What dissociates in an aqueous solution. Electrolytic dissociation

The fundamental support of chemistry, along with the periodic system of D. I. Mendeleev, the structure of organic compounds of A. M. Butlerov, and other significant discoveries, is the theory of electrolytic dissociation. It was developed by Svante Arrhenius in 1887 to explain the specific behavior of electrolytes in water and other polar liquids and melts. He found a compromise between two categorically different theories about solutions existing at that time - physical and chemical. The first argued that the solute and the solvent do not interact with each other in any way, forming a simple mechanical mixture. The second is that there is a chemical bond between them. It turned out that in fact, solutions have both properties.

In subsequent stages of the development of science, many scientists continued research and development in this area, relying on available information about the structure of atoms and the nature of chemical bonds between them. In particular, I. A. Kablukov studied the issue of solvation processes, V. A. Kistyakovsky determined the dependence of the rise of a liquid column in a capillary under boiling temperature conditions on molecular weight.

Modern interpretation of the theory

Before this discovery, many properties and circumstances of the splitting processes were not studied, like the solutions themselves. Electrolytic dissociation is the process of decomposition of a substance into its constituent ions in water or other polar liquids, interaction of particles of the compound with solvent molecules, and the appearance of mobility of cations and anions at the nodes of the crystal lattice due to melting. As a result of this, the formed substances acquire a new property - electrical conductivity.

Ions, being in a free state of solution or melt, interact with each other. Likely charged ones repel, oppositely charged ones attract. Charged particles are solvated by solvent molecules - each is tightly surrounded by strictly oriented dipoles in accordance with the Coulomb attractive forces; in particular, they are hydrated if the medium is aqueous. Cations always have larger radii than anions due to the specific arrangement of particles around them with charges localized at the edges.

Composition, classification and names of charged particles in the light of electrolytic dissociation

An ion is an atom or group of atoms that carries a positive or negative charge. They are characterized by a conventional division into simple (K (+), Ca (2+), H (+) - consisting of one chemical element), complex and complex (OH (-), SO 4 (2-), HCO 3 (- ) - from several). If a cation or anion is associated with a solvent molecule, it is called solvated, and with a dipole of the H 2 O molecule - hydrated.

When electrolytic dissociation of water occurs, two charged particles H (+) and OH (-) are formed. A hydrogen proton accepts a lone electron pair of oxygen from another water molecule into a vacant orbital, resulting in the formation of a hydronium ion H 3 O (+).

The main provisions of Arrhenius's discovery

All representatives of the classes of inorganic compounds, except for oxides, disintegrate in solutions of oriented dipoles of liquids; in chemical terms, they dissociate into their constituent ions to a greater or lesser extent. This process does not require the presence of electric current; the equation of electrolytic dissociation is its schematic notation.

Once in a solution or melt, ions can be exposed to an electric current and move directionally towards the cathode (negative electrode) and anode (positive). The latter attract oppositely charged atomic aggregates. This is where the particles got their names - cations and anions.

In parallel and simultaneously with the decomposition of the substance, the reverse process occurs - the association of ions into the original molecules, therefore one hundred percent dissolution of the substance does not occur. This equation for the reaction of electrolytic dissociation contains an equal sign between its right and left sides. Electrolytic dissociation, like any other reaction, is subject to laws governing chemical equilibrium, and the law of mass action is no exception. It states that the rate of decomposition into ions is proportional to the concentration of the electrolyte.

Classification of substances during dissociation

Chemical terminology divides substances into insoluble, slightly soluble and soluble. The last two are weak and strong electrolytes. Information on the solubility of certain compounds is summarized in the solubility table. The dissociation of strong electrolytes is an irreversible process; they completely disintegrate into ions. Weak - only partially, they are characterized by the phenomenon of association, and therefore, the equilibrium of the processes taking place.

It is important to note that there is no direct relationship between solubility and electrolyte strength. In strong people it can be weakly expressed. Just like weak electrolytes, they can be highly soluble in water.

Examples of compounds whose solutions conduct electric current

The class of “strong electrolytes” includes all well-dissociating acids, such as nitric, hydrochloric, bromic, sulfuric, perchloric and others. To the same extent, alkalis are alkaline hydroxides and individual representatives of the “alkaline earth metals” group. The electrolytic dissociation of salts is intense, except for certain cyanates and thiocyanates, as well as mercury (II) chloride.

The class of “weak electrolytes” is represented by other mineral and almost all organic acids: carbonic, sulfide, boric, nitrous, sulfurous, silicon, acetic and others. As well as poorly soluble and hydrocarbon bases and amphoteric hydroxides (hydroxides of magnesium, beryllium, iron, zinc in the oxidation state (2+)). In turn, water molecules are very weak electrolytes, but still break down into ions.

Quantitative description of dissociating processes

The degree of electrolytic dissociation actually characterizes the scale of the splitting process. It can be calculated - the number of particles split into ions must be divided by the total number of molecules of the dissolved substance in the system. This value is denoted by the letter “alpha”.

It is logical that for strong electrolytes “α” is equal to one, or one hundred percent, since the number of decayed particles is equal to their total number. For the weak - always less than one. Complete disintegration of the original molecules into ions in an aqueous environment does not occur, and the reverse process occurs.

The main factors influencing the completeness of decay

The degree of electrolytic dissociation is influenced by a number of undeniable factors. First of all, the nature of the solvent and the substance that disintegrates in it is important. For example, all strong electrolytes have a covalent, highly polar or ionic type of bond between their constituent particles. Liquids are represented by dipoles, in particular water, in the molecules there is a separation of charges, and as a result of their specific orientation, electrolytic dissociation of the dissolved substance occurs.

The alpha value is inversely affected by concentration. As it increases, the value of the degree of dissociation decreases, and vice versa. The process itself is entirely endothermic, that is, a certain amount of heat is required to initiate it. The influence of the temperature factor is justified as follows: the higher it is, the greater the degree of dissociation.

Minor Factors

Polybasic acids, such as phosphoric acid, and bases containing several hydroxyl groups, for example, Fe(OH) 3, decompose into ions in stages. A dependence has been determined - each subsequent stage of dissociation is characterized by a degree that is thousands or tens of thousands of times less than the previous one.

The degree of decomposition can also be changed by the addition of other electrolytes to the system, changing the concentration of one of the ions of the main solute. This entails a shift of equilibrium to the side, which is determined by the Le Chatelier-Brown rule - the reaction proceeds in the direction in which the neutralization of the influence exerted on the system from the outside is observed.

Classical equilibrium process constant

To characterize the process of decomposition of a weak electrolyte, in addition to its degree, the electrolytic dissociation constant (K d) is used, which is expressed by the ratio of the concentrations of cations and anions to the quantitative content of the original molecules in the system. In essence, it is the usual chemical equilibrium constant for the reversible reaction of the splitting of a dissolved substance into ions.

For example, for the process of decomposition of a compound into its constituent particles, the dissociation constant (K d) will be determined by the quotient of the constant concentrations of cations and anions in the solution, raised to powers corresponding to the numbers preceding them in the chemical equation, and the total number of remaining non-dissociated formula units dissolved substance. There is a dependence - the higher (K ​​d), the greater the number of cations and anions in the system.

The relationship between the concentration of a weak decaying compound, the degree of dissociation and the constant is determined using the Ostwald dilution law by the equation: K d = α 2 s.

Water as a weakly dissociating substance

Dipole molecules disintegrate to an extremely small extent into charged particles, since this is energetically unfavorable. Still, splitting into hydrogen cations and hydroxyl anions occurs. Taking into account hydration processes, we can talk about the formation of hydronium and OH (-) ions from two water molecules.

Permanent dissociation is determined by the ratio of the product of hydrogen protons and hydroxide groups, called the ionic product of water, to the equilibrium concentration of undissociated molecules in solution.

The electrolytic dissociation of water determines the presence of H (+) in the system, which characterize its acidity, and the presence of OH (-) - basicity. If the concentrations of the proton and hydroxyl group are equal, such a medium is called neutral. There is a so-called hydrogen index - this is a negative logarithm of the total quantitative content of H (+) in the solution. A pH less than 7 indicates that the environment is acidic, more indicates that it is alkaline. This is a very important value; biological, biochemical and chemical reactions of various water systems - lakes, ponds, rivers and seas - are analyzed based on its experimental value. The relevance of the hydrogen index for industrial processes is also undeniable.

Recording reactions and notation

The equation of electrolytic dissociation using chemical symbols describes the processes of decomposition of molecules into corresponding particles and is called ionic. It is many times simpler than the standard molecular one and has a more general appearance.

When drawing up such an equation, it must be taken into account that substances that precipitate or are removed from the reacting mixture as part of gas vapor during the reaction must always be written only in molecular form, in contrast to electrolyte compounds, strong representatives of which are included in the composition only in the form of split into ions solutions. Electrolytic dissociation for them is an irreversible process, since association is impossible due to the formation of non-fissionable substances or gases. For this type of equation, the same rules apply as for other chemical reactions - the sums of the coefficients of the left and right sides must be equal to each other to maintain material balance.

Electrolytic dissociation of acids and bases can occur in several stages if the substances are polybasic or polyacidic. For each subreaction, its own equation is written.

Role in chemical science and its development

The creation of the theory of Svante Arrhenius was of the greatest importance for the general process of formation of physical and, in particular, electrochemical science. Based on the discovery of such a phenomenon as electrolytic dissociation, electrode processes, the specificity of the passage of currents through various media, and the theory of inducing cathode-anode potentials received intensive development. In addition, the theory of solutions has advanced significantly. Unprecedented discoveries awaited chemical kinetics, the field of corrosion of metals and alloys, as well as work to find new means of protection against it.

There is still so much new and unknown in the modern world. Every day, scientists are moving further in their knowledge of such a great discipline as chemistry. Electrolytic dissociation, as well as its creators and followers, have forever occupied an honorable place in the context of the development of world science.

1. ELECTROLYTES

1.1. Electrolytic dissociation. Degree of dissociation. Electrolyte Power

According to the theory of electrolytic dissociation, salts, acids, and hydroxides, when dissolved in water, completely or partially disintegrate into independent particles - ions.

The process of decomposition of substance molecules into ions under the influence of polar solvent molecules is called electrolytic dissociation. Substances that dissociate into ions in solutions are called electrolytes. As a result, the solution acquires the ability to conduct electric current, because mobile electric charge carriers appear in it. 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 ions of acidic residues and hydroxide ions.

To quantitatively characterize the dissociation process, the concept of the degree of dissociation was introduced. The degree of dissociation of an electrolyte (α) is the ratio of the number of its molecules disintegrated into ions in a given solution ( n ), to the total number of its molecules in solution ( N), or

α = .

The degree of electrolytic dissociation is usually expressed either in fractions of a unit or as a percentage.

Electrolytes with a degree of dissociation greater than 0.3 (30%) are usually called strong, with a degree of dissociation from 0.03 (3%) to 0.3 (30%) - medium, less than 0.03 (3%) - weak electrolytes. So, for a 0.1 M solution CH3COOH α = 0.013 (or 1.3%). Therefore, acetic acid is a weak electrolyte. The degree of dissociation shows what part of the dissolved molecules of a substance has broken up into ions. The degree of electrolytic dissociation of an electrolyte in aqueous solutions depends on the nature of the electrolyte, its concentration and temperature.

By their nature, electrolytes can be divided into two large groups: strong and weak. Strong electrolytes dissociate almost completely (α = 1).

Strong electrolytes include:

1) acids (H 2 SO 4, HCl, HNO 3, HBr, HI, HClO 4, H M nO 4);

2) bases – metal hydroxides of the first group of the main subgroup (alkali) – LiOH, NaOH, KOH, RbOH, CsOH , as well as hydroxides of alkaline earth metals – Ba (OH) 2, Ca (OH) 2, Sr (OH) 2;.

3) salts soluble in water (see solubility table).

Weak electrolytes dissociate into ions to a very small extent; in solutions they are found mainly in an undissociated state (in molecular form). For weak electrolytes, an equilibrium is established between undissociated molecules and ions.

Weak electrolytes include:

1) inorganic acids ( H 2 CO 3, H 2 S, HNO 2, H 2 SO 3, HCN, H 3 PO 4, H 2 SiO 3, HCNS, HClO, etc.);

2) water (H 2 O);

3) ammonium hydroxide ( NH 4 OH);

4) most organic acids

(for example, acetic CH 3 COOH, formic HCOOH);

5) insoluble and slightly soluble salts and hydroxides of some metals (see solubility table).

Process electrolytic dissociation depicted using chemical equations. For example, dissociation of hydrochloric acid (HC l ) is written as follows:

HCl → H + + Cl – .

Bases dissociate to form metal cations and hydroxide ions. For example, the dissociation of KOH

KOH → K + + OH – .

Polybasic acids, as well as bases of polyvalent metals, dissociate stepwise. For example,

H 2 CO 3 H + + HCO 3 – ,

HCO 3 – H + + CO 3 2– .

The first equilibrium - dissociation according to the first step - is characterized by the constant

.

For second stage dissociation:

.

In the case of carbonic acid, the dissociation constants have the following values: K I = 4.3× 10 –7, K II = 5.6 × 10–11. For stepwise dissociation always K I > K II > K III >... , because the energy that must be expended to separate an ion is minimal when it is separated from a neutral molecule.

Average (normal) salts, soluble in water, dissociate to form positively charged metal ions and negatively charged ions of the acid residue

Ca(NO 3) 2 → Ca 2+ + 2NO 3 –

Al 2 (SO 4) 3 → 2Al 3+ +3SO 4 2–.

Acid salts (hydrosalts) are electrolytes containing hydrogen in the anion, which can be split off in the form of the hydrogen ion H +. Acid salts are considered as a product obtained from polybasic acids in which not all hydrogen atoms are replaced by a metal. Dissociation of acid salts occurs in stages, for example:

KHCO 3 → K + + HCO 3 – (first stage)

There are two main reasons for the passage of electric current through conductors: either due to the transfer of electrons, or due to the transfer of ions. Electronic conductivity is inherent primarily in metals. Ionic conductivity is inherent in many chemical compounds that have an ionic structure, for example, salts in solid or molten states, as well as many aqueous and non-aqueous solutions.

All substances by their behavior in solutions usually divided into two categories:

a) substances whose solutions have ionic conductivity (electrolytes);

b) substances whose solutions do not have ionic conductivity (non-electrolytes).

Electrolytes include most inorganic acids, bases and salts. Non-electrolytes include many organic compounds, such as alcohols and carbohydrates.

It turned out that electrolyte solutions have lower melting points and higher boiling points compared to the corresponding values ​​for a pure solvent or for a solution of a non-electrolyte in the same solvent. To explain these facts, Arrhenius proposed theory of electrolytic dissociation.

Under electrolytic dissociation refers to the breakdown of electrolyte molecules in solution with the formation of positively and negatively charged ions - cations and anions. For example, a molecule of acetic acid dissociates in an aqueous solution like this:

CH 3 COOH CH 3 COO - + H +

The dissociation process in all cases is reversible, therefore, when writing the equations for the dissociation reaction, the reversibility sign is used. Different electrolytes dissociate into ions to varying degrees. The completeness of decomposition depends on the nature of the electrolyte, its concentration, the nature of the solvent, and temperature.

Strong and weak electrolytes. Degree of dissociation. Dissociation constant. Degree of dissociation α called - the ratio of the number of molecules disintegrated into ions (n) to the total number of dissolved molecules (n 0).

α = (n/n 0)?100

The degree of dissociation can vary from 0 to 1, from no dissociation to complete dissociation. Depending on the degree of dissociation, weak and strong electrolytes are distinguished. TO weak electrolytes include substances whose degree of dissociation in 0.1 M solutions is less than 3%; if the degree of dissociation in a 0.1 M solution exceeds 30%, then such an electrolyte is called strong. Electrolytes, the degree of dissociation of which lies in the range from 3% to 30%, are called electrolytes medium strength.

Strong electrolytes include most salts, some acids - HCl, HBr, HI, HNO 3, HClO 4, H 2 SO 4 and bases of alkali and alkaline earth metals - alkalis LiOH, NaOH, KOH, RbOH, CsOH, Ca(OH) 2, Sr(OH) 2, Ba(OH) 2.

The reaction equation for the dissociation of the electrolyte AA into K + cations and A - anions can be generally presented as follows:

KA K + + A -

and degree of dissociation α in this case can be expressed as the ratio of the molar concentration of the formed ions [K + ] or [A - ] to original molar concentration of the electrolyte [AK] o, i.e.

With increasing solution concentration, the degree of electrolyte dissociation decreases.

Polybasic acids and bases dissociate stepwise - first one of the ions is split off from the molecule, then another, etc. Each dissociation step is characterized by its own dissociation constant.

Stage I: H 2 SO 4 → H + + HSO 4 -

Stage II: НSO 4 - Н + + SO 4 2-

General equation: H 2 SO 4 2H + + SO 4 2-

The process of electrolytic dissociation is characterized by dissociation constant (K) . So, for the reaction KA K + + A is the dissociation constant:

K = [K + ] ? [A - ]/[KA]

There is a quantitative relationship between the constant and the degree of electrolytic dissociation. In the example given, we denote the total concentration of the dissolved substance With , and the degree of dissociation α . Then [K + ] = [A - ] = α?с and accordingly the concentration of undissociated particles [CA] = (1 - α )With .

Substituting the values ​​into the expression for the dissociation constant, we obtain the relation

Since the molar concentration is C = 1/V, then

This equation is a mathematical expression of Ostwald's dilution law: the dissociation constant of the electrolyte does not depend on the dilution of the solution.

Ionic product of water. pH solution. Water dissociation constant value K H2O = 1·10 -14 . This constant for water is called ionic product of water, which depends only on temperature.

According to the reaction H 2 OH + + OH -, during the dissociation of water, one OH - ion is formed for each H + ion, therefore, in pure water the concentrations of these ions are the same: [H + ] = [OH - ] = 10 -7 .

pH = -log[H + ]

Aqueous solutions have a pH value in the range from 1 to 14. Based on the ratio of the concentrations of these ions, three types of media are distinguished: neutral, acidic and alkaline.

Neutral environment- an environment in which the ion concentrations [H + ] = [OH - ] = 10 -7 mol/l (pH = 7).

Acidic environment- a medium in which the concentration of [H + ] ions is greater than the concentration of [OH - ] ions, i.e. [H + ] > 10 -7 mol/l (pH< 7).

Alkaline environment- an environment in which the concentration of [H + ] ions is less than the concentration of [OH - ] ions, i.e. [H+]< 10 -7 моль/л (рН > 7).

Qualitatively, the reaction of the medium and the pH of aqueous solutions of electrolytes are determined using indicators and a pH meter.

For example, if the ion concentration = 10 -4 mol/l, then pH = - log10 -4 = 4 and the solution medium is acidic, and if the ion concentration is [OH - ] = 10 -4 mol/l, then [H + ] = TO(H 2 O) - [OH - ] = 10 -14 - 10 -4 = 10 -10, and pH = - log10 -10 = 10 and the solution is alkaline.

Product of solubility. The dissolution of a solid in water stops when a saturated solution is formed, i.e. equilibrium is established between a solid substance and particles of the same substance in solution. So, for example, in a saturated solution of silver chloride the equilibrium is established:

AgCl solid Ag + aq + Cl - aq

In a saturated electrolyte solution, the product of the concentrations of its ions is a constant value at a given temperature and this value quantitatively characterizes the ability of the electrolyte to dissolve, it is called solubility product(ETC).

PR(AgCl) = [Ag + ]

Solubility product - this is a constant value equal to the product of the concentrations of ions of a poorly soluble electrolyte in its saturated solution. In general, for a slightly soluble electrolyte of composition A m B n we can write: A m B n ↔ mA + nB

PR AmBn = [A] m ? [B]n

Knowing the values ​​of solubility products, it is possible to solve issues related to the formation or dissolution of precipitation during chemical reactions, which is especially important for analytical chemistry.

The dissolution of any substance in water is accompanied by the formation of hydrates. If at the same time no formula changes occur in the particles of the dissolved substance in the solution, then such substances are classified as non-electrolytes. They are, for example, gas nitrogen N 2, liquid chloroform CHCl 3, solid sucrose C 12 H 22 O 11, which in aqueous solution exist in the form of hydrates of these molecules.
Many substances are known (in the general form MA), which, after dissolving in water and forming hydrates of MA nH 2 O molecules, undergo significant formula changes. As a result, hydrated ions appear in the solution - cations M + * nH 2 O and anions A * nH 2 O:
MA * nH 2 O → M + * nH 2 O + A - * nH 2 O
Such substances belong to electrolytes.
The process of appearance of hydrated ions in an aqueous solution called electrolytic dissociation(S. Arrhenius 1887).
Electrolytic dissociation of ionic crystalline substances (M +)(A -) in water is irreversible reaction:
(M +)(A -) (t) →(M +)(A -) (p) =(M +) (p) + (A -) (p)
Such substances are considered strong electrolytes, these are many bases and salts, for example:

NaOH = Na + + OH - K 2 SO 4 = 2K + + SO 4 -
Ba(OH) 2 = Ba 2+ + 2OH - Na 2 = 2Na + + S 2-
Electrolytic dissociation of MA substances consisting of polar covalent molecules is reversible reaction:
(M-A) (g,f,t) → (M-A) (p) ↔ M + (p) A - (p)
Such substances are classified as weak electrolytes; they include many acids and some bases, for example:
a) HNO 2 ↔ H + + NO 2-
b) CH 3 COOH ↔ H + + CH 3 COO —
c) H 2 CO 3 ↔ H + + HCO 3 - (first stage)
HCO 3 — ↔ H + + CO 3 2- (second stage)
d) NH 3 * H 2 O ↔ NH 4 + OH -
In dilute aqueous solutions of weak electrolytes, we will always find both the original molecules and the products of their dissociation - hydrated ions.
The qualitative characteristic of the dissociation of electrolytes is called the degree of dissociation and is denoted ɑ 1, always ɑ › 0.
For strong electrolytes ɑ = 1 by definition (the dissociation of such electrolytes is complete).
For weak electrolytes degree of dissociation - the ratio of the minor concentration of the dissociated substance (c d) to the total concentration of the substance in solution (c):

The degree of dissociation is a fraction of one from 100%. For weak electrolytes ɑ ˂ C 1 (100%). For weak acids H n A, the degree of dissociation at each subsequent step decreases sharply compared to the previous one:
H 3 PO 4 ↔ H + + H 2 PO 4 — = 23.5%
H 2 PO 4 — ↔ H + + HPO 4 2- = 3*10 -4%
HPO 4 2- ↔ H + + PO 4 3- = 2*10 -9%
The degree of dissociation depends on the nature and concentration of the electrolyte, as well as on the temperature of the solution; it grows with decrease concentration of the substance in the solution (i.e. when the solution is diluted) at heating.
IN diluted solutions strong acids H n A their hydrothions H n -1 A do not exist, for example:
H 2 SO 4 = H + + (1 → 1)
= H + + SO 4 -2 (1 → 1)
As a result: H 2 SO 4 (dil.) = 2H + + SO 4 -2
V concentrated In solutions, the content of hydroanions (and even parent molecules) becomes noticeable:
H 2 SO 4 — (conc.) ↔ H + + HSO 4 — (1 ˂ 1)
HSO 4 — ↔ H + + SO 4 2- (2 ˂ 1 ˂ 1)
(it is impossible to summarize the equations for the stages of reversible dissociation!). When heated, the values ​​of 1 and 2 increase, which promotes the occurrence of reactions involving concentrated acids.
Acids are electrolytes that, upon dissociation, supply hydrogen cations to an aqueous solution and do not form any other positive anions:
* the letter indicates the degree of occurrence of any reversible reactions, including the degree of hydrolysis.
H 2 SO 4 = 2H + = SO 4 2-, HF ↔ H + + F —
Common strong acids:
Oxygen-containing acids

Anoxic acids
HCl, HBr, HI, HNCS
In a dilute aqueous solution (conditionally up to 10% or 0.1 molar) these acids dissociate completely. For strong acids H n A, the list includes their hydrothiones(anions of acid salts), also dissociating completely under these conditions.
Common weak acids:
Oxygen-containing acids

Anoxic acids
Bases are electrolytes that, when dissociated, supply hydroxide ions to an aqueous solution and do not form any other negative ions:
KOH = K + + OH - , Ca(OH) 2 = Ca 2+ + 2OH -
Dissociation sparingly soluble bases Mg(OH) 2 , Cu(OH) 2 , Mn(OH) 2 , Fe(OH) 2 and others are of no practical importance.
TO strong reasons ( alkalis) include NaOH, KOH, Ba(OH) 2 and some others. The most famous weak base is ammonia hydrate NH 3 H 2 O.
Medium salts are electrolytes that, upon dissociation, supply any cations exceptH + , and any anions exceptOH:
Cu(NO 3) 2 = Cu 2+ + 2NO 3 -
Al 2 (SO 4) 3 = 2Al 3+ + 3SO 4 2-
Na(CH 3 COO) = Na + + CH 3 COO —
BaCl 2 = Ba 2+ + 2Cl
K 2 S = 2K + + S 2-
Mg(CN) 2 = Mg 2+ + 2CN -
We are not just talking about highly soluble salts. Dissociation sparingly soluble and practically insoluble salts don't matter.
Dissociate similarly double salts:
KAl(SO 4) 2 = K + + Al 3+ + 2SO 4 2-
Fe(NH 4) 2 (SO 4) 2 = Fe 2+ + 2NH 4 + 2SO 4 2-
Acid salts(most of them are soluble in water) dissociate completely according to the type of medium salts:
KHSO 4 = K + + HSO 4 —
KHCr 2 O 7 = K + + HCr 2 O 7 —
KH 2 PO 4 = K + + H 2 PO 4 —
NaHCO 3 = Na + + HCO 3 -
The resulting hydroanions are, in turn, exposed to water:
a) if the hydroanion belongs to a strong acid, then it itself also dissociates completely:
HSO 4 - = H + + HSO 4 2-, HCr 2 O 7 - = H + + Cr 2 O 7 2-
and the complete equation for the dissociation reaction will be written as:
KHSO 4 = K + + H + + SO 4 2-
KHCr 2 O 7 = K + + H + Cr 2 O 7 2-
(solutions of these salts will necessarily be acidic, as well as solutions of the corresponding acids);
b) if hydrothione belongs to a weak acid, then its behavior in water is dual - or incomplete dissociation like a weak acid:
H 2 PO 4 — ↔ H + + HPO 4 2- (1)
HCO 3 - ↔ H + CO 3 2- (1)

Or interaction with water (called reversible hydrolysis):
H 2 PO 4 - + H 2 O ↔ H 3 PO 4 + OH - (2)
HCO 3 - + H 2 O ↔ H 2 CO 3 + OH - (2)
At 1 2, dissociation predominates (and the solution will be acidic), and at 1 2, hydrolysis predominates (and the salt solution will be alkaline). So, solutions of salts with the anions HSO 3 -, H 2 PO 4 -, H 2 AsO 4 - and HSeO 3 will be acidic, solutions of salts with other anions (the majority of them) will be alkaline. In other words, the name “acidic” for salts with a majority of hydroanions does not imply that these anions will behave as acids in solution (hydrolysis of hydroanions and calculation of the ratio between 1 and 2 are studied only in high school)

Basic salts MgCl (OH), CuCO 3 (OH) 2 and others are mostly practically insoluble in water, and it is impossible to discuss their behavior in an aqueous solution.

Substances whose solutions (or melts) conduct electric current are called e l e c t r o l i t a m i. Often the solutions of these substances themselves are called electrolytes. These solutions (melts) of electrolytes are conductors of the second kind, since the transmission of electricity is carried out in them by movement and about n about in - charged particles. A positively charged particle is called cation (Ca +2), a particle carrying a negative charge - anion (HE -). Ions can be simple (Ca +2, H +) and complex (PO 4 ־ 3, HCO 3 ־ 2).

The founder of the theory of electrolytic dissociation is the Swedish scientist S. Arrhenius. According to theory electrolytic dissociation is the disintegration of molecules into ions when they are dissolved in water, and this occurs without the influence of an electric current. However, this theory did not answer the questions: what reasons determine the appearance of ions in solutions and why positive ions, when colliding with negative ones, do not form neutral particles.

Russian scientists made their contribution to the development of this theory: D.I. Mendeleev, I.A. Kablukov - supporters of the chemical theory of solutions, who paid attention to the influence of the solvent in the dissociation process. Kablukov argued that the solute interacts with the solvent ( solvation process ) forming products of variable composition ( salts ).

A solvate is an ion surrounded by solvent molecules (solvation shell), of which there can be different numbers (this is how a variable composition is achieved). If the solvent is water, then the process of interaction between the molecules of the solute and the solvent is called g i d r a t a t i e y, and the product of interaction is g i d r a t o m.

Thus, the cause of electrolytic dissociation is solvation (hydration). And it is the solvation (hydration) of ions that prevents them from recombining into neutral molecules.

Quantitatively, the dissociation process is characterized by the value degree of electrolytic dissociation ( α ), which is the ratio of the amount of substance disintegrated into ions to the total amount of dissolved substance. It follows that for strong electrolytes α = 1 or 100% (solute ions are present in the solution), for weak electrolytes 0< α < 1 (в растворе присутствуют наряду с ионами растворенного вещества и его недиссоциированные молекулы), для неэлектролитов α = 0 (there are no ions in the solution). In addition to the nature of the solute and solvent, the quantity α depends on the concentration of the solution and temperature.

If the solvent is water, strong electrolytes include:

1) all salts;

2) the following acids: HCl, HBr, HI, H2SO4, HNO3, HClO4;

3) the following bases: LiOH, NaOH, KOH, RbOH, CsOH, Ca(OH) 2, Sr(OH) 2, Ba(OH) 2.

The process of electrolytic dissociation is reversible, therefore, it can be characterized by the value of the equilibrium constant, which, in the case of a weak electrolyte, is called dissociation constant (K D ) .

The greater this value, the easier the electrolyte breaks down into ions, the more its ions are in the solution. For example: HF ═ H + + F־

This value is constant at a given temperature and depends on the nature of the electrolyte and solvent.

Polybasic acids and polyacid bases dissociate stepwise. For example, sulfuric acid molecules primarily eliminate one hydrogen cation:

H 2 SO 4 ═ H + + HSO 4 ־ .

Elimination of the second ion according to the equation

HSO 4 ־ ═ N + + SO 4 ־ 2

is already much more difficult, since it has to overcome the attraction from the doubly charged SO 4 ־ 2 ion, which, of course, attracts the hydrogen ion more strongly than the singly charged HSO 4 ־ ion. Therefore, the second stage of dissociation occurs to a much lesser extent than the first.

Bases containing more than one hydroxyl group in the molecule also dissociate stepwise. For example:

Ba(OH) 2 ═ BaOH + + OH - ;

BaOH + = Ba 2+ + OH - .

Medium (normal) salts always dissociate into metal ions and acid residues:

CaCl 2 = Ca 2+ + 2Cl - ;

Na 2 SO 4 = 2Na + + SO 4 2- .

Acid salts, like polybasic acids, dissociate stepwise. For example:

NaHCO 3 = Na + + HCO 3 - ;

HCO 3 - = H + + CO 3 2- .

However, the degree of dissociation in the second step is very small, so that the acid salt solution contains only a small number of hydrogen ions.

Basic salts dissociate into basic and acidic ions. For example:

Fe(OH)Cl 2 = FeOH 2+ + 2Cl - .

Almost no secondary dissociation of basic residue ions into metal and hydroxyl ions occurs.