Lecture
By the early 19th century, natural scientists had established that all solutions could be divided into two large groups — those that do not conduct electric current and those that do. The German physicist G. S. Ohm found that conducting solutions carry current according to the same laws as metallic conductors. The English physicist M. Faraday studied the phenomenon of electrolysis and established its basic laws. The phenomena and patterns observed in solutions were explained in 1887 by the Swedish scientist Svante Arrhenius, who proposed the theory of electrolytic dissociation. Its essence can be expressed as follows:
electrolytes break down into ions in melts under the action of high temperature, and in solutions as a result of interaction with the solvent (Fig. 56).
The process of an electrolyte breaking down into ions is called electrolytic dissociation.
Svante August
Arrhenius —
Swedish physical
chemist, author of the theory
of electrolytic
dissociation (1887)
In the solid state, most ionic and molecular compounds are dielectrics that do not conduct current. Thus, the electrical conductivity of solid NaCl is 1000 billion times lower than that of metals or graphite. At the same time, aqueous solutions of electrolytes conduct current because ionic and many molecular compounds, under the influence of polar water molecules, break down (dissociate) into cations and anions upon dissolution.
An electrolyte — is a substance that conducts electric current as a result of dissociation into cations and anions in a melt or in solution.
The electrical conductivity of electrolytes is caused by the mobility of cations and anions in melts and solutions.
Electrolytes include salts, alkalis, and acids, that is, compounds with an ionic or covalent polar chemical bond.
Ionic crystals consist of cations and anions bound to one another by electrostatic forces of interaction in an ordered crystal structure. Under the action of polar water molecules, ionic crystals break down and dissolve, forming hydrates of ions (Fig. 57).
Study Figure 57 and pay attention to the relative size of Na+ cations, Cl– anions, and water molecules, as well as the signs of the partial charges on the H and O atoms in the water molecule. Note which atom of the water molecule is turned toward the Na+ cation and which toward the Cl– anion. Why is the water molecule oriented toward the Na+ cation with its oxygen atom, and toward the Cl– anion with its hydrogen atom?
A special feature of the dissociation of ionic compounds (salts and alkalis) is that these substances already consist of ions, and polar water molecules only separate the cations and anions. The energy expended on dissociation is compensated by the energy released during hydration of the ions, and overall the process proceeds spontaneously.
Being bases that are highly soluble in water, alkalis dissociate in aqueous solutions into hydroxide anions and metal cations:
As a result of the dissociation of salts in water, their ionic crystals break down into metal (or ammonium) cations and anions of acid residues:
Most organic and inorganic acids are soluble in water and are electrolytes. In acids, hydrogen atoms are linked by covalent polar bonds to the atoms of acid residues. There are no ready-made ions in polar molecules. Let us consider the process of acid dissociation using the example of a 
molecule.
First, the polar acid molecules are attracted to the polar water molecules and become even more polarized.
Hydrogen bonds form between the acid and water molecules (Fig. 58).
Then the covalent polar bond 
lengthens and becomes less strong. As a result, it breaks, that is, ions are formed:
The H+ cation formed upon dissociation of the acid attaches to a water molecule to form a hydronium cation H3O+ by a donor-acceptor mechanism, with which you became acquainted while studying the material in § 13. The chlorine ion formed during ionization is also hydrated, as shown in Figure 57.
To simplify equations of reactions in ionic form, the hydrated hydrogen cation is usually written as H+.
In a similar way, under the action of polar water molecules, other compounds with a covalent polar type of bond also dissociate.
Dissociation includes the processes of hydration of the initial molecules, their polarization, the breaking of bonds within the molecules, and their breakdown to form hydrated ions.
Acid molecules, upon dissociation in aqueous solutions, dissociate fully or partially into hydrogen cations and anions of acid residues:
Atoms and molecules are electrically neutral particles. Metal cations, for example K+, as you know, are formed as a result of metal atoms giving up electrons. Anions of nonmetals, for example Br–, are formed as a result of the addition of electrons to nonmetal atoms.
The properties of cations and anions differ from the properties of the corresponding electrically neutral atoms. Thus, in the electric field created by electrodes, an ordered movement of the electrolyte's cations and anions occurs — an electric current. Positive ions — metal cations, H+, 
, and others move toward the negative electrode (the cathode).
Negative ions — anions of acid residues and OH– move toward the positive electrode (the anode) (Fig. 58.1).
Unlike neutral atoms (in particular, metal atoms), cations, for example Cu2+, Fe3+, and others, have fewer electrons in their outer electron levels and a smaller size, and can always act as an oxidizing agent and accept electrons:
CuSO4 + Zn = ZnSO4 + Cu↓
or Cu2+ + Zn = Zn2+ + Cu↓.
Compared with neutral atoms (I, S, O), simple anions, for example, I–, S2– , O2– , and others, have more electrons in their outer electron levels and a larger size, and can act as a reducing agent and give up electrons:
2NaI + Cl2 = 2NaCl + I2
or 2I– + Cl2 = 2Cl– + I2.
Strong and weak electrolytes
Simultaneously with electrolytic dissociation in an electrolyte solution, processes of ion association also occur. Therefore, in electrolytes only a certain fraction of the substance exists in the form of free ions, that is, the substance may not decompose completely into ions.
To quantitatively characterize electrolytic dissociation, the concept of the degree of dissociation α is used.
The degree of dissociation α — is the ratio of the number of molecules or formula units of the electrolyte that have decomposed into ions to the total number of its dissolved molecules or formula units:
.
The number of particles and their amount (mol) are related to one another by the relationship:
, hence, 
,
where n is the amount (mol) of the corresponding molecules or formula units.
The value of α is expressed in fractions of unity or as a percentage. The degree of dissociation shows what part of the electrolyte has decomposed and exists in solution in the form of free ions. For example, the degree of dissociation of acetic acid in a solution with a molar concentration of 0.1 mol/dm3 is equal to 1.3 %. This means that out of 1000 acid molecules, only 13 decomposed into H+ cations and CH3COO– anions.
Ion association — is the combination of oppositely charged ions under the action of electrostatic forces into neutral or charged particles.
Examples of association may be irreversible exchange reactions (they will be examined in § 27), written in ionic form:
Ag+ + Cl– = AgCl↓,
or reversible dissociation reactions of weak acids:
.
It might seem that strong acids in aqueous solutions dissociate completely into ions. Nevertheless, moistened litmus paper turns red when placed above 20–30% solutions of HCl and HNO3. This indicates the presence, in the solution and in the vapors above the solution, of undissociated HCl and HNO3 molecules, which dissociate upon contact with the moist paper. However, above 5–10% solutions of these strong electrolytes, litmus paper no longer turns red. This indicates the complete dissociation of HCl and HNO3 molecules into non-volatile ions even in such relatively concentrated solutions.
One can assess how strong dissociation is in aqueous solutions of ammonia. Moist litmus paper placed above a dilute ammonia solution turns blue. This is explained by the fact that most of the ammonia in water is in molecular form and evaporates. Thus, above a 4% solution of NH3, the number of NH3 molecules is one and a half times greater than the number of water vapor molecules.
The probability of ion association decreases as the electrolyte concentration decreases, and consequently the degree of dissociation in dilute solutions is greater than in concentrated ones. This means that the degree of dissociation depends not only on the nature of the substance but also on the concentration of the electrolyte.
Depending on the numerical value of α, electrolytes are conventionally divided into strong and weak.
Strong electrolytes — are chemical compounds that in dilute solutions are practically completely dissociated into ions. The degree of dissociation α of strong electrolytes is close to 1.
Strong electrolytes include soluble ionic compounds and some substances with a covalent polar type of bond; in their case the process of association of the formed ions is practically absent. Therefore, in equations reflecting electrolytic dissociation, instead of an equals sign, a one-directional arrow is placed. As examples, let us give the dissociation equations of soluble salts (a), alkalis (b), and strong acids (c) — H2SO4, HCl, HBr, HI, HClO4, HNO3:
The degree of dissociation of weak electrolytes is, as a rule, less than 5 %. Weak electrolytes include water (
), as well as:
Weak electrolytes — chemical compounds that dissociate into ions only to an insignificant extent even in dilute solutions. The ions of weak electrolytes are in mobile equilibrium with molecules that have not dissociated, or with a substance of non-molecular structure that has not dissolved.
The dissociation of weak electrolytes is depicted by a decomposition equation with arrows in two directions: 
.
It is quite difficult to draw a clear boundary between strong and weak acids. Hydrofluoric acid (HF) is conventionally considered a weak acid, although in dilute solutions it dissociates to the extent of 8–25 %. Phosphoric acid H3PO4, sulfurous acid H2SO3, and oxalic acid HOOC—COOH are also considered weak. At the first dissociation step, they dissociate to the extent of 10–30 % in concentrated solutions and more than 50 % in dilute ones.
For example, in a dilute solution H3PO4 dissociates according to the equation:
The degree of dissociation at this step can reach 50 %. At each subsequent step, the degree of dissociation decreases significantly, and dissociation proceeds according to the equations:
The decrease in the degree of dissociation is related to the strengthening of the interaction of hydrogen ions 
with the acid residues in the series 
, 
and 
.
The degree of dissociation α is usually determined experimentally from the electrical conductivity of the solution.
The dissociation of weak electrolytes, as a reversible reaction, obeys Le Chatelier's principle. Thus, the H+ cations formed at the first dissociation step, according to Le Chatelier's principle, shift the equilibrium of the dissociation reaction at the second and third steps to the left. This thereby decreases the degree of dissociation of the anions 
and 
and the concentration of the anions and 
respectively.
As a rule, at each subsequent step the degree of dissociation decreases.
In a solution of H3PO4 with a molar concentration of 1 mol/dm3, the fractions of H3PO4 molecules and of the anions 
, and formed at various stages amount to:
| H3PO4 | H+ | |
|
|
| 92 % | ~8 % | ~8 % | ~6 · 10–6 % | ~10–16 % |
All polybasic weak acids dissociate practically only at the first step. Let us consider one more example: the dissociation of weak hydrosulfuric acid. At the first step, a small part of the H2S molecules dissociate to form hydrogen ions and hydrosulfide ions, and at the second step sulfide ions are formed from the hydrosulfide ions:
H2S 
H+ + HS–;
HS– H+ + S2–.
Dissociation at the second step proceeds significantly more weakly, because:
Electrolytes are substances whose aqueous solutions or melts conduct electric current.
The process of dissociation of molecular compounds with covalent polar bonds differs from the process of dissociation of ionic compounds by the presence of a stage of molecule polarization and cleavage of the covalent bond.
According to the degree of dissociation, strong and weak electrolytes are distinguished.
The degree of dissociation depends not only on the nature of the substances but also on the concentration of the solutions, and increases with their dilution.
1. Why do salts conduct electric current after melting?
2. How does electric current in metals differ from electric current in solutions or melts of electrolytes?
3. Write the dissociation equations for the following substances: KHCO3, HClO4, Sr(OH)2, LiOH, КHSO4, NH4H2PO4, HNO2.
4. Why does electrolytic dissociation in solutions proceed spontaneously?
5. Compare the oxidizing and reducing abilities:
6. Why, for the anions of most polybasic acids, for example , 
, 
, 
, does dissociation at the second step proceed more weakly than at the first? How will the degree of dissociation change when hydrogen ions are added to solutions of these acids?
7. An aqueous solution contains 50 mol of hydrogen fluoride. What is the total number of ions formed upon its dissociation, if α(HF) = 9 %?
8. Determine the total amount of cations and anions of the salt in a solution containing sodium sulfate with a mass of 2.84 g.
9. What is the molar concentration of H+ cations in an aqueous solution of acetic acid with a volume of 4 dm3, if the degree of dissociation of the acid is 2.6 % and the mass of the acid is 6 g?
10. In a solution with a volume of 2.5 dm3, the mass of Ba(OH)2 is 14 g. Calculate the molar concentration of OH– ions in this solution.
*Self-check
1. The dissociation of both NaCl and HCl includes:
2. The F– ion differs from the fluorine atom F in:
3. A solution of a certain salt contains 2 mol of cations and 3 mol of anions. This salt could be:
4. Dissociation equations of weak electrolytes:
2H++ 
; H+ + ОH–; С6Н5О– + Н+;5. In a solution of nitrous acid, the number of undissociated molecules is 4 times greater than the number of dissociated ones. The degree of dissociation of the acid is equal to:
By the beginning of the 19th century, natural scientists had established that all solutions can be divided into two large groups — those that do not conduct electric current and those that do. The German physicist G. S. Ohm found that conducting solutions carry current according to the same laws as conducting metals. The English physicist M. Faraday studied the phenomenon of electrolysis and established its basic laws. The phenomena and regularities observed in solutions were explained in 1887 by the Swedish scientist Svante Arrhenius, who proposed the theory of electrolytic dissociation. Its essence can be expressed as follows:
electrolytes break down into ions in melts under the action of high temperature, and in solutions as a result of interaction with the solvent (fig. 56).
The process of an electrolyte breaking down into ions is called electrolytic dissociation.
In the solid state, most ionic and molecular compounds are dielectrics that do not conduct current. Thus, the electrical conductivity of solid NaCl is 1000 billion times lower than that of metals or graphite. At the same time, aqueous solutions of electrolytes conduct current because ionic and many molecular compounds, under the influence of polar water molecules, break down (dissociate) into cations and anions upon dissolution.
An electrolyte — is a substance that conducts electric current as a result of dissociation into cations and anions in a melt or in a solution.
The electrical conductivity of electrolytes is due to the mobility of cations and anions in melts and solutions.
Electrolytes include salts, alkalis, and acids, that is, compounds with an ionic or covalent polar chemical bond.
Dissociation of compounds with an ionic type of bond
Ionic crystals consist of cations and anions bound to each other by electrostatic interaction forces in an ordered crystal structure. Under the action of polar water molecules, ionic crystals are destroyed and dissolve with the formation of ion hydrates (fig. 57).
Study figure 57 and note the relative size of the Na+ cations, Cl– anions, and water molecules, as well as the signs of the partial charges on the H and O atoms in the water molecule. Note which atom of the water molecule is turned toward the Na+ cation, and which toward the Cl– anion. Why is the water molecule oriented toward the Na+ cation with its oxygen atom, and toward the Cl– anion with its hydrogen atom?
A feature of the dissociation of ionic compounds (salts and alkalis) is that these substances already consist of ions, and polar water molecules only separate the cations and anions. The energy expended on dissociation is compensated by the energy released during hydration of the ions, and overall the process proceeds spontaneously.
Being readily soluble bases in water, alkalis dissociate in aqueous solutions into hydroxide anions and metal cations:
As a result of the dissociation of salts in water, their ionic crystals break down into metal (or ammonium) cations and acid-residue anions:
Dissociation of compounds with a covalent polar type of bond
Most organic and inorganic acids are soluble in water and are electrolytes. In acids, hydrogen atoms are bound to acid-residue atoms by covalent polar bonds. There are no ready-made ions in polar molecules. Let us consider the process of acid dissociation using the example of the molecule.
First, the polar acid molecules are attracted to the polar water molecules and become even more polarized.
Hydrogen bonds form between the acid and water molecules (fig. 58).
Then the covalent polar bond lengthens and becomes less strong. As a result, it breaks, that is, ions are formed:
The H+ cation produced by the dissociation of the acid attaches to a water molecule to form a hydronium cation H3O+ by the donor-acceptor mechanism, which you became familiar with when studying the material in § 13. The chlorine ion formed during ionization is also hydrated, as shown in figure 57.
To simplify equations of reactions in ionic form, the hydrated hydrogen cation is usually written as H+.
In a similar way, under the action of polar water molecules, other compounds with a covalent polar type of bond also dissociate.
Dissociation includes the processes of hydration of the initial molecules, their polarization, the breaking of bonds in the molecules, and their breakdown with the formation of hydrated ions.
Acid molecules, upon dissociation in aqueous solutions, fully or partially dissociate into hydrogen cations and acid-residue anions:
*The difference in chemical properties between ions and atoms/molecules
Atoms and molecules are electroneutral particles. Metal cations, for example K+, as you know, are formed as a result of metal atoms giving up electrons. Nonmetal anions, for example Br–, are formed as a result of nonmetal atoms gaining electrons.
The properties of cations and anions differ from the properties of the corresponding electroneutral atoms. Thus, in the electric field created by electrodes, an ordered movement of the electrolyte's cations and anions occurs — an electric current. Positive ions — metal cations, H+, and others — move toward the negative electrode (the cathode).
Negative ions — acid-residue anions and OH– — move toward the positive electrode (the anode) (fig. 58.1).
Unlike neutral atoms (in particular, metal atoms), cations, for example Cu2+, Fe3+, and others, have fewer electrons in their outer electron levels and a smaller size, and can always act as an oxidizing agent and accept electrons:
CuSO4 + Zn = ZnSO4 + Cu↓
or Cu2+ + Zn = Zn2+ + Cu↓.
Compared with neutral atoms (I, S, O), simple anions, for example I–, S2– , O2– and others, have more electrons in their outer electron levels and a larger size, and can act as a reducing agent and give up electrons:
2NaI + Cl2 = 2NaCl + I2
or 2I– + Cl2 = 2Cl– + I2.
Strong and weak electrolytes
Simultaneously with electrolytic dissociation, ion association processes take place in an electrolyte solution. Therefore, in electrolytes only a certain fraction of the substance exists in the form of free ions, that is, the substance may not break down completely into ions.
To quantitatively characterize electrolytic dissociation, the concept of degree of dissociation α is used.
Degree of dissociation α — is the ratio of the number of molecules or formula units of the electrolyte that have broken down into ions to the total number of its dissolved molecules or formula units:
.
The number of particles and their amount (moles) are related by the relation:
, hence, ,
where n is the amount (moles) of the corresponding molecules or formula units.
The value of α is expressed as a fraction of unity or as a percentage. The degree of dissociation shows what part of the electrolyte has broken down and exists in solution in the form of free ions. For example, the degree of dissociation of acetic acid in a solution with a molar concentration of 0.1 mol/dm3 is equal to 1.3 %. This means that out of 1000 acid molecules, only 13 have broken down into H+ cations and CH3COO– anions.
Ion association — the combination of oppositely charged ions under the action of electrostatic forces into neutral or charged particles.
Examples of association can be irreversible exchange reactions (they will be discussed in § 27), written in ionic form:
Ag+ + Cl– = AgCl↓,
or reversible dissociation reactions of weak acids:
.
One would think that strong acids fully dissociate into ions in aqueous solutions. Nevertheless, moistened litmus paper turns red when placed above 20–30% solutions of HCl and HNO3. This indicates the presence, in the solution and in the vapors above it, of undissociated HCl and HNO3 molecules, which dissociate upon reaching the moist paper. However, above 5–10% solutions of these strong electrolytes, the litmus paper no longer turns red. This indicates the complete dissociation of HCl and HNO3 molecules into nonvolatile ions even in such relatively concentrated solutions.
One can assess how strong the dissociation is in aqueous ammonia solutions. Moist litmus paper placed above a dilute ammonia solution turns blue. This is explained by the fact that most of the ammonia in water is in the molecular state and volatilizes. Thus, above a 4% solution of NH3, the number of NH3 molecules is one and a half times greater than the number of water vapor molecules.
The probability of ion association decreases as the electrolyte concentration decreases, and consequently the degree of dissociation in dilute solutions is greater than in concentrated ones. This means that the degree of dissociation depends not only on the nature of the substance but also on the concentration of the electrolyte.
Depending on the numerical value of α, electrolytes are conventionally divided into strong and weak.
Strong electrolytes — are chemical compounds that are practically completely dissociated into ions in dilute solutions. The degree of dissociation α of strong electrolytes is close to 1.
Strong electrolytes include soluble ionic compounds and some substances with a covalent polar type of bond; in their case, the process of association of the resulting ions is practically absent. Therefore, in equations reflecting electrolytic dissociation, an arrow pointing in one direction is used instead of an equals sign. As examples, let us give the dissociation equations of soluble salts (a), alkalis (b), and strong acids (c) — H2SO4, HCl, HBr, HI, HClO4, HNO3:
The degree of dissociation of weak electrolytes is, as a rule, less than 5 %. Weak electrolytes include water (), as well as:
Weak electrolytes — are chemical compounds that dissociate into ions only to a small extent even in dilute solutions. The ions of weak electrolytes are in mobile equilibrium with the molecules that have not dissociated, or with a substance of non-molecular structure that has not dissolved.
The dissociation of weak electrolytes is represented by a breakdown equation with arrows in two directions: .
It is quite difficult to draw a clear boundary between strong and weak acids. Hydrofluoric acid (HF) is conventionally considered weak, although in dilute solutions it dissociates to 8–25 %. Phosphoric H3PO4, sulfurous H2SO3, and oxalic HOOC—COOH acids are also considered weak. At the first dissociation step, they dissociate to 10–30 % in concentrated solutions and more than 50 % in dilute solutions.
For example, in a dilute solution H3PO4 dissociates according to the equation:
The degree of dissociation at this step can reach 50 %. At each subsequent step, the degree of dissociation decreases significantly, and dissociation proceeds according to the equations:
The decrease in the degree of dissociation is related to the strengthening of the interaction of hydrogen ions with the acid residues in the series , and .
The degree of dissociation α is usually determined experimentally from the electrical conductivity of the solution.
The dissociation of weak electrolytes, being a reversible reaction, obeys Le Chatelier's principle. Thus, the H+ cations formed at the first dissociation step, according to Le Chatelier's principle, shift the equilibrium of the dissociation reaction at the second and third steps to the left. This reduces the degree of dissociation of the and anions and the concentration of the and anions, respectively.
As a rule, the degree of dissociation decreases at each subsequent step.
In a solution of H3PO4 with a molar concentration of 1 mol/dm3, the fractions of H3PO4 molecules and of the anions , and formed at different stages are:
| H3PO4 | H+ | |
|
|
| 92 % | ~8 % | ~8 % | ~6 · 10–6 % | ~10–16 % |
All polybasic weak acids dissociate practically only at the first step. Let us consider one more example — the dissociation of weak hydrosulfuric acid. At the first step, a small fraction of H2S molecules dissociate to form hydrogen ions and hydrosulfide ions; at the second step, sulfide ions are formed from hydrosulfide ions:
H2S H+ + HS–;
HS– H+ + S2–.
Dissociation at the second step proceeds significantly more weakly, because:
Electrolytes are substances whose aqueous solutions or melts conduct electric current.
The process of dissociation of molecular compounds with covalent polar bonds differs from the process of dissociation of ionic compounds by the presence of a stage of molecule polarization and covalent bond breaking.
According to their degree of dissociation, strong and weak electrolytes are distinguished.
The degree of dissociation depends not only on the nature of the substances, but also on the concentration of the solutions, and increases with their dilution.
Questions, tasks, problems
1. Why do salts conduct electric current after melting?
2. How does electric current in metals differ from electric current in solutions or melts of electrolytes?
3. Write the dissociation equations for the following substances: KHCO3, HClO4, Sr(OH)2, LiOH, КHSO4, NH4H2PO4, HNO2.
4. Why does electrolytic dissociation in solutions proceed spontaneously?
5. Compare the oxidizing and reducing ability of:
6. Why, for the anions of most polybasic acids, for example , , , , does dissociation at the second step proceed more weakly than at the first? How will the degree of dissociation change when hydrogen ions are added to solutions of these acids?
7. An aqueous solution contains hydrogen fluoride in an amount of 50 mol. What is the total number of ions formed upon its dissociation, if α(HF) = 9 %?
8. Determine the total number of cations and anions of the salt in a solution containing sodium sulfate with a mass of 2.84 g.
9. What is the molar concentration of H+ cations in an aqueous solution of acetic acid with a volume of 4 dm3, if the degree of dissociation of the acid is 2.6 %, and the mass of the acid is 6 g?
10. In a solution with a volume of 2.5 dm3, the mass of Ba(OH)2 is 14 g. Calculate the molar concentration of OH– ions in this solution.
*Self-check
1. The dissociation of both NaCl and HCl includes:
2. The F– ion differs from the fluorine atom F in:
3. A solution of a certain salt contains 2 mol of cations and 3 mol of anions. This salt could be:
4. Dissociation equations of weak electrolytes:
2H++ ; H+ + ОH–; С6Н5О– + Н+;5. In a solution of nitrous acid, the number of undissociated molecules is 4 times greater than the number of dissociated ones. The degree of dissociation of the acid is equal to:
By the beginning of the 19th century, natural scientists had established that all solutions can be divided into two large groups — those that do not conduct electric current and those that do. The German physicist G. S. Ohm found that conducting solutions carry current according to the same laws as conducting metals. The English physicist M. Faraday studied the phenomenon of electrolysis and established its basic laws. The phenomena and regularities observed in solutions were explained in 1887 by the Swedish scientist Svante Arrhenius, who proposed the theory of electrolytic dissociation. Its essence can be expressed as follows:
electrolytes break down into ions in melts under the action of high temperature, and in solutions as a result of interaction with the solvent (fig. 56).
The process of an electrolyte breaking down into ions is called electrolytic dissociation.
In the solid state, most ionic and molecular compounds are dielectrics that do not conduct current. Thus, the electrical conductivity of solid NaCl is 1000 billion times lower than that of metals or graphite. At the same time, aqueous solutions of electrolytes conduct current because ionic and many molecular compounds, under the influence of polar water molecules, break down (dissociate) into cations and anions upon dissolution.
An electrolyte — is a substance that conducts electric current as a result of dissociation into cations and anions in a melt or in a solution.
The electrical conductivity of electrolytes is due to the mobility of cations and anions in melts and solutions.
Electrolytes include salts, alkalis, and acids, that is, compounds with an ionic or covalent polar chemical
продолжение следует...
Часть 1 5.25. Electrolytic dissociation of substances in solutions
Часть 2 - 5.25. Electrolytic dissociation of substances in solutions
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