Degree of dissociation. Strong and weak electrolytes. About the classification of electrolytes

If electrolytes completely dissociated into ions, then the osmotic pressure (and other quantities proportional to it) would always be an integer number of times greater than the values ​​observed in non-electrolyte solutions. But van't Hoff also established that the coefficient i

expressed in fractional numbers, which increase with dilution of the solution, approaching whole numbers.

Arrhenius explained this fact by the fact that only a part of the electrolyte dissociates into ions in solution, and introduced the concept of the degree of dissociation. The degree of dissociation of an electrolyte is the ratio of the number of its molecules decomposed into ions in a given solution to the total number of its molecules in the solution.

Later it was found that electrolytes can be divided into two groups: strong and weak electrolytes. Strong electrolytes in aqueous solutions are almost completely dissociated. The concept of the degree of dissociation is essentially inapplicable to them, and the deviation of the isotonic coefficient i

from integer values ​​due to other reasons. Weak electrolytes in aqueous solutions dissociate only partially. Therefore, the number of ions in solutions of strong electrolytes is greater than in solutions of weak ones of the same concentration. And if in solutions of weak electrolytes FROM

ions is small, the distance between them is large and the interaction of ions with each other is insignificant, then in not very dilute solutions of strong electrolytes, the average distance between ions due to a significant concentration is relatively small. In such solutions, the ions are not completely free, their movement is constrained by mutual attraction to each other. Due to this attraction, each ion is, as it were, surrounded by a spherical swarm of oppositely charged ions, called the "ionic atmosphere".

All salts belong to strong electrolytes; the most important acids and bases are HNO3, H2SO4, HClO4, HCl, HBr, HI, KOH, NaOH, Ba(OH)2, and Ca(OH)2.

Weak electrolytes include most organic acids, and among the most important inorganic compounds, they include H2CO3, H2S, HCN, H2SiO3, and NH4OH.

The degree of dissociation is usually denoted by the Greek letter a and expressed either in fractions of a unit or as a percentage.

The process of electrolytic dissociation is reversible, therefore, in the electrolyte solution, along with its ions, there are also molecules. The ratio of the content of these particles is determined by the degree of electrolytic dissociation, which is a quantitative characteristic of the dissociation process.

Degree of dissociation(α) is the ratio of the number of electrolyte molecules decomposed into ions ( n ) to the total number of dissolved molecules ( ):

The degree of dissociation is determined empirically and is expressed in fractions of a unit or as a percentage:

If α = 0, then there is no dissociation. If α = 100%, then the electrolyte completely decomposes into ions. If α = 1.3%, then out of 1000 electrolyte molecules only 13 dissociate into ions.

Factors affecting the degree of electrolytic dissociation:

1. The nature of the electrolyte: the polarity of the chemical bond in the compound, the increase of which contributes to the increase in α.

2. Solution concentration: α increases with decreasing solution concentration.

3. Temperature: α increases with increasing solution temperature.

All electrolytes are usually divided into 3 groups according to the degree of electrolytic dissociation: strong, weak and medium strength (Table 7.1.).

When writing dissociation equations, the strength of the electrolyte should be taken into account. According to the theory of electrolytic dissociation, strong electrolytes dissociate in one stage into ions, which make up the electrolyte molecule. For example:

H 2 SO 4 ↔ 2H + + SO 4 2-.

Weak electrolytes dissociate stepwise, with the ions of the first stage (step) predominating. For example:

Stage I H 2 S ↔ H + + NS -

Stage II HS - ↔ H + + S 2-.

In solutions of weak electrolytes, chemical equilibrium always takes place, which is expressed in the equality of the rates of the dissociation and association reactions. Using the law of mass action (6.8.), for such electrolytes, the equilibrium can be quantitatively expressed by the value of the dissociation constant (K diss)[‡] . For example, for electrolyte HA ↔ H + + A - dissociation constant:

. (7.10)

Table 7.1.

Classification of electrolytes depending on the value of α[§]

Ionic equations

According to the theory of electrolytic dissociation, all reactions in aqueous electrolyte solutions are reactions between ions. They're called ionic reactions , and the equations of these reactions are ionic equations .

When studying the processes occurring in electrolyte solutions, one should be guided by the following rule:

Reactions between ions in electrolyte solutions go almost to the end in the direction of the formation of precipitates, gases and weak electrolytes.

In ionic equations, it is customary to write in a form that is not dissociated into ions (in the form of molecules) the formulas of sparingly soluble compounds, non-electrolytes, electrolytes of weak and medium strength. The record of ionic reactions can be represented in the form of molecular, full and reduced ionic equations. When writing the equation, the sign ↓ at the formula means that the substance is removed from the reaction sphere in the form of an insoluble compound, the sign shows that the substance is released in the form of a gas.

molecular equation BaCl 2 + H 2 SO 4 = BaSO 4 ↓ + 2HCl

complete ionic equation for Ba 2+ + 2Cl- + 2H++ SO 4 2- = BaSO 4 ↓ + 2H+ + 2Cl-

abbreviated ionic equation Ba 2+ + SO 4 2- = BaSO 4 ↓.

When writing ionic equations, we use the data from the table "Solubility of salts, acids and bases in water" (Appendix, Table 4.).

1 Strong electrolytes include electrolytes in which α > 30%:

a) all alkalis (bases formed by s-family metals, with the exception of beryllium and magnesium): LiOH, NaOH, KOH, RbOH, CsOH, FrOH, Ca (OH) 2, Ba (OH) 2 - practically completely dissociate.

Foundations dissociate in solution to form a metal cation and hydroxide ions.

NaOH → Na + + OH - ;

Ba(OH) 2 → BaOH + + OH - .

The dissociation of many strong electrolytes in the second stage is not as active as in the first. Therefore, this process can be written as follows:

BaOH + Ba 2 + + OH - .

The overall equation of the process:

Ba (OH) 2 → Ba 2 + + 2OH -.

b) some acids, for example: HCl, HClO 4, HBr, HJ, HNO 3, H 2 SO 4.

Acids dissociate in solution with the formation of hydrogen ions and ions of acid residues (because the bond between the hydrogen cation and the acid residue is more polar than between the ions in the acid residue itself).

HCl → H + + Cl – ;

H 2 SO 4 → H + + HSO 4 -;

HSO 4 - H + + SO 4 2– (sulfuric acid dissociates worse in the second stage than in the first, so the reversibility sign “ ” is put).

Conventionally, the overall equation of the process can be written as:

H 2 SO 4 → 2H + + SO 4 2–.

c) soluble salts (α ~ 100%)

In salts, the constituents are metal atoms and acidic residues. It is into these ions that salts decompose when melted or dissolved in water.

Na 3 PO 3 → 3Na + + PO 3 3 – .

d) acidic, basic and complex salts during dissociation in the first stage.

Acid salts decompose into metal cations and anions of the acid residue:

K 2 HRO 3 → 2K + + HRO 3 2 - .

Moreover, according to the principle of electrostatic attraction, the hydrogen ion (s) (H +) remains next to the anion of the acid residue (KO n -), and not to the metal cation (Me n +).

Basic salts decompose into metal cations associated with the hydroco group and anions of the acid residue:

Al(OH) 2 Cl → Al(OH) 2 + + Cl – .

Moreover, according to the principle of electrostatic attraction, the hydroxo group (s) (OH -) remains next to the metal cation (Me n +), and not to the anion of the acid residue (KO n -).

Complex salts decompose into ions of the outer sphere and a complex ion (since the bond between the ion of the outer sphere and the complex ion, as a rule, is covalent polar or ionic, and between ions or molecules in the complex ion itself is more often donor-acceptor).

K 3 → 3K + + 3– .

2 Weak electrolytes dissociate poorly, their α< 3%.

The dissociation of weak electrolytes proceeds reversibly, and if three or more ions are formed during the decay of a molecule, then it is also stepwise.

Weak electrolytes include:

a) all other grounds:

NH 4 OH NH 4 + + OH – ;

Be(OH) 2 BeOH + + OH – ;

BeOH + Be 2 + + OH – ;

b) most other acids:

HCN H + + CN – ;

H 2 CO 3 H + + HCO 3 -;

HCO 3 - H + + CO 3 2–;

c) all insoluble soluble salts:

AgCl Ag + + Cl – ;

BaSO 4 Ba 2 + + SO 4 2– ;

d) acidic, basic and complex salts during dissociation in the second and subsequent stages (the first stage, as we remember, proceeds irreversibly).

Acid salt:

K 2 HRO 3 → 2K + + HRO 3 2 – ;

HRO 3 2 – H + + RO 3 3 – .

The number of steps following the first is determined by the number of hydrogen ions remaining near the acid residue.

Basic salt:

Al(OH) 2 Cl → Al(OH) 2 + + Cl – ;

Al(OH) 2 + AlOH 2 + + OH - ;

AlOH + Al 3 + + OH - .

The number of stages following the first is determined by the number of hydroxo groups remaining near the metal cation.

Complex salts:

3– Fe 3+ + 6CN – .

3 Electrolytes of medium strength have α from 3% to 30%

1.1.5 Dissociation constant. The process of dissociation of weak electrolytes is reversible and there is a dynamic equilibrium in the system, which can be described by an equilibrium constant expressed in terms of the concentrations of the formed ions and non-dissociated molecules, called the dissociation constant. Those. the electrolytic dissociation constant is nothing but the chemical equilibrium constant applicable to the decomposition of a weak electrolyte. For some electrolyte that decomposes into ions in solution in accordance with the equation:

A a B b aA x + + bB y –

the dissociation constant is expressed by the following relationship:

The dissociation constant (K D or simply K) 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 K D, the greater the concentration of ions in the solution.

In polybasic acids and polyacid bases, dissociation occurs in steps, and each step is characterized by its own degree of dissociation. So, phosphoric acid dissociates in three steps (table 1).

Table 1 - Dissociation of phosphoric acid

As you can see, K D 1 > K D 2 > K D 3 . Consequently, the most complete dissociation proceeds through the first stage, since: 1) it is easier to detach an ion from a neutral molecule than from a charged ion: hydrogen ions H + are much stronger attracted to a three-charged PO 3– ion and a doubly charged HPO 2– ion than to a singly charged HPO - ; 2) the dissociation occurring in the second and subsequent stages is suppressed by the ions formed during the decay of the molecule in the first stage (the dissociation equilibrium shifts to the left due to the ions of the same name, in the case of phosphoric acid, hydrogen ions).

It follows that the decomposition of the electrolyte into ions proceeds mainly through the first stage, and in the solution of orthophosphoric acid there will be mainly H + and HPO 2– ions.

1.1.6 Relationship between the dissociation constant and the degree of dissociation. Ostwald's dilution law. Let us write again the dissociation equation for a binary compound related to weak electrolytes:

AB A + + B - .

We write the expression for its dissociation constant:

K =
.

If the total concentration of a weak electrolyte is denoted FROM, then the equilibrium concentrations A + and B - are α ·FROM, and the concentration of undissociated AB molecules is ( FROM – α ·FROM) = (1 – α )∙FROM. Then expression (2) in this case can be rewritten as follows:

Thus, the degree of dissociation of a weak electrolyte is inversely proportional to the concentration and directly proportional to the dilution of the solution; expression (5) is called the Ostwald dilution law: the degree of dissociation of a weak electrolyte in a solution is higher, the more dilute the solution.

1.1.7 Shift in the equilibrium of dissociation of a weak electrolyte. Equilibrium in electrolyte solutions, like any chemical equilibrium, remains unchanged until the conditions that determine it change, and a change in conditions entails a violation of the equilibrium.

Thus, the equilibrium is disturbed when the concentration of one of the ions participating in this equilibrium changes: when it increases, a process occurs during which these ions bind. For example, if in a solution of hypochlorous acid, which is a weak electrolyte and reversibly dissociates according to the scheme

HClO H + + Cl –

introduce any salt of this acid, which is a strong electrolyte and dissociates irreversibly (for example, NaCl → Na + + Cl -) and thereby increase the concentration of Cl - ions, then, in accordance with the Le Chatelier principle, the equilibrium shifts to the left, i.e. The degree of dissociation of hypochlorous acid decreases. It follows from this that the introduction of ions of the same name into a solution of a weak electrolyte (i.e., ions identical with one of the ions of the electrolyte) reduces the degree of dissociation of this electrolyte . In this case, a decrease in the degree of dissociation of hypochlorous acid will also occur if any strong acid containing H + hydrogen ions is added to it.

On the contrary, a decrease in the concentration of one of the ions causes the dissociation of a new number of molecules. For example, when hydroxide ions (formed during dissociation, for example, NaOH → Na + + OH –) that bind hydrogen ions are introduced into the solution of the indicated acid, the dissociation of the acid increases due to the shift of the dissociation equilibrium to the right.

Based on the examples considered, a general conclusion can be drawn. A prerequisite for reactions between electrolytes is the removal of certain ions from the solution, for example, due to the formation of weakly dissociating substances or substances escaping from the solution in the form of a precipitate or gas. In other words, reactions in electrolyte solutions always go in the direction of the formation of the least dissociated or least soluble substances.

1. ELECTROLYTES

1.1. electrolytic dissociation. Degree of dissociation. The strength of electrolytes

According to the theory of electrolytic dissociation, salts, acids, hydroxides, dissolving in water, completely or partially decompose into independent particles - ions.

The process of disintegration of molecules of substances into ions under the action of polar solvent molecules is called electrolytic dissociation. Substances that dissociate into ions in solution are called electrolytes. As a result, the solution acquires the ability to conduct an electric current, because. mobile carriers of electric charge appear in it. According to this theory, when dissolved in water, electrolytes decompose (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 acid residues and hydroxide ions.

For a quantitative characteristic of the dissociation process, the concept of the degree of dissociation is introduced. The degree of dissociation of an electrolyte (α) is the ratio of the number of its molecules decomposed 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 decomposed 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 - hydroxides of metals of the first group of the main subgroup (alkalis) - 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 table of solubility).

Weak electrolytes dissociate into ions to a very small extent, in solutions they are 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 ( NH4OH);

4) most organic acids

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

5) insoluble and sparingly soluble salts and hydroxides of certain metals (see table of solubility).

Process electrolytic dissociation depicted using chemical equations. For example, the 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 in steps. For example,

H 2 CO 3 H + + HCO 3 -,

HCO 3 - H + + CO 3 2–.

The first equilibrium - dissociation along the first stage - is characterized by a constant

.

For dissociation in the second step:

.

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 detach an ion is minimal when it is detached from a neutral molecule.

Medium (normal) salts, soluble in water, dissociate with the formation of 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) - electrolytes containing hydrogen in the anion, capable of splitting off in the form of a 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. The dissociation of acid salts occurs in stages, for example:

KHCO3 → K + + HCO 3 - (first stage)