Highly water-soluble electrolytes are examples. Theory of electrolytic dissociation

All substances can be divided into electrolytes and non-electrolytes. Electrolytes include substances whose solutions or melts conduct electric current (for example, aqueous solutions or melts of KCl, H 3 PO 4 , Na 2 CO 3). Non-electrolyte substances do not conduct electric current when melted or dissolved (sugar, alcohol, acetone, etc.).

Electrolytes are divided into strong and weak. Strong electrolytes in solutions or melts completely dissociate into ions. When writing the equations of chemical reactions, this is emphasized by an arrow in one direction, for example:

HCl → H + + Cl -

Ca (OH) 2 → Ca 2+ + 2OH -

Strong electrolytes include substances with a heteropolar or ionic crystal structure (table 1.1).

Table 1.1 Strong electrolytes

Weak electrolytes decompose into ions only partially. Along with ions, in melts or solutions of these substances, the vast majority of non-dissociated molecules are present. In solutions of weak electrolytes, in parallel with dissociation, the reverse process proceeds - association, that is, the combination of ions into molecules. When writing the reaction equation, this is emphasized by two oppositely directed arrows.

CH 3 COOH D CH 3 COO - + H +

Weak electrolytes include substances with a homeopolar type of crystal lattice (table 1.2).

Table 1.2 Weak electrolytes

The equilibrium state of a weak electrolyte in an aqueous solution is quantitatively characterized by the degree of electrolytic dissociation and the electrolytic dissociation constant.

The degree of electrolytic dissociation α is the ratio of the number of molecules decomposed into ions to the total number of dissolved electrolyte molecules:

The degree of dissociation shows what part of the total amount of the dissolved electrolyte decomposes into ions and depends on the nature of the electrolyte and solvent, as well as on the concentration of the substance in the solution, has a dimensionless value, although it is usually expressed as a percentage. With infinite dilution of the electrolyte solution, the degree of dissociation approaches unity, which corresponds to the complete, 100%, dissociation of the solute molecules into ions. For solutions of weak electrolytes α<<1. Сильные электролиты в растворах диссоциируют полностью (α =1). Если известно, что в 0,1 М растворе уксусной кислоты степень электрической диссоциации α =0,0132, это означает, что 0,0132 (или 1,32%) общего количества растворённой уксусной кислоты продиссоциировало на ионы, а 0,9868 (или 98,68%) находится в виде недиссоциированных молекул. Диссоциация слабых электролитов в растворе подчиняется закону действия масс.

In general, a reversible chemical reaction can be represented as:

a A+ b B D d D+ e E

The reaction rate is directly proportional to the product of the concentration of reacting particles in powers of their stoichiometric coefficients. Then for the direct reaction

V 1 = k 1[A] a[B] b,

and the rate of the reverse reaction

V 2 = k 2[D] d[E] e.

At some point in time, the rates of the forward and reverse reactions will equalize, i.e.

This state is called chemical equilibrium. From here

k 1[A] a[B] b=k 2[D] d[E] e

Grouping the constants on one side and the variables on the other side, we get:

Thus, for a reversible chemical reaction in a state of equilibrium, the product of the equilibrium concentrations of the reaction products in powers of their stoichiometric coefficients, related to the same product for the starting substances, is a constant value at a given temperature and pressure. Numerical value of the chemical equilibrium constant To does not depend on the concentration of reactants. For example, the equilibrium constant for the dissociation of nitrous acid, in accordance with the law of mass action, can be written as:

HNO 2 + H 2 OD H 3 O + + NO 2 -

the value K a called the dissociation constant of the acid, in this case nitrous.

The dissociation constant of a weak base is expressed similarly. For example, for the ammonia dissociation reaction:

NH 3 + H 2 O DNH 4 + + OH -

the value K b called the dissociation constant of the base, in this case ammonia. The higher the dissociation constant of the electrolyte, the more the electrolyte dissociates and the higher the concentration of its ions in solution at equilibrium. There is a relationship between the degree of dissociation and the dissociation constant of a weak electrolyte:

This is a mathematical expression of the Ostwald dilution law: when a weak electrolyte is diluted, the degree of its dissociation increases. For weak electrolytes at To≤1∙10 -4 and With≥0.1 mol/l use the simplified expression:

To= α 2 ∙With or α

Example1. Calculate the degree of dissociation and concentration of ions and [ NH 4 + ] in 0.1 M ammonium hydroxide solution if To NH 4 OH \u003d 1.76 ∙ 10 -5

Given: NH 4 OH

To NH 4 OH \u003d 1.76 ∙ 10 -5

Decision:

Since the electrolyte is rather weak ( To NH 4 OH =1,76∙10 –5 <1∙ 10 - 4) и раствор его не слишком разбавлен, можно принять, что:

or 1.33%

The concentration of ions in a binary electrolyte solution is equal to C∙α, since the binary electrolyte ionizes with the formation of one cation and one anion, then \u003d [ NH 4 + ] \u003d 0.1 1.33 10 -2 \u003d 1.33 10 -3 (mol / l).

Answer:α=1.33%; \u003d [ NH 4 + ] \u003d 1.33 ∙ 10 -3 mol / l.

Theory of strong electrolytes

Strong electrolytes in solutions and melts completely dissociate into ions. However, experimental studies of the electrical conductivity of solutions of strong electrolytes show that its value is somewhat underestimated compared to the electrical conductivity that should be at 100% dissociation. This discrepancy is explained by the theory of strong electrolytes proposed by Debye and Hueckel. According to this theory, in solutions of strong electrolytes, there is an electrostatic interaction between ions. Around each ion, an “ionic atmosphere” is formed from ions of opposite charge, which slows down the movement of ions in solution when a direct electric current is passed. In addition to the electrostatic interaction of ions, in concentrated solutions it is necessary to take into account the association of ions. The influence of interionic forces creates the effect of incomplete dissociation of molecules, i.e. apparent degree of dissociation. The value of α determined experimentally is always somewhat lower than the true α. For example, in a 0.1 M Na 2 SO 4 solution, the experimental value α = 45%. To take into account electrostatic factors in solutions of strong electrolytes, the concept of activity is used (a). The activity of an ion is called the effective or apparent concentration, according to which the ion acts in solution. Activity and true concentration are related by the expression:

where f- activity coefficient, which characterizes the degree of deviation of the system from the ideal due to electrostatic interactions of ions.

The activity coefficients of ions depend on the value of µ, called the ionic strength of the solution. The ionic strength of a solution is a measure of the electrostatic interaction of all ions present in a solution and is equal to half the sum of the products of the concentrations (with) of each of the ions present in the solution per square of its charge number (z):

In dilute solutions (µ<0,1М) коэффициенты активности меньше единицы и уменьшаются с ростом ионной силы. Растворы с очень низкой ионной силой (µ < 1∙10 -4 М) можно считать идеальными. В бесконечно разбавленных растворах электролитов активность можно заменить истинной концентрацией. В идеальной системе a = c and the activity factor is 1. This means that there are practically no electrostatic interactions. In very concentrated solutions (µ>1M), the activity coefficients of ions can be greater than unity. The relationship of the activity coefficient with the ionic strength of the solution is expressed by the formulas:

at µ <10 -2

at 10 -2 ≤ µ ≤ 10 -1

+ 0,1z2µ at 0.1<µ <1

The equilibrium constant expressed in terms of activities is called thermodynamic. For example, for the reaction

a A+ b B d D+ e E

the thermodynamic constant has the form:

It depends on temperature, pressure and the nature of the solvent.

Since the activity of the particle , then

where To C is the concentration equilibrium constant.

Meaning To C depends not only on temperature, the nature of the solvent and pressure, but also on the ionic strength m. Since thermodynamic constants depend on the smallest number of factors, they are, therefore, the most fundamental characteristics of equilibrium. Therefore, in reference books, it is thermodynamic constants that are given. The values of thermodynamic constants of some weak electrolytes are given in the appendix of this manual. \u003d 0.024 mol / l.

With an increase in the charge of the ion, the activity coefficient and the activity of the ion decrease.

Questions for self-control:

What is an ideal system? Name the main reasons for the deviation of a real system from an ideal one.
What is the degree of dissociation of electrolytes?
Give examples of strong and weak electrolytes.
What is the relationship between the dissociation constant and the degree of dissociation of a weak electrolyte? Express it mathematically.
What is activity? How are the activity of an ion and its true concentration related?
What is an Activity Factor?
How does the charge of an ion affect the value of the activity coefficient?
What is the ionic strength of a solution, its mathematical expression?
Write down the formulas for calculating the activity coefficients of individual ions depending on the ionic strength of the solution.
Formulate the law of mass action and express it mathematically.
What is the thermodynamic equilibrium constant? What factors influence its value?
What is the concentration equilibrium constant? What factors influence its value?
How are thermodynamic and concentration equilibrium constants related?
To what extent can the value of the activity coefficient change?
What are the main provisions of the theory of strong electrolytes?

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.

Decision

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.

Decision

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

Decision

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)

Thus, 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 .

Decision

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 ,
hence, 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?

Decision

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 .

Decision

When mixed, the volume of the solution doubles and the concentration of each of the substances will decrease by half, i.e. will become 0.1 M or 1.0 10 -1 mol/l. These are there will be concentrationsPb 2+ andCl - . Hence,[ 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 the foregoing, it can be concluded 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 precipitateBaWithO 3 . The solution in contact with this precipitate is saturated withBaWithO 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 toBaWithO 3 and part of the sedimentBaWithO 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

The theory of electrolytic dissociation proposed by the Swedish scientist S. Arrhenius in 1887.

Electrolytic dissociation- this is the breakdown of electrolyte molecules with the formation of positively charged (cations) and negatively charged (anions) ions in solution.

For example, acetic acid dissociates like this in an aqueous solution:

CH 3 COOH⇄H + + CH 3 COO - .

Dissociation is a reversible process. But different electrolytes dissociate differently. The degree depends on the nature of the electrolyte, its concentration, the nature of the solvent, external conditions (temperature, pressure).

Degree of dissociation α - the ratio of the number of molecules decomposed into ions to the total number of molecules:

α=v´(x)/v(x).

The degree can vary from 0 to 1 (from the absence of dissociation to its complete completion). Indicated as a percentage. It is determined experimentally. During the dissociation of the electrolyte, the number of particles in the solution increases. The degree of dissociation indicates the strength of the electrolyte.

Distinguish strong and weak electrolytes.

Strong electrolytes- these are electrolytes, the degree of dissociation of which exceeds 30%.

Medium Strength Electrolytes- these are those whose degree of dissociation divides in the range from 3% to 30%.

Weak electrolytes- the degree of dissociation in an aqueous 0.1 M solution is less than 3%.

Examples of weak and strong electrolytes.

Strong electrolytes in dilute solutions completely decompose into ions, i.e. α = 1. But experiments show that dissociation cannot be equal to 1, it has an approximate value, but is not equal to 1. This is not a true dissociation, but an apparent one.

For example, let some connection α = 0.7. Those. according to the Arrhenius theory, 30% of non-dissociated molecules “float” in the solution. And 70% formed free ions. And the electrostatic theory gives a different definition to this concept: if α \u003d 0.7, then all molecules are dissociated into ions, but the ions are only 70% free, and the remaining 30% are bound by electrostatic interactions.

The apparent degree of dissociation.

The degree of dissociation depends not only on the nature of the solvent and solute, but also on the concentration of the solution and temperature.

The dissociation equation can be represented as follows:

AK ⇄ A- + K + .

And the degree of dissociation can be expressed as follows:

With an increase in the concentration of the solution, the degree of dissociation of the electrolyte decreases. Those. the degree value for a particular electrolyte is not a constant value.

Since dissociation is a reversible process, the reaction rate equations can be written as follows:

If dissociation is equilibrium, then the rates are equal and as a result we get equilibrium constant(dissociation constant):

K depends on the nature of the solvent and on the temperature, but does not depend on the concentration of the solutions. It can be seen from the equation that the more non-dissociated molecules, the lower the value of the electrolyte dissociation constant.

Polybasic acids dissociate in steps, and each step has its own value of the dissociation constant.

If a polybasic acid dissociates, then the first proton is most easily split off, and as the charge of the anion increases, the attraction increases, and therefore the proton is split off much more difficult. For example,

The dissociation constants of phosphoric acid at each stage should be very different:

I - stage:

II - stage:

III - stage:

At the first stage, phosphoric acid is an acid of medium strength, and at the 2nd stage it is weak, at the 3rd stage it is very weak.

Examples of equilibrium constants for some electrolyte solutions.

Consider an example:

If metallic copper is added to a solution containing silver ions, then at the moment of equilibrium, the concentration of copper ions should be greater than the concentration of silver.

But the constant has a low value:

AgCl⇄Ag + +Cl - .

Which suggests that by the time equilibrium was reached, very little silver chloride had dissolved.

The concentration of metallic copper and silver are introduced into the equilibrium constant.

Ionic product of water.

The table below contains data:

This constant is called ion product of water, which depends only on temperature. According to dissociation, there is one hydroxide ion for 1 H + ion. In pure water, the concentration of these ions is the same: [ H + ] = [Oh - ].

Hence, [ H + ] = [Oh- ] = = 10-7 mol/l.

If a foreign substance, such as hydrochloric acid, is added to water, the concentration of hydrogen ions will increase, but the ion product of water does not depend on the concentration.

And if you add alkali, then the concentration of ions will increase, and the amount of hydrogen will decrease.

Concentration and are interconnected: the more one value, the less the other.

The acidity of the solution (pH).

The acidity of solutions is usually expressed by the concentration of ions H + . In acidic environments pH<10 -7 моль/л, в нейтральных - pH\u003d 10 -7 mol / l, in alkaline - pH> 10 -7 mol/l.
The acidity of a solution is expressed in terms of the negative logarithm of the concentration of hydrogen ions, calling it pH.

pH = -lg[ H + ].

The relationship between the constant and the degree of dissociation.

Consider an example of the dissociation of acetic acid:

Let's find a constant:

Molar concentration С=1/V, we substitute into the equation and get:

These equations are by the breeding law of W. Ostwald, according to which the dissociation constant of the electrolyte does not depend on the dilution of the solution.

ELECTROLYTES Substances whose solutions or melts conduct electricity.

NON-ELECTROLYTES Substances whose solutions or melts do not conduct electricity.

Dissociation- decomposition of compounds into ions.

Degree of dissociation is the ratio of the number of molecules dissociated into ions to the total number of molecules in the solution.

STRONG ELECTROLYTES when dissolved in water, they almost completely dissociate into ions.

When writing the equations of dissociation of strong electrolytes put an equal sign.

Strong electrolytes include:

Soluble salts ( see solubility table);

Many inorganic acids: HNO 3, H 2 SO 4, HClO 3, HClO 4, HMnO 4, HCl, HBr, HI ( look acids-strong electrolytes in the solubility table);

Bases of alkali (LiOH, NaOH, KOH) and alkaline earth (Ca (OH) 2, Sr (OH) 2, Ba (OH) 2) metals ( see strong electrolyte bases in the solubility table).

WEAK ELECTROLYTES in aqueous solutions only partially (reversibly) dissociate into ions.

When writing the dissociation equations for weak electrolytes, the sign of reversibility is put.

Weak electrolytes include:

Almost all organic acids and water (H 2 O);

Some inorganic acids: H 2 S, H 3 PO 4, HClO 4, H 2 CO 3, HNO 2, H 2 SiO 3 ( look acids-weak electrolytes in the solubility table);

Insoluble metal hydroxides (Mg (OH) 2, Fe (OH) 2, Zn (OH) 2) ( see basescweak electrolytes in the solubility table).

The degree of electrolytic dissociation is influenced by a number of factors:

the nature of the solvent and electrolyte: strong electrolytes are substances with ionic and covalent strongly polar bonds; good ionizing ability, i.e. the ability to cause dissociation of substances, have solvents with a high dielectric constant, the molecules of which are polar (for example, water);

temperature: since dissociation is an endothermic process, an increase in temperature increases the value of α;

concentration: when the solution is diluted, the degree of dissociation increases, and with increasing concentration, it decreases;

stage of the dissociation process: each subsequent stage is less effective than the previous one, approximately 1000–10,000 times; for example, for phosphoric acid α 1 > α 2 > α 3:

H3PO4⇄Н++H2PO−4 (first stage, α 1),

H2PO−4⇄H++HPO2−4 (second stage, α 2),

НPO2−4⇄Н++PO3−4 (third stage, α 3).

For this reason, in a solution of this acid, the concentration of hydrogen ions is the highest, and the concentration of PO3−4 phosphate ions is the lowest.

1. Solubility and the degree of dissociation of a substance are not related to each other. For example, a weak electrolyte is acetic acid, which is highly (unrestrictedly) soluble in water.

2. A solution of a weak electrolyte contains less than others those ions that are formed at the last stage of electrolytic dissociation

The degree of electrolytic dissociation is also affected by addition of other electrolytes: e.g. degree of dissociation of formic acid

HCOOH ⇄ HCOO − + H+

decreases if a little sodium formate is added to the solution. This salt dissociates to form formate ions HCOO − :

HCOONa → HCOO − + Na +

As a result, the concentration of HCOO– ions in the solution increases, and according to the Le Chatelier principle, an increase in the concentration of formate ions shifts the equilibrium of the formic acid dissociation process to the left, i.e. the degree of dissociation decreases.

Ostwald dilution law- ratio expressing the dependence of the equivalent electrical conductivity of a dilute solution of a binary weak electrolyte on the concentration of the solution:

Here, is the dissociation constant of the electrolyte, is the concentration, and are the values of the equivalent electrical conductivity at concentration and at infinite dilution, respectively. The ratio is a consequence of the law of mass action and equality

where is the degree of dissociation.

The Ostwald dilution law was developed by W. Ostwald in 1888 and confirmed by him experimentally. The experimental establishment of the correctness of the Ostwald dilution law was of great importance for substantiating the theory of electrolytic dissociation.

Electrolytic dissociation of water. Hydrogen indicator pH Water is a weak amphoteric electrolyte: H2O H+ + OH- or, more precisely: 2H2O \u003d H3O + + OH- The dissociation constant of water at 25 ° C is: can be considered constant and equal to 55.55 mol / l (water density 1000 g / l, mass 1 l 1000 g, amount of water substance 1000g: 18g / mol \u003d 55.55 mol, C \u003d 55.55 mol: 1 l \u003d 55 .55 mol/l). Then This value is constant at a given temperature (25 ° C), it is called the ion product of water KW: The dissociation of water is an endothermic process, therefore, with an increase in temperature, in accordance with the Le Chatelier principle, dissociation increases, the ion product increases and reaches a value of 10-13 at 100 ° C. In pure water at 25°C, the concentrations of hydrogen and hydroxyl ions are equal to each other: = = 10-7 mol/l Solutions in which the concentrations of hydrogen and hydroxyl ions are equal to each other are called neutral. If acid is added to pure water, the concentration of hydrogen ions will increase and become more than 10-7 mol / l, the medium will become acidic, while the concentration of hydroxyl ions will instantly change so that the ion product of water retains its value of 10-14. The same thing will happen when alkali is added to pure water. The concentrations of hydrogen and hydroxyl ions are related to each other through the ion product, therefore, knowing the concentration of one of the ions, it is easy to calculate the concentration of the other. For example, if = 10-3 mol/l, then = KW/ = 10-14/10-3 = 10-11 mol/l, or if = 10-2 mol/l, then = KW/ = 10-14 /10-2 = 10-12 mol/l. Thus, the concentration of hydrogen or hydroxyl ions can serve as a quantitative characteristic of the acidity or alkalinity of the medium. In practice, it is not the concentrations of hydrogen or hydroxyl ions that are used, but the hydrogen pH or hydroxyl pOH indicators. The hydrogen index pH is equal to the negative decimal logarithm of the concentration of hydrogen ions: pH = - lg The hydroxyl index pOH is equal to the negative decimal logarithm of the concentration of hydroxyl ions: pOH = - lg It is easy to show by pronouncing the ionic product of water that pH + pOH = 14 the medium is neutral, if less than 7 - acidic, and the lower the pH, the higher the concentration of hydrogen ions. pH greater than 7 - alkaline environment, the higher the pH, the higher the concentration of hydroxyl ions.

Instruction

The essence of this theory is that when melted (dissolved in water), almost all electrolytes decompose into ions, which are both positively and negatively charged (which is called electrolytic dissociation). Under the influence of an electric current, negative (“-”) towards the anode (+), and positively charged (cations, “+”), move towards the cathode (-). Electrolytic dissociation is a reversible process (the reverse process is called "molarization").

The degree (a) of electrolytic dissociation is dependent on the electrolyte itself, the solvent, and their concentration. This is the ratio of the number of molecules (n) that have decayed into ions to the total number of molecules introduced into the solution (N). You get: a = n / N

Thus, strong electrolytes are substances that completely decompose into ions when dissolved in water. Strong electrolytes, as a rule, are substances with highly polar or bonds: these are salts that are highly soluble (HCl, HI, HBr, HClO4, HNO3, H2SO4), as well as strong bases (KOH, NaOH, RbOH, Ba (OH) 2 , CsOH, Sr(OH)2, LiOH, Ca(OH)2). In a strong electrolyte, the substance dissolved in it is mostly in the form of ions ( ); there are practically no molecules that are undissociated.

Weak electrolytes are substances that only partially dissociate into ions. Weak electrolytes, along with ions in solution, contain undissociated molecules. Weak electrolytes do not give a strong concentration of ions in solution.

The weak ones are:
- organic acids (almost all) (C2H5COOH, CH3COOH, etc.);
- some of the acids (H2S, H2CO3, etc.);
- almost all salts, slightly soluble in water, ammonium hydroxide, as well as all bases (Ca3 (PO4) 2; Cu (OH) 2; Al (OH) 3; NH4OH);
- water.

They practically do not conduct electric current, or conduct, but poorly.

note

Although pure water conducts electricity very poorly, it still has a measurable electrical conductivity, due to the fact that water dissociates slightly into hydroxide ions and hydrogen ions.

Helpful advice

Most electrolytes are corrosive substances, so when working with them, be extremely careful and follow safety regulations.

A strong base is an inorganic chemical compound formed by a hydroxyl group -OH and an alkali (group I elements of the periodic system: Li, K, Na, RB, Cs) or alkaline earth metal (group II elements Ba, Ca). They are written as formulas LiOH, KOH, NaOH, RbOH, CsOH, Ca(OH) ₂, Ba(OH) ₂.

You will need

evaporating cup
burner
indicators
metal rod
H₃RO₄

Instruction

Strong bases exhibit, characteristic of all. The presence in the solution is determined by the change in color of the indicator. Add phenolphthalein to the sample with the test solution or omit litmus paper. Methyl orange is yellow, phenolphthalein is purple, and litmus paper is blue. The stronger the base, the more intense the color of the indicator.

If you need to find out which alkalis are presented to you, then conduct a qualitative analysis of the solutions. The most common strong bases are lithium, potassium, sodium, barium, and calcium. Bases react with acids (neutralization reactions) to form salt and water. In this case, Ca(OH) ₂, Ba(OH) ₂ and LiOH can be distinguished. When with acid, insoluble ones are formed. The remaining hydroxides will not give precipitation, tk. all K and Na salts are soluble.
3 Ca(OH) ₂ + 2 H₃RO₄ --→ Ca₃(PO₄)₂↓+ 6 H₂О

3 Va(OH) ₂ +2 H₃RO₄ --→ Va₃(PO₄)₂↓+ 6 H₂О

3 LiOH + Н₃РО₄ --→ Li₃РО₄↓ + 3 H₂О
Strain them and dry them. Inject the dried sediments into the flame of the burner. Lithium, calcium and barium ions can be qualitatively determined by changing the color of the flame. Accordingly, you will determine where which hydroxide is. Lithium salts color the burner flame carmine red. Barium salts - in green, and calcium salts - in raspberry.

The remaining alkalis form soluble orthophosphates.

3 NaOH + Н₃РО₄--→ Na₃РО₄ + 3 H₂О

3 KOH + H₃PO₄--→ K₃PO₄ + 3 H₂О

Evaporate the water to a dry residue. Evaporated salts on a metal rod alternately bring into the burner flame. There, sodium salt - the flame will turn bright yellow, and potassium - pink-purple. Thus, having a minimum set of equipment and reagents, you have determined all the strong reasons given to you.

An electrolyte is a substance that in the solid state is a dielectric, that is, does not conduct electric current, however, in a dissolved or molten form it becomes a conductor. Why is there such a drastic change in properties? The fact is that electrolyte molecules in solutions or melts dissociate into positively charged and negatively charged ions, due to which these substances in such a state of aggregation are able to conduct electric current. Most salts, acids, bases have electrolytic properties.

Instruction

What substances are strong? Such substances, in solutions or melts of which almost 100% of the molecules are exposed, and regardless of the concentration of the solution. The list includes the absolute majority of soluble alkalis, salts and some acids, such as hydrochloric, bromine, iodine, nitric, etc.

And how do the weak ones behave in solutions or melts? electrolytes? Firstly, they dissociate to a very small extent (no more than 3% of the total number of molecules), and secondly, they go the worse and slower, the higher the concentration of the solution. Such electrolytes include, for example, (ammonium hydroxide), most organic and inorganic acids (including hydrofluoric - HF) and, of course, the familiar water to all of us. Since only a negligible fraction of its molecules decomposes into hydrogen ions and hydroxyl ions.

Remember that the degree of dissociation and, accordingly, the strength of the electrolyte are dependent on factors: the nature of the electrolyte itself, the solvent, and temperature. Therefore, this division itself is to a certain extent conditional. After all, the same substance can, under different conditions, be both a strong electrolyte and a weak one. To assess the strength of the electrolyte, a special value was introduced - the dissociation constant, determined on the basis of the law of mass action. But it is applicable only to weak electrolytes; strong electrolytes they do not obey the law of the acting masses.

Sources:

strong electrolytes list

salt- these are chemicals consisting of a cation, that is, a positively charged ion, a metal and a negatively charged anion - an acid residue. There are many types of salts: normal, acidic, basic, double, mixed, hydrated, complex. It depends on the compositions of the cation and anion. How can you determine base salt?