process is and what changes the substance undergoes.
In the 18th century, studies of dissolution processes led scientists to the conclusion that a solution is formed as a result of the chemical interaction of a solute and a solvent. This point of view replaced the corpuscular theory of dissolution, which dominated the works of chemists of the late 17th and early 18th centuries. In 1722, F. Hoffmann proved that during dissolution, the solvent combines with the solute. G. Burhaave shared the same point of view. Based on the study of the physical and chemical properties of solutions, K. Berthollet at the beginning of the 19th century came to the general conclusion that any type of dissolution is a joining process. According to his views, solutions are indeterminate compounds of a solute and a solvent.
Thus, the idea of the manifestation of chemistry in solutions found many supporters. The theory of solutions received the greatest development in the works of D.I. Mendeleev.
3.1 CHEMICAL THEORY OF SOLUTIONS D.I. MENDELEEV
In 1865, D.I. Mendeleev’s doctoral dissertation “On the combination of alcohol with water” was published. What tasks did the scientist set for himself? He primarily sought to improve the method for determining the density of solutions of two liquids - alcohol and water, which was of great practical importance. The study of measurements of the density of solutions allowed Mendeleev to find out the dependence of changes in the properties of a solution on its composition. He found that at a certain ratio of components, a noticeable compression of solutions occurs. He explained the reason for this compression by the formation of the compound C 2 H 5 OH? 3H 2 O. On this basis, the scientist came to the general conclusion that solutions are obtained as a result of the interaction of the components forming the solution. As a result of this interaction, certain chemical compounds - hydrates - are formed in the solution.
Mendeleev, Dmitry Ivanovich - Russian chemist Dmitry Ivanovich Mendeleev was born in Tobolsk in the family of a gymnasium director. While studying at the gymnasium, Mendeleev had very mediocre grades, especially in Latin. In 1850, he entered the Department of Natural Sciences of the Faculty of Physics and Mathematics of the Main Pedagogical Institute in St. Petersburg. Among the professors of the institute at that time were such outstanding scientists as the physicist E. H. Lenz, the chemist A. A. Voskresensky, and the mathematician N. V. Ostrogradsky. In 1855, Mendeleev graduated from the institute with a gold medal and was appointed senior teacher at a gymnasium in Simferopol, but due to the outbreak of the Crimean War, he transferred to Odessa, where he worked as a teacher at the Richelieu Lyceum.
In 1856, Mendeleev defended his master's thesis at St. Petersburg University, in 1857 he was approved as a private lecturer at this university and taught a course in organic chemistry there. In 1859-1861 Mendeleev was on a scientific trip to Germany, where he worked in the laboratory of R. Bunsen and G. Kirchhoff at the University of Heidelberg. One of Mendeleev’s important discoveries dates back to this period - the determination of the “absolute boiling point of liquids,” now known as the critical temperature. In 1860, Mendeleev, together with other Russian chemists, took part in the International Congress of Chemists in Karlsruhe, at which S. Cannizzaro presented his interpretation of the molecular theory of A. Avogadro. This speech and discussion regarding the distinction between the concepts of atom, molecule and equivalent served as an important prerequisite for the discovery of the periodic law.
In 1864, Mendeleev was elected professor of chemistry at the St. Petersburg Institute of Technology. In 1865, he defended his doctoral dissertation “On the combination of alcohol with water” (the topic of the dissertation is often used to substantiate the legend about his invention of 40-proof vodka). In the same year, Mendeleev was confirmed as a professor of technical chemistry at St. Petersburg University, and two years later he headed the department of inorganic chemistry.
Mendeleev developed the doctrine of periodicity until the end of his life. Among Mendeleev’s other scientific works, one can note a series of works on the study of solutions and the development of the hydration theory of solutions (1865–1887). In 1872 he began studying the elasticity of gases, the result of which was proposed in 1874. generalized equation of state of an ideal gas (Cliperon–Mendeleev equation). In 1880–1885 Mendeleev dealt with the problems of oil refining and proposed the principle of its fractional distillation.
Mendeleev was one of the founders of the Russian Chemical Society (1868) and was repeatedly elected its president. In 1876, Mendeleev became a corresponding member of the St. Petersburg Academy of Sciences, but Mendeleev’s candidacy for academicianship was rejected in 1880.
D.I. Mendeleev was a member of more than 90 academies of sciences, scientific societies, and universities in different countries. Chemical element No. 101 (mendelevium), an underwater mountain range and a crater on the far side of the Moon, and a number of educational institutions and scientific institutes are named after Mendeleev. In 1962, the USSR Academy of Sciences established a prize and a Gold Medal named after. Mendeleev for the best work in chemistry and chemical technology, in 1964 Mendeleev's name was included on the honor board of the University of Bridgeport in the USA along with the names of Euclid, Archimedes, N. Copernicus, G. Galileo, I. Newton, A. Lavoisier.
In 1865-1867, D.I. Mendeleev noted that the dissolution process is divided into two stages: the dissolved substance, together with one part of the liquid, forms a certain chemical compound, and this latter dissolves in the rest of the same liquid. Each stage is determined by the action of chemical forces, but of varying intensity. After much searching, he came to the conclusion that this interaction was in the nature of an exchange of particles of the compound with particles of excess solvent. Mendeleev believed that solutions can be reconciled with the atomistic theory if we introduce the concepts of “association” and “dissociation”, if we consider solutions as the most general case of chemical interaction, when weak forces of chemical affinity appear.
The basic principles of the doctrine of solutions, the interaction of substances that make up the solution, the formation of certain compounds that are in a state of dissociation and mobile equilibrium, subject to the law of mass action, were developed by Mendeleev in 1883-1887. He collected and systematized a large amount of factual material, which he presented in the fundamental monograph “Study of aqueous solutions by specific gravity” (1887).
While studying two-component systems, D.I. Mendeleev examined the dependence of density on composition for 233 substances. He studied the dependence of density on composition for aqueous solutions of alkalis, nitric acid, oxygen, nitrogen, carbonic acid, solutions of various organic compounds at different temperatures and concentrations.
Mendeleev studied solutions in which the properties of the substances forming the solution changed greatly. Without recognizing the interaction of the components of the solution, according to Mendeleev, it is impossible to explain changes in properties when the composition of solutions changes.
“Now it is clear to me and already beyond doubt,” wrote D.I. Mendeleev, “that solutions are governed by the usual laws of chemical action, that they contain the same specific compounds that make chemistry so powerful.”
3.2 OSMOTIC THEORY OF VANT HOFF
etc.................
Arrhenius drew attention to the close connection between the ability of solutions of salts, acids and bases to conduct electric current and the deviations of solutions of these substances from van't Hoff and Raoult's laws. He showed that from the electrical conductivity of a solution it is possible to calculate its osmotic pressure, and therefore the correction factor L. The values of i that he calculated from the electrical conductivity coincided well with the values found for the same solutions by other methods.
The reason for the excessively high osmotic pressure of electrolyte solutions is, according to Arrhenius, the dissociation of electrolytes into ions. As a result, on the one hand, the total number of particles in the solution increases, and consequently, the osmotic pressure, decrease in vapor pressure and changes in boiling and freezing temperatures increase, on the other hand, the ions determine the ability of the solution to conduct electric current.
These assumptions were later developed into a coherent theory called theory of electrolytic dissociation.
According to this theory, when dissolved in water, electrolytes break up (dissociate) into positively and negatively charged ions. Positively charged ions are called cations, these include, for example, hydrogen and metal ions. Negatively charged ions are called anions: these include acid residue ions and hydroxide ions. Like solvent molecules, ions in solution are in a state of disordered thermal motion.
The process of electrolytic dissociation is depicted using chemical equations. For example, the dissociation of HCl is expressed by the equation:
The breakdown of electrolytes into ions explains the deviations from Van't Hoff's and Raoult's laws, which were discussed at the beginning of this chapter. As an example, we cited the decrease in the freezing point of the NaCL solution. Now it is not difficult to understand why the decrease in the freezing point of this solution is so great. Sodium chloride goes into solution in the form of Na + and Cl - ions. In this case, from one mole of NaCl, not 6.02 IO 23 particles are obtained, but twice as many. Therefore, the decrease in freezing temperature in a NaCl solution should be twice as large as in a non-electrolyte solution of the same concentration.
Similarly, in a very dilute solution of barium chloride, dissociating according to the equation
the osmotic pressure turns out to be 3 times greater than that calculated by Van't Hoff's law, since the number of particles in the solution is 3 times greater than if barium chloride were in it in the form of BaCl 2 molecules.
Thus, the features of aqueous solutions of electrolytes, which at first glance contradict the laws of Van't Hoff and Raoult, were explained on the basis of these same laws.
However, Arrhenius' theory did not take into account the complexity of phenomena in solutions. In particular, she considered ions as free particles independent of solvent molecules. Arrhenius's theory was opposed by Mendeleev's chemical, or hydrate, theory of solutions, which was based on the idea of the interaction of a solute with a solvent. In overcoming the apparent contradiction of both theories, great merit belongs to the Russian scientist I.A. Kablukov, who first suggested the hydration of ions. The development of this idea subsequently led to the unification of the theories of Arrhenius and Mendeleev.
- Ivan Alekseevich Kablukov (1857-1942) studied the electrical conductivity of solutions. His work “Modern theories of solutions (Van't Hoffai Arrhenius) in connection with the doctrine of chemical equilibrium” had a great influence on the development of physical chemistry in Russia and contributed to the deepening of the theory of electrolytic dissociation.
Have you ever wondered why some solutions conduct electricity and others do not? For example, everyone knows that it is better not to take a bath while blow-drying your hair. After all, water is a good conductor of electric current, and if a working hair dryer falls into the water, it cannot be avoided. In fact, water is not such a good conductor of current. There are solutions that conduct electricity much better. Such substances are called electrolytes. These include acids, alkalis and water-soluble salts.
Electrolytes - who are they?
The question arises: why do solutions of some substances transmit electricity, while others do not? It's all about charged particles - cations and anions. When dissolved in water, electrolytes break up into ions, which, when exposed to electric current, move in a given direction. Positively charged cations move towards the negative pole, the cathode, and negatively charged anions move towards the positive pole, the anode. The process of decomposition of a substance into ions when melted or dissolved in water is proudly called electrolytic dissociation.
This term was coined by the Swedish scientist S. Arrhenius when he studied the properties of solutions to transmit electricity. To do this, he short-circuited some substance through a solution and monitored whether the light bulb came on or not. If an incandescent light bulb lights up, it means the solution conducts electricity, which leads to the conclusion that this substance is an electrolyte. If the light bulb remains extinguished, then the solution does not conduct electricity, therefore this substance is a non-electrolyte. Non-electrolytes include solutions of sugar, alcohol, and glucose. But solutions of table salt and sulfuric acid conduct electricity well, therefore electrolytic dissociation occurs in them.
How does dissociation occur?
Subsequently, the theory of electrolytic dissociation was developed and supplemented by Russian scientists I.A. Kablukov and V.A. Kistyakovsky, applying to its justification the chemical theory of solutions by D.I. Mendeleev.
These scientists found that the electrolytic dissociation of acids, alkalis and salts occurs as a result of the hydration of the electrolyte, that is, its interaction with water molecules. The ions, cations and anions formed as a result of this process will be hydrated, that is, associated with water molecules that surround them in a dense ring. Their properties differ significantly from unhydrated ions.
So, in a solution of strontium nitrate Sr(NO3)2, as well as in solutions of cesium hydroxide CsOH, electrolytic dissociation occurs. Examples of this process can be expressed as follows:
Sr(NO3)2 = Sr2+ + 2NO3 -,
those. upon dissociation of one molecule of strontium nitrate, one strontium cation and 2 nitrate anions are formed;
CsOH = Cs+ + OH-,
those. the dissociation of one cesium hydroxide molecule produces one cesium cation and one hydroxide anion.
Electrolytic dissociation of acids occurs similarly. For hydroiodic acid, this process can be expressed by the following equation:
those. the dissociation of one molecule of hydroiodic acid produces one hydrogen cation and one iodine anion.
Dissociation mechanism.
Electrolytic dissociation of electrolyte substances occurs in several stages. For substances with an ionic type of bond, such as NaCl, NaOH, this process includes three sequential processes:
First, water molecules, which have 2 opposite poles (positive and negative) and represent a dipole, are oriented at the ions of the crystal. They are attached with their positive pole to the negative ion of the crystal, and vice versa, with the negative pole - to the positive ion of the crystal;
then the crystal ions are hydrated by water dipoles,
and only after this the hydrated ions seem to diverge in different directions and begin to move randomly in the solution or melt until they are affected by an electric field.
For substances such as HCl and other acids, the dissociation process is similar, except that at the initial stage, a transition from a covalent bond to an ionic bond occurs due to the action of water dipoles. These are the main points of the theory of dissociation of substances.
The theory of electrolytes dates back to the first half of the 19th century, when M. Faraday conducted his famous experiments with solutions of table salt. He established that absolutely pure water It conducts electricity very poorly, but if you add a few salt crystals to it, the conductivity immediately increases. Even then, the assumption was born that salt disintegrates in water into certain particles that are capable of conducting electric current, however, full theory, which describes all these processes in solutions, appeared much later.
Electrolytic dissociation theory
The theory, the founder of which was Svante Arrhenius in the period 1883-1887, is based on the idea that when molecules of a soluble substance (electrolyte) enter a polar or non-polar liquid, they dissociation into ions. Electrolytes are compounds that spontaneously disintegrate in solution into ions capable of independent existence. The number of ions formed, their structure and the magnitude of the charge depend only on the nature of the dissociated molecule.
To use the theory in describing the properties of dissolution, a number of assumptions are used, namely: it is assumed that the dissociation is incomplete, the ions (their electron shells) do not react with each other, and their behavior can be described by the law of mass action under ideal conditions. If we consider theoretical system, where the electrolyte CA is in phase equilibrium with the products of its dissociation - the K+ cation and the A- anion, then, according to the law of mass action, it is possible to construct an equation for the dissociation reaction:
KA = K+ + A- (1)
The equilibrium constant, written in terms of the concentrations of substances under isothermal conditions, will have the following value:
Kd = x / (2)
In this case (in equation 2), the equilibrium constant Kd will be nothing more than the dissociation constant, the values , , on the right side are the equilibrium concentrations of the electrolyte and its dissociation products.
Taking into account the assumptions of the Arrhenius theory, which were applied by the author, in particular, about the incompleteness of dissociation, the concept of the degree of dissociation - α is introduced. Thus, if we express the concentration of the solution C (mol/l), then per liter of solution there is αC mole of electrolyte (CA), and its equilibrium concentration can be expressed as (1-α)C mol/l. From the reaction equation (1) it is obvious that per αC mole of electrolyte (CA) the same amount of K+ and A- ions is formed. If we substitute all these quantities into equation (2) and carry out a number of simplifications, we obtain the formula dissociation constants(degree of dissociation formula):
Kd = ∝ 2 x C /1-∝ (3)
This equation allows us to quantify the degree of electrolytic dissociation in different solutions.
Arrhenius's theory gave rise to many scientific directions in chemistry: with its help, the first theories of acids and bases were created, explanations were given for physicochemical processes in homogeneous systems. However, it is not without its shortcomings, which mainly relate to the fact that the theory does not take into account interionic interactions.
Classification of electrolytes with examples
Electrolytes are classified into weak and strong, periodically distinguishing a group of medium-strength electrolytes. Strong electrolytes are characterized by the fact that they disintegrate in solution fully. As a rule, these are strong mineral acids, for example:
- Nitric acid - HNO3.
- Hydrochloric acid - HCl.
- Perchloric acid - HClO4.
- Phosphoric acid - H3PO4.
Strong electrolytes can be bases, for example:
- Potassium hydroxide - KOH.
The bulk of strong electrolytes are the vast majority of salts (NaCl, Na2SO4, Ca (NO3)2, CH3COONa, chlorides, sulfides).
Weak electrolytes, on the contrary, are partially hydrated in solutions. This group should include inorganic acids (H2CO3, H3BO3, H3AsO4), weak bases (ammonium), some salts (HgCl2), organic acids (CH3COOH, C6H5COOH), phenols and amines. IN non-aqueous solutions the same compounds can be both strong and weak electrolytes, thus depending on the nature of the solvent.
Dissociation of acids, bases and salts
Patterns for acids
During the electrical dissociation of acids in aqueous solutions, positively charged hydrogen ions (H+) are necessarily formed as cations:
HNO3 → H+ + NO3-
If the acid is polybasic (for example: the dissociation equation of H2SO4), then dissociation occurs sequentially, each time eliminating one hydrogen ion:
H2SO4 → H + + HSO4- first stage - hydrogen sulfate ion
HSO4- → H + + SO4- second stage - sulfate ion
The process for a polybasic acid, as a rule, proceeds as much as possible in the first stage, the degree of dissociation of subsequent ones is much less.
Characteristics of the process for alkalis
When alkalis dissociate in aqueous solutions, a negatively charged hydroxyl ion (OH-) is necessarily formed:
NaOH → Na+ + OH-
The process for polyacid bases (for example, the mechanism of dissociation of magnesium hydroxide) proceeds in a multi-step manner similar to polybasic acids:
Mg (OH)2 → OH- + Mg (OH)+ first stage
Mg (OH)+ → OH- + Mg2+ second stage
There are also cases when both hydrogen cations and hydroxyl anions can be formed during the dissociation process (during the dissociation of ampholytes or amphoteric compounds, for example, Zn, Al):
2OH- + Zn2+ + 2H2O ←→ Zn (OH)2 + H2O ←→ 2- + 2H+
Flow rules for acidic and basic salts
For acidic salts, the main pattern is as follows - cations (positively charged metals) dissociate first, and only then hydrogen cations:
KHSO4 → K+ + HSO4- first stage
HSO4 - → H+ + SO4- second stage
In basic salts, first of all, acid residues go into solution, and only then the hydroxyl ion:
BaOHCl → Cl- + Ba (OH)+ first stage
Ba (OH)+ → OH- + Ba2+ second stage
pH value
Definition, essence and meaning
Dissociation processes can occur not only for dissolved substances, but also a solvent. Thus, water is itself a weak electrolyte and is characterized by dissociation to a very small extent. The process equation can be written as follows:
H2O= H3O+ + OH-
One water molecule dissociates into positively charged hydrogen ions and negatively charged hydronium anions. It is the concentration of these ions that determines the acidity level of the solution - the more hydronium ions, the more acidic the solution.
The concentration of hydronium ions in real solutions is, as a rule, very small (for example: 5 × 10−6 g/l) and therefore, for convenience, this value is taken logarithmically, and in order to obtain a positive value, it is taken with the opposite sign. Let us briefly formulate a strict definition of the concept of “hydrogen index” or pH.
pH (hydrogen index) is the negative natural logarithm of the concentration of hydronium ions, reflecting the acidity of the solution.
pH= - log
The pH values are usually assessed on a scale from 0 to 14, where 0 is the most acidic solution, and 14 is the most alkaline. A neutral solution (corresponding to the pH of pure water) is considered to be a solution with a value of 7. For example, here are several typical solutions with characteristic pH values:
Much less often they resort to using another indicator - pOH. In its meaning, it is absolutely similar to the hydrogen index, except that the concentration of hydroxyl ions is taken as the basis.