Which are in dynamic equilibrium with undissociated molecules. Weak electrolytes include most organic acids and many organic bases in aqueous and non-aqueous solutions.

Weak electrolytes are:

• almost all organic acids and water;
• some inorganic acids: HF, HClO, HClO 2 , HNO 2 , HCN, H 2 S, HBrO, H 3 PO 4 , H 2 CO 3 , H 2 SiO 3 , H 2 SO 3 and others;
• some sparingly soluble metal hydroxides: Fe(OH) 3 , Zn(OH) 2 and others; as well as ammonium hydroxide NH 4 OH.

Literature

• M. I. Ravich-Sherbo. V. V. Novikov "Physical and colloidal Chemistry" M: Higher school, 1975

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See what "Weak electrolytes" is in other dictionaries:

weak electrolytes- - electrolytes, slightly dissociating in aqueous solutions into ions. The process of dissociation of weak electrolytes is reversible and obeys the law of mass action. General chemistry: textbook / A. V. Zholnin ... Chemical terms

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In a broad sense, liquid or solid in va and systems, in which ions are present in a noticeable concentration, causing the passage of electricity through them. current (ionic conductivity); in the narrow sense into va, which decay into ions in pre. When dissolving E. ... ... Physical Encyclopedia

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In wa, in k ryh in a noticeable concentration there are ions that cause the passage of electric. current (ionic conductivity). E. also called. conductors of the second kind. In the narrow sense of the word, E. in va, molecules to ryh in p re due to electrolytic ... ... Chemical Encyclopedia

- (from Electro ... and Greek lytos decomposable, soluble) liquid or solid substances and systems in which ions are present in any noticeable concentration, causing the passage of electric current. In a narrow sense, E. ... ... Great Soviet Encyclopedia

This term has other meanings, see Dissociation. Electrolytic dissociation is the process of breaking down an electrolyte into ions when it dissolves or melts. Contents 1 Dissociation in solutions 2 ... Wikipedia

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The dissociation of the electrolyte is quantitatively characterized by the degree of dissociation. Degree of dissociation a–is the ratio of the number of molecules dissociated into ions N diss.,to the total number of dissolved electrolyte molecules N :

a =

a is the fraction of electrolyte molecules decomposed into ions.

The degree of electrolyte dissociation depends on many factors: the nature of the electrolyte, the nature of the solvent, the concentration of the solution, and the temperature.

According to the ability to dissociate, electrolytes are conditionally divided into strong and weak. Electrolytes that exist in solution only as ions are called strong . Electrolytes, which in a dissolved state are partly in the form of molecules and partly in the form of ions, are called weak .

Strong electrolytes include almost all salts, some acids: H 2 SO 4, HNO 3, HCl, HI, HClO 4, hydroxides of alkali and alkaline earth metals (see appendix, table 6).

The process of dissociation of strong electrolytes goes to the end:

HNO 3 \u003d H + + NO 3 -, NaOH \u003d Na + + OH -,

and equal signs are put in the dissociation equations.

As applied to strong electrolytes, the concept of "degree of dissociation" is conditional. " Apparent" degree of dissociation (a each) below the true (see appendix, table 6). With an increase in the concentration of a strong electrolyte in a solution, the interaction of oppositely charged ions increases. When approaching each other sufficiently, they form associates. The ions in them are separated by layers of polar water molecules surrounding each ion. This affects the decrease in the electrical conductivity of the solution, i.e. the effect of incomplete dissociation is created.

To take this effect into account, the activity coefficient g is introduced, which decreases with increasing solution concentration, varying from 0 to 1. To quantitatively describe the properties of solutions of strong electrolytes, a quantity called activity (a).

The activity of an ion is understood as that effective concentration of it, according to which it acts in chemical reactions.

Ion activity ( a) is equal to its molar concentration ( FROM) multiplied by the activity factor (g):

a = g FROM.

The use of activity instead of concentration makes it possible to apply to solutions the regularities established for ideal solutions.

Weak electrolytes include some mineral (HNO 2, H 2 SO 3, H 2 S, H 2 SiO 3, HCN, H 3 PO 4) and most organic acids (CH 3 COOH, H 2 C 2 O 4, etc.) , ammonium hydroxide NH 4 OH and all poorly soluble bases in water, organic amines.

The dissociation of weak electrolytes is reversible. In solutions of weak electrolytes, an equilibrium is established between ions and undissociated molecules. In the corresponding dissociation equations, the sign of reversibility ("") is put. For example, the dissociation equation for weak acetic acid is written as follows:

CH 3 COOH « CH 3 COO - + H + .

In a solution of a weak binary electrolyte ( KA) the following equilibrium is established, characterized by an equilibrium constant called the dissociation constant To d:

KA « K + + A - ,

.

If dissolved in 1 liter of solution FROM moles of electrolyte KA and the degree of dissociation is equal to a, which means that dissociated aС moles of electrolyte and each ion was formed according to aС moles. remains in the undissociated state ( FROM – aС) moles KA.

KA « K + + A - .

C - aC aC aC

Then the dissociation constant will be equal to:

(6.1)

Since the dissociation constant does not depend on the concentration, the derived relation expresses the dependence of the degree of dissociation of a weak binary electrolyte on its concentration. Equation (6.1) shows that a decrease in the concentration of a weak electrolyte in a solution leads to an increase in the degree of its dissociation. Equation (6.1) expresses Ostwald's dilution law .

For very weak electrolytes (at a<<1), уравнение Оствальда можно записать следующим образом:

To d a 2 C, or a» (6.2)

The dissociation constant for each electrolyte is constant at a given temperature, it does not depend on the concentration of the solution and characterizes the ability of the electrolyte to decompose into ions. The higher Kd, the more the electrolyte dissociates into ions. The dissociation constants of weak electrolytes are tabulated (see Appendix, Table 3).

Measurement of the degree of dissociation of various electrolytes showed that individual electrolytes at the same normal concentration of solutions dissociate into ions very differently.

The difference in the values ​​of the degree of dissociation of acids is especially great. For example, nitric and hydrochloric acids in 0.1 N. solutions almost completely decompose into ions; carbonic, hydrocyanic and other acids dissociate under the same conditions only to a small extent.

Of the water-soluble bases (alkalis), ammonium oxide hydrate is weakly dissociating, the remaining alkalis dissociate well. All salts, with a few exceptions, also dissociate well into ions.

The difference in the values ​​of the degree of dissociation of individual acids is due to the nature of the valence bond between the atoms that form their molecules. The more polar the bond between hydrogen and the rest of the molecule, the easier it is to split off, the more the acid will dissociate.

Electrolytes that dissociate well into ions are called strong electrolytes, in contrast to weak electrolytes, which form only a small number of ions in aqueous solutions. Solutions of strong electrolytes retain high electrical conductivity even at very high concentrations. Conversely, the electrical conductivity of solutions of weak electrolytes rapidly decreases with increasing concentration. strong electrolytes include acids such as hydrochloric, nitric, sulfuric and some others, then alkalis (except NH 4 OH) and almost all salts.

Polyoonic acids and polyacid bases dissociate in steps. So, for example, sulfuric acid molecules first of all dissociate according to the equation

H 2 SO 4 ⇄ H + HSO 4 '

or more precisely:

H 2 SO 4 + H 2 O ⇄ H 3 O + HSO 4 '

Elimination of the second hydrogen ion according to the equation

HSO 4 ‘⇄ H + SO 4 »

or

HSO 4 '+ H 2 O ⇄ H 3 O + SO 4 "

it is already much more difficult, since it has to overcome the attraction from the doubly charged ion SO 4 ”, which, of course, attracts the hydrogen ion to itself more strongly than the singly charged ion HSO 4 '. Therefore, the second stage of dissociation or, as they say, secondary dissociation occurs in a much smallerdegree than the primary one, and ordinary sulfuric acid solutions contain only a small number of SO 4 ions "

Phosphoric acid H 3 RO 4 dissociates in three steps:

H 3 PO 4 ⇄ H + H 2 PO 4 '

H 2 PO 4 ⇄ H + HPO4 »

HPO 4 » ⇄ H + PO 4 »’

H 3 RO 4 molecules strongly dissociate into H and H 2 RO 4 ions. Ions H 2 PO 4 ' behave like a weaker acid, and dissociate into H and HPO 4 "to a lesser extent. HPO 4 ions, on the other hand, dissociate as a very weak acid, and almost do not give H ions

and PO four "'

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

Va(OH) 2 ⇄ BaOH + OH'

VaOH ⇄ Va + OH'

As for salts, normal salts always dissociate into metal ions and acid residues. For example:

CaCl 2 ⇄ Ca + 2Cl 'Na 2 SO 4 ⇄ 2Na + SO 4 "

Acid salts, like polybasic acids, dissociate in steps. For example:

NaHCO 3 ⇄ Na + HCO 3 '

HCO 3 ‘⇄ H + CO 3 »

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

Basic salts dissociate into ions of basic and acid residues. For example:

Fe(OH)Cl 2 ⇄ FeOH + 2Cl"

The secondary dissociation of ions of the main residues into metal and hydroxyl ions almost does not occur.

In table. 11 shows the numerical values ​​of the degree of dissociation of some acids, bases and salts in 0 , 1 n. solutions.

Decreases with increasing concentration. Therefore, in very concentrated solutions, even strong acids are relatively weakly dissociated. For

Table 11

Acids, bases and salts in 0.1 N.solutions at 18°

Electrolyte	Formula	Degree of dissociation in %
acids
Salt	HCl	92
Hydrobromic	HBr	92
Hydroiodide	HJ	. 92
Nitrogen	HNO3	92
sulfuric	H 2 SO 4	58
sulphurous	H 2SO3	34
Phosphoric	H 3 PO 4	27
Hydrofluoric	HF	8,5
Acetic	CH3COOH	1,3
Coal	H2 CO3	0,17
Hydrogen sulfide	H 2 S	0,07
hydrocyanic	HCN	0,01
Bornaya	H 3 BO 3	0,01
Foundations
barium hydroxide	Ba (OH) 2	92
caustic potash	con	89
Sodium hydroxide	NaON	84
ammonium hydroxide	NH4OH	1,3
salt
Chloride	Kcl	86
Ammonium chloride	NH4Cl	85
Chloride	NaCl	84
Nitrate	KNO 3	83
	AgNO3	81
acetic acid	NaCH 3 COO	79
Chloride	ZnCl 2	73
sulfate	Na 2 SO 4	69
sulfate	ZnSO4	40
Sulfate

Strong and weak electrolytes

In solutions of some electrolytes, only a part of the molecules dissociate. For a quantitative characteristic of the strength of the electrolyte, the concept of the degree of dissociation was introduced. The ratio of the number of molecules dissociated into ions to the total number of molecules of the solute is called the degree of dissociation a.

where C is the concentration of dissociated molecules, mol/l;

C 0 - the initial concentration of the solution, mol / l.

According to the degree of dissociation, all electrolytes are divided into strong and weak. Strong electrolytes include those whose degree of dissociation is greater than 30% (a > 0.3). These include:

strong acids (H 2 SO 4, HNO 3, HCl, HBr, HI);

· soluble hydroxides, except for NH 4 OH;

soluble salts.

Electrolytic dissociation of strong electrolytes proceeds irreversibly

HNO 3 ® H + + NO - 3 .

Weak electrolytes have a dissociation degree of less than 2% (a< 0,02). К ним относятся:

Weak inorganic acids (H 2 CO 3, H 2 S, HNO 2, HCN, H 2 SiO 3, etc.) and all organic, for example, acetic acid (CH 3 COOH);

· insoluble hydroxides, as well as soluble hydroxide NH 4 OH;

insoluble salts.

Electrolytes with intermediate values ​​of the degree of dissociation are called electrolytes of medium strength.

The degree of dissociation (a) depends on the following factors:

on the nature of the electrolyte, that is, on the type of chemical bonds; dissociation most easily occurs at the site of the most polar bonds;

from the nature of the solvent - the more polar the latter, the easier the dissociation process goes in it;

on temperature - an increase in temperature enhances dissociation;

on the concentration of the solution - when the solution is diluted, the dissociation also increases.

As an example of the dependence of the degree of dissociation on the nature of chemical bonds, consider the dissociation of sodium hydrosulfate (NaHSO 4), in the molecule of which there are the following types of bonds: 1-ion; 2 - polar covalent; 3 - the bond between sulfur and oxygen atoms is of low polarity. The rupture occurs most easily at the site of the ionic bond (1):

Na 1 O 3 O S 3 H 2 O O

1. NaHSO 4 ® Na + + HSO - 4, 2. then at the place of the polar bond of a lesser degree: HSO - 4 ® H + + SO 2 - 4. 3. the acid residue does not dissociate into ions.

The degree of electrolyte dissociation strongly depends on the nature of the solvent. For example, HCl strongly dissociates in water, weaker in ethanol C 2 H 5 OH, almost does not dissociate in benzene, in which it practically does not conduct electric current. Solvents with high permittivity (e) polarize solute molecules and form solvated (hydrated) ions with them. At 25 0 С e (H 2 O) \u003d 78.5, e (C 2 H 5 OH) \u003d 24.2, e (C 6 H 6) \u003d 2.27.

In solutions of weak electrolytes, the dissociation process proceeds reversibly and, therefore, the laws of chemical equilibrium are applicable to the equilibrium in solution between molecules and ions. So, for the dissociation of acetic acid

CH 3 COOH « CH 3 COO - + H + .

The equilibrium constant K with will be determined as

K c \u003d K d \u003d CCH 3 COO - · C H + / CCH 3 COOH.

The equilibrium constant (K c) for the dissociation process is called the dissociation constant (K d). Its value depends on the nature of the electrolyte, solvent and temperature, but it does not depend on the concentration of the electrolyte in the solution. The dissociation constant is an important characteristic of weak electrolytes, since it indicates the strength of their molecules in solution. The smaller the dissociation constant, the weaker the electrolyte dissociates and the more stable its molecules. Given that the degree of dissociation, in contrast to the dissociation constant, changes with the concentration of the solution, it is necessary to find a relationship between K d and a. If the initial concentration of the solution is taken equal to C, and the degree of dissociation corresponding to this concentration a, then the number of dissociated molecules of acetic acid will be equal to a C. Since

CCH 3 COO - \u003d C H + \u003d a C,

then the concentration of undecayed acetic acid molecules will be equal to (C - a C) or C (1- a C). From here

K d \u003d aC a C / (C - a C) \u003d a 2 C / (1- a). (one)

Equation (1) expresses the Ostwald dilution law. For very weak electrolytes a<<1, то приближенно К @ a 2 С и

a = (K / C). (2)

As can be seen from formula (2), with a decrease in the concentration of the electrolyte solution (when diluted), the degree of dissociation increases.

Weak electrolytes dissociate in stages, for example:

1 stage H 2 CO 3 "H + + HCO - 3,

2 stage HCO - 3 "H + + CO 2 - 3.

Such electrolytes are characterized by several constants - depending on the number of stages of decomposition into ions. For carbonic acid

K 1 \u003d CH + CHCO - 2 / CH 2 CO 3 \u003d 4.45 × 10 -7; K 2 \u003d CH + · CCO 2- 3 / CHCO - 3 \u003d 4.7 × 10 -11.

As can be seen, the decomposition into carbonic acid ions is determined mainly by the first stage, while the second can manifest itself only when the solution is highly diluted.

The total equilibrium H 2 CO 3 « 2H + + CO 2 - 3 corresponds to the total dissociation constant

K d \u003d C 2 n + · CCO 2- 3 / CH 2 CO 3.

The values ​​of K 1 and K 2 are related to each other by the relation

K d \u003d K 1 K 2.

The bases of multivalent metals dissociate in a similar manner. For example, two steps of dissociation of copper hydroxide

Cu (OH) 2 "CuOH + + OH -,

CuOH + "Cu 2+ + OH -

correspond to dissociation constants

K 1 \u003d CCuOH + SON - / CCu (OH) 2 and K 2 \u003d Ccu 2+ SON - / CCuOH +.

Since strong electrolytes are completely dissociated in solution, the very term dissociation constant for them is meaningless.

Dissociation of various classes of electrolytes

From the point of view of the theory of electrolytic dissociation acid a substance is called, during the dissociation of which only a hydrated hydrogen ion H 3 O (or simply H +) is formed as a cation.

foundation A substance is called a substance that, in an aqueous solution, forms OH hydroxide ions as an anion and no other anions.

According to Bronsted's theory, an acid is a proton donor and a base is a proton acceptor.

The strength of bases, like the strength of acids, depends on the value of the dissociation constant. The larger the dissociation constant, the stronger the electrolyte.

There are hydroxides that can interact and form salts not only with acids, but also with bases. Such hydroxides are called amphoteric. These include Be(OH) 2 , Zn(OH) 2 , Sn(OH) 2 , Pb(OH) 2 , Cr(OH) 3 , Al(OH) 3. Their properties are due to the fact that they dissociate to a weak degree according to the type of acids and the type of bases.

H++RO- « ROH « R + + OH -.

This equilibrium is explained by the fact that the strength of the bond between metal and oxygen differs slightly from the strength of the bond between oxygen and hydrogen. Therefore, when beryllium hydroxide reacts with hydrochloric acid, beryllium chloride is obtained

Be (OH) 2 + HCl \u003d BeCl 2 + 2H 2 O,

and when interacting with sodium hydroxide - sodium beryllate

Be (OH) 2 + 2NaOH \u003d Na 2 BeO 2 + 2H 2 O.

salt can be defined as electrolytes that dissociate in solution to form cations other than hydrogen cations and anions other than hydroxide ions.

Medium salts, obtained with the complete replacement of hydrogen ions of the corresponding acids with metal cations (or NH + 4), completely dissociate Na 2 SO 4 "2Na + + SO 2- 4.

Acid salts dissociate in steps

1 stage NaHSO 4 « Na + + HSO - 4 ,

2 stage HSO - 4 "H + + SO 2-4.

The degree of dissociation in the 1st stage is greater than in the 2nd stage, and the weaker the acid, the lower the degree of dissociation in the 2nd stage.

basic salts, obtained by incomplete replacement of hydroxide ions with acidic residues, also dissociate in steps:

1 step (CuOH) 2 SO 4 "2 CuOH + + SO 2- 4,

2 stage CuOH + "Cu 2+ + OH -.

Basic salts of weak bases dissociate mainly in the 1st step.

complex salts, containing a complex complex ion that retains its stability upon dissolution, dissociate into a complex ion and ions of the outer sphere

K 3 « 3K + + 3 - ,

SO 4 "2+ + SO 2 - 4.

In the center of the complex ion is an atom - the complexing agent. This role is usually performed by metal ions. Near the complexing agents are located (coordinated) polar molecules or ions, and sometimes both together, they are called ligands. The complexing agent, together with the ligands, constitutes the inner sphere of the complex. Ions located far from the complexing agent, less strongly associated with it, are in the external environment of the complex compound. The inner sphere is usually enclosed in square brackets. The number indicating the number of ligands in the inner sphere is called coordinating. Chemical bonds between complex and simple ions are relatively easily broken in the process of electrolytic dissociation. The bonds leading to the formation of complex ions are called donor-acceptor bonds.

The ions of the outer sphere are easily split off from the complex ion. This dissociation is called primary. The reversible disintegration of the inner sphere is much more difficult and is called secondary dissociation.

Cl " + + Cl - - primary dissociation,

+ « Ag + +2 NH 3 - secondary dissociation.

secondary dissociation, like the dissociation of a weak electrolyte, is characterized by an instability constant

To nest. \u003d × 2 / [ + ] \u003d 6.8 × 10 -8.

The instability constants (K inst.) of various electrolytes is a measure of the stability of the complex. The less K nest. , the more stable the complex.

So, among the same type of compounds:

-	+	+	+
K nest \u003d 1.3 × 10 -3	K nest \u003d 6.8 × 10 -8	K nest \u003d 1 × 10 -13	K nest \u003d 1 × 10 -21

the stability of the complex increases with the transition from - to + .

The values ​​of the instability constant are given in reference books on chemistry. Using these values, one can predict the course of reactions between complex compounds with a strong difference in the instability constants, the reaction will go towards the formation of a complex with a lower instability constant.

A complex salt with an unstable complex ion is called double salt. Double salts, unlike complex ones, dissociate into all the ions that make up their composition. For example:

KAl(SO 4) 2 "K + + Al 3+ + 2SO 2- 4,

NH 4 Fe (SO 4) 2 "NH 4 + + Fe 3+ + 2SO 2- 4.

SOLUTIONS
THEORY OF ELECTROLYTIC DISSOCIATION

ELECTROLYTIC DISSOCIATION
ELECTROLYTES AND NON-ELECTROLYTES

Theory of electrolytic dissociation

(S. Arrhenius, 1887)

1. When dissolved in water (or melted), electrolytes decompose into positively and negatively charged ions (subject to electrolytic dissociation).

2. Under the action of an electric current, cations (+) move towards the cathode (-), and anions (-) move towards the anode (+).

3. Electrolytic dissociation is a reversible process (the reverse reaction is called molarization).

4. Degree of electrolytic dissociation ( a ) depends on the nature of the electrolyte and solvent, temperature and concentration. It shows the ratio of the number of molecules decomposed into ions ( n ) to the total number of molecules introduced into the solution ( N).

a = n / N0< a <1

Mechanism of electrolytic dissociation of ionic substances

When dissolving compounds with ionic bonds ( e.g. NaCl ) the hydration process begins with the orientation of water dipoles around all ledges and faces of salt crystals.

Orienting around the ions of the crystal lattice, water molecules form either hydrogen or donor-acceptor bonds with them. This process releases a large amount of energy, which is called hydration energy.

The energy of hydration, the value of which is comparable to the energy of the crystal lattice, goes to the destruction of the crystal lattice. In this case, hydrated ions pass layer by layer into the solvent and, mixing with its molecules, form a solution.

Mechanism of electrolytic dissociation of polar substances

Substances whose molecules are formed according to the type of polar covalent bond (polar molecules) also dissociate similarly. Around each polar molecule of matter ( e.g. HCl ), the dipoles of water are oriented in a certain way. As a result of interaction with water dipoles, the polar molecule becomes even more polarized and turns into an ionic molecule, then free hydrated ions are easily formed.

Electrolytes and non-electrolytes

The electrolytic dissociation of substances, proceeding with the formation of free ions, explains the electrical conductivity of solutions.

The process of electrolytic dissociation is usually written in the form of a diagram, without revealing its mechanism and omitting the solvent ( H2O ), although he is a major contributor.

CaCl 2 "Ca 2+ + 2Cl -

KAl(SO 4) 2 "K + + Al 3+ + 2SO 4 2-

HNO 3 "H + + NO 3 -

Ba (OH) 2 "Ba 2+ + 2OH -

From the electrical neutrality of molecules it follows that the total charge of cations and anions must be equal to zero.

For example, for

Al 2 (SO 4) 3 ––2 (+3) + 3 (-2) = +6 - 6 = 0

KCr(SO 4) 2 ––1 (+1) + 3 (+3) + 2 (-2) = +1 + 3 - 4 = 0

Strong electrolytes

These are substances that, when dissolved in water, almost completely decompose into ions. As a rule, strong electrolytes include substances with ionic or highly polar bonds: all highly soluble salts, strong acids ( HCl, HBr, HI, HClO 4, H 2 SO 4, HNO 3 ) and strong bases ( LiOH, NaOH, KOH, RbOH, CsOH, Ba (OH) 2, Sr (OH) 2, Ca (OH) 2).

In a solution of a strong electrolyte, the solute is found mainly in the form of ions (cations and anions); undissociated molecules are practically absent.

Weak electrolytes

Substances that partially dissociate into ions. Solutions of weak electrolytes, along with ions, contain undissociated molecules. Weak electrolytes cannot give a high concentration of ions in solution.

Weak electrolytes include:

1) almost all organic acids ( CH 3 COOH, C 2 H 5 COOH, etc.);

2) some inorganic acids ( H 2 CO 3 , H 2 S, etc.);

3) almost all water-soluble salts, bases and ammonium hydroxide(Ca 3 (PO 4 ) 2 ; Cu (OH ) 2 ; Al (OH ) 3 ; NH 4 OH ) ;

4) water.

They poorly (or almost do not conduct) electricity.

CH 3 COOH « CH 3 COO - + H +

Cu (OH) 2 "[CuOH] + + OH - (first stage)

[CuOH] + "Cu 2+ + OH - (second step)

H 2 CO 3 "H + + HCO - (first stage)

HCO 3 - "H + + CO 3 2- (second stage)

Non-electrolytes

Substances whose aqueous solutions and melts do not conduct electricity. They contain covalent non-polar or low-polar bonds that do not break down into ions.

Gases, solids (non-metals), organic compounds (sucrose, gasoline, alcohol) do not conduct electric current.

Degree of dissociation. Dissociation constant

The concentration of ions in solutions depends on how completely the given electrolyte dissociates into ions. In solutions of strong electrolytes, the dissociation of which can be considered complete, the concentration of ions can be easily determined from the concentration (c) and the composition of the electrolyte molecule (stoichiometric indices), for example :

Ion concentrations in solutions of weak electrolytes are qualitatively characterized by the degree and dissociation constant.

Degree of dissociation (a) is the ratio of the number of molecules decayed into ions ( n ) to the total number of dissolved molecules ( N):

a = n / N

and is expressed in fractions of a unit or in% ( a \u003d 0.3 - conditional division boundary into strong and weak electrolytes).

Example

Determine the molar concentration of cations and anions in 0.01 M solutions KBr, NH 4 OH, Ba (OH) 2, H 2 SO 4 and CH 3 COOH.

The degree of dissociation of weak electrolytes a = 0.3.

Solution

KBr, Ba (OH) 2 and H 2 SO 4 - strong electrolytes that dissociate completely(a = 1).

KBr « K + + Br -

0.01M

Ba (OH) 2 "Ba 2+ + 2OH -

0.01M

0.02M

H 2 SO 4 "2H + + SO 4

0.02M

[SO 4 2-] = 0.01 M

NH 4 OH and CH 3 COOH - weak electrolytes(a=0.3)

NH 4 OH + 4 + OH -

0.3 0.01 = 0.003M

CH 3 COOH « CH 3 COO - + H +

[H +] \u003d [CH 3 COO -] \u003d 0.3 0.01 \u003d 0.003 M

The degree of dissociation depends on the concentration of the weak electrolyte solution. When diluted with water, the degree of dissociation always increases, because the number of solvent molecules increases ( H2O ) per solute molecule. According to Le Chatelier's principle, the equilibrium of electrolytic dissociation in this case should shift in the direction of product formation, i.e. hydrated ions.

The degree of electrolytic dissociation depends on the temperature of the solution. Usually, with increasing temperature, the degree of dissociation increases, because bonds in molecules are activated, they become more mobile and easier to ionize. The concentration of ions in a weak electrolyte solution can be calculated knowing the degree of dissociationaand the initial concentration of the substancec in solution.

Example

Determine the concentration of non-dissociated molecules and ions in a 0.1 M solution NH4OH if the degree of dissociation is 0.01.

Solution

Molecule concentrations NH4OH , which will decay into ions by the moment of equilibrium, will be equal toac. Ion concentration NH 4 - and OH - - will be equal to the concentration of dissociated molecules and equal toac(according to the electrolytic dissociation equation)

NH4OH		NH4+		oh-
c - a c

A c = 0.01 0.1 = 0.001 mol/l

[NH 4 OH] \u003d c - a c \u003d 0.1 - 0.001 \u003d 0.099 mol / l

Dissociation constant ( KD ) is the ratio of the product of the equilibrium concentrations of ions to the power of the corresponding stoichiometric coefficients to the concentration of undissociated molecules.

It is the equilibrium constant of the process of electrolytic dissociation; characterizes the ability of a substance to decompose into ions: the higher KD , the greater the concentration of ions in the solution.

The dissociations of weak polybasic acids or polyacid bases proceed in stages, respectively, for each stage there is its own dissociation constant:

First stage:

H 3 PO 4 « H + + H 2 PO 4 -

K D 1 = () / = 7.1 10 -3

Second step:

H 2 PO 4 - « H + + HPO 4 2-

K D 2 = () / = 6.2 10 -8

Third step:

HPO 4 2- « H + + PO 4 3-

K D 3 = () / = 5.0 10 -13

K D 1 > K D 2 > K D 3

Example

Get an equation relating the degree of electrolytic dissociation of a weak electrolyte ( a ) with dissociation constant (Ostwald's dilution law) for a weak monobasic acid ON THE .

HA «H++A+

K D = () /

If the total concentration of a weak electrolyte is denotedc, then the equilibrium concentrations H + and A - are equal ac, and the concentration of undissociated molecules ON - (c - a c) \u003d c (1 - a)

K D \u003d (a c a c) / c (1 - a) \u003d a 2 c / (1 - a)

In the case of very weak electrolytes ( a £ 0.01 )

K D = c a 2 or a = \ é (K D / c )

Example

Calculate the degree of dissociation of acetic acid and the concentration of ions H + in 0.1 M solution if K D (CH 3 COOH) = 1.85 10 -5

Solution

Let's use the Ostwald dilution law

\ é (K D / c ) = \ é ((1.85 10 -5) / 0.1 )) = 0.0136 or a = 1.36%

[ H + ] \u003d a c \u003d 0.0136 0.1 mol / l

Solubility product

Definition

Put some sparingly soluble salt into a beaker, e.g. AgCl and add distilled water to the precipitate. At the same time, ions Ag+ and Cl- , experiencing attraction from the surrounding dipoles of water, gradually break away from the crystals and go into solution. Colliding in solution, ions Ag+ and Cl- form molecules AgCl and deposited on the crystal surface. Thus, two mutually opposite processes occur in the system, which leads to dynamic equilibrium, when the same number of ions pass into the solution per unit time Ag+ and Cl- how many are deposited. Ion accumulation Ag+ and Cl- stops in solution, it turns out saturated solution. Therefore, we will consider a system in which there is a precipitate of a sparingly soluble salt in contact with a saturated solution of this salt. In this case, two mutually opposite processes take place:

1) The transition of ions from the precipitate to the solution. The rate of this process can be considered constant at a constant temperature: V 1 = K 1 ;

2) Precipitation of ions from solution. The speed of this process V 2 depends on ion concentration Ag + and Cl - . According to the law of mass action:

V 2 \u003d k 2

Since the system is in equilibrium, then

V1 = V2

k2 = k1

K 2 / k 1 = const (at T = const)

In this way, the product of ion concentrations in a saturated solution of a sparingly soluble electrolyte at a constant temperature is constant magnitude. This value is calledsolubility product(ETC ).

In the given example ETC AgCl = [Ag+][Cl-] . In cases where the electrolyte contains two or more identical ions, the concentration of these ions must be raised to the appropriate power when calculating the solubility product.

For example , PR Ag 2 S = 2 ; PR PbI 2 = 2

In the general case, the expression for the solubility product for an electrolyte is A m B n

PR A m B n = [A] m [B] n .

The values ​​of the solubility product for different substances are different.

For example , PR CaCO 3 = 4.8 10 -9 ; PR AgCl \u003d 1.56 10 -10.

ETC easy to calculate, knowing c creativity of the compound at a given t°.

Example 1

The solubility of CaCO 3 is 0.0069 or 6.9 10 -3 g/l. Find PR CaCO 3 .

Solution

We express the solubility in moles:

S CaCO 3 = ( 6,9 10 -3 ) / 100,09 = 6.9 10 -5 mol/l

M CaCO3

Since every molecule CaCO3 gives one ion each when dissolved Ca 2+ and CO 3 2-, then
[ Ca 2+ ] \u003d [ CO 3 2- ] \u003d 6.9 10 -5 mol / l ,
Consequently, PR CaCO 3 \u003d [ Ca 2+ ] [ CO 3 2- ] \u003d 6.9 10 -5 6.9 10 -5 \u003d 4.8 10 -9

Knowing the value of PR , you can in turn calculate the solubility of the substance in mol / l or g / l.

Example 2

Solubility product PR PbSO 4 \u003d 2.2 10 -8 g / l.

What is the solubility PbSO4?

Solution

Denote the solubility PbSO 4 via X mol/l. Going into solution X moles of PbSO 4 will give X Pb 2+ ions and X ionsSO 4 2- , i.e.:

== X

ETCPbSO 4 = = = X X = X 2

X=\ é(ETCPbSO 4 ) = \ é(2,2 10 -8 ) = 1,5 10 -4 mol/l.

To go to the solubility, expressed in g / l, we multiply the found value by the molecular weight, after which we get:

1,5 10 -4 303,2 = 4,5 10 -2 g/l.

Precipitation formation

If a

[ Ag + ] [ Cl - ] < ПР AgCl- unsaturated solution

[ Ag + ] [ Cl - ] = PRAgCl- saturated solution

[ Ag + ] [ Cl - ] > PRAgCl- supersaturated solution

A precipitate is formed when the product of the ion concentrations of a sparingly soluble electrolyte exceeds the value of its solubility product at a given temperature. When the ion product becomes equal toETC, precipitation stops. Knowing the volume and concentration of the mixed solutions, it is possible to calculate whether the resulting salt will precipitate.

Example 3

Does a precipitate form when mixing equal volumes of 0.2MsolutionsPb(NO 3 ) 2 andNaCl.
ETCPbCl 2 = 2,4 10 -4 .

Solution

When mixed, the volume of the solution doubles and the concentration of each of the substances decreases by half, i.e. will become 0.1 M or 1.0 10 -1 mol/l. These are there will be concentrationsPb 2+ andCl - . Consequently,[ Pb 2+ ] [ Cl - ] 2 = 1 10 -1 (1 10 -1 ) 2 = 1 10 -3 . The resulting value exceedsETCPbCl 2 (2,4 10 -4 ) . So part of the saltPbCl 2 precipitates out. From all of the above, we can conclude that various factors influence the formation of precipitation.

Influence of the concentration of solutions

Sparingly soluble electrolyte with a sufficiently large valueETCcannot be precipitated from dilute solutions.For example, precipitatePbCl 2 will not fall out when mixing equal volumes 0.1MsolutionsPb(NO 3 ) 2 andNaCl. When mixing equal volumes, the concentrations of each of the substances will become0,1 / 2 = 0,05 Mor 5 10 -2 mol/l. Ionic product[ Pb 2+ ] [ Cl 1- ] 2 = 5 10 -2 (5 10 -2 ) 2 = 12,5 10 -5 .The resulting value is lessETCPbCl 2 hence no precipitation will occur.

Influence of the amount of precipitant

For the most complete precipitation, an excess of precipitant is used.

For example, precipitate saltBaCO 3 : BaCl 2 + Na 2 CO 3 ® BaCO 3 ¯ + 2 NaCl. After adding an equivalent amountNa 2 CO 3 ions remain in solutionBa 2+ , whose concentration is determined by the quantityETC.

Increasing the concentration of ionsCO 3 2- caused by the addition of excess precipitant(Na 2 CO 3 ) , will entail a corresponding decrease in the concentration of ionsBa 2+ in solution, i.e. will increase the completeness of the deposition of this ion.

Influence of the ion of the same name

The solubility of sparingly soluble electrolytes decreases in the presence of other strong electrolytes having similar ions. If to an unsaturated solutionBaSO 4 add solution little by littleNa 2 SO 4 , then the ionic product, which was initially less than ETCBaSO 4 (1,1 10 -10 ) , will gradually reachETCand exceed it. Precipitation will begin.

Temperature effect

ETCis constant at constant temperature. With increasing temperature ETC increases, so precipitation is best done from cooled solutions.

Dissolution of precipitation

The solubility product rule is important for transferring sparingly soluble precipitates into solution. Suppose we need to dissolve the precipitateBaFROMO 3 . The solution in contact with this precipitate is saturated withBaFROMO 3 .
It means that[ Ba 2+ ] [ CO 3 2- ] = PRBaCO 3 .

If an acid is added to the solution, then the ionsH + bind the ions present in the solutionCO 3 2- into weak carbonic acid molecules:

2H + + CO 3 2- ® H 2 CO 3 ® H 2 O+CO 2 ­

As a result, the concentration of the ion will sharply decrease.CO 3 2- , the ion product becomes less thanETCBaCO 3 . The solution will be unsaturated with respect toBaFROMO 3 and part of the sedimentBaFROMO 3 goes into solution. With the addition of a sufficient amount of acid, the entire precipitate can be brought into solution. Consequently, the dissolution of the precipitate begins when, for some reason, the ion product of a sparingly soluble electrolyte becomes less thanETC. In order to dissolve the precipitate, an electrolyte is introduced into the solution, the ions of which can form a slightly dissociated compound with one of the ions of a sparingly soluble electrolyte. This explains the dissolution of sparingly soluble hydroxides in acids.

Fe(OH) 3 + 3HCl® FeCl 3 + 3H 2 O

ionsOh - bind into poorly dissociated moleculesH 2 O.

Table.Solubility product (SP) and solubility at 25AgCl

1,25 10 -5

1,56 10 -10

AgI

1,23 10 -8

1,5 10 -16

Ag 2 CrO4

1,0 10 -4

4,05 10 -12

BaSO4

7,94 10 -7

6,3 10 -13

CaCO3

6,9 10 -5

4,8 10 -9

PbCl 2

1,02 10 -2

1,7 10 -5

PbSO 4

1,5 10 -4

2,2 10 -8