Activity and activity coefficient of electrolytes. Average ionic activity and average ionic activity coefficient

In connection with the electrostatic interaction in solution, even for dilute solutions of strong electrolytes, concentrations in thermodynamic equations must be replaced by activities. For example, if for an ion, as well as for a component in a solution, the expression is true:

where with i– concentration i ion in an ideal solution, then for a real solution we will have:

where a i = c i f i - activity of the i-th ion in solution,

f i - activity coefficient.

Then the interaction energy of an ion with surrounding ions per 1 mole of ions is equal to

f i →1 at с→0

Τᴀᴋᴎᴍ ᴏϬᴩᴀᴈᴏᴍ, the value of the activity coefficient, mainly depending on the strength of the electrostatic interaction of ions, as well as a number of other effects, characterizes the degree of deviation of the properties of real electrolyte solutions from ideal solutions. According to the meaning of f i, this is the work of transferring an ion from an ideal solution to a real one.

Distinguish between the activity of the electrolyte and the activity of the ions. For any electrolyte, the dissociation process can be written as follows:

where n + and n - - the number of ions BUT with charge z+ and ions B with charge z– into which the original particle decays.

For the electrolyte solution as a whole, we can write:

m salt = m 0 salt + RT ln a salt, (9)

On the other hand, the chemical potential of the electrolyte is the sum of the chemical potentials of the ions, since the electrolyte decomposes into ions:

m salt = n + m + + n - m - , (10)

m + and m - refer to one mole of ions, m salts - to one mole of electrolyte. Let us substitute expression (10) into (9):

n + m + + n - m - = m 0 salt + RT ln a salt (11)

For each type of ions, we can write an equation like (9):

m + = m 0 + + RT ln a +

m - = m 0 - + RT ln a - (12)

We substitute equation (12) into the left side of equation (11) and swap the right and left sides.

m 0 s + RT ln aс = n + m 0 + + n + RT ln a+ + n – m 0 – + n - RT ln a - (13)

Combine all the terms with m 0 on the left side

(m 0 s - n + m 0 + - n - m 0 -) = n + RT ln a+ + n - RT·ln a- - RT·ln a salt (14)

If we take into account that by analogy with formula (10)

m 0 С \u003d n + m 0 + + n - m 0 - (15)

then m 0 С - n + m 0 + - n - m 0 - = 0 (16)

Equation (15) is similar to equation (10), but it refers to the standard state when ( a C = a + = a- = 1). In equation (14), the right side is equal to zero, and it will be rewritten as follows:

RT ln a c = n + RT ln a+ + n - RT·ln a -

ln a c = ln a+ n + + ln a+n-

This is the relationship of the activity of the electrolyte in solution with the activities of the ions

where aС - electrolyte activity͵ a+ and a– - activities of positive and negative ions. For example, for binary electrolytes AB, the following is true:

Hence

It is impossible in principle to find experimentally the activities of individual ions, since one would have to deal with a solution of one kind of ions. It's impossible. For this reason, the concept of average ionic activity (), which is the geometric mean of the activity of individual ions, was introduced:

or substituting expression (17) we have:

The cryoscopic method and the method based on the determination of vapor pressure make it possible to determine the activity of the electrolyte as a whole ( a C) and using equation (19) find the average ionic activity.

In all cases, when it becomes extremely important to substitute the value a+ or a- into some equation, these values ​​are replaced by the average activity of a given electrolyte a± , for example,

a ± ʼʼ a + ʼʼ a –

As is known, activity is related to concentration by the ratio a=f∙m. The average ionic activity coefficient () is determined by an expression similar to the expression for the average ionic activity

There are tabular values ​​for different ways of expressing concentrations (molality, molarity, mole fractions). It is worth saying that for them it has numerically different values. Experimentally, the values ​​are determined by the cryoscopic method, the method of measuring vapor pressure, the method of measuring the EMF of galvanic cells, etc.

Similarly, the average ionic stoichiometric coefficient n ± is determined from the expression:

b) Concentrations of ions formed during complete dissociation Na2CO3, are equal

Since n + = 2, n – = 1, then .

Activity and activity coefficient of electrolytes. Average ionic activity and average ionic activity coefficient - concept and types. Classification and features of the category "Activity and activity coefficient of electrolytes. Average ionic activity and average ionic activity coefficient" 2017, 2018.

Average ionic activity, activity coefficient, concentration.

The total concentration of ions in a solution is the molar concentration of the dissolved electrolyte, taking into account its degree of dissociation into ions and the number of ions into which the electric stove molecule dissociates in solution.

For strong electrolytes, α = 1; therefore, the total concentration of ions is determined by the molar concentration of the electrolyte and the number of ions into which a strong electrolyte molecule decomposes in solution.

So, in the case of dissociation of a strong electrolyte - sodium chloride in an aqueous solution

NaCl → Na + + Cl -

at initial electrolyte concentration with(NaCl) \u003d 0.1 mol / l, the ion concentrations turn out to be equal to the same value: c (Na +) \u003d 0.1 mol / l and c (Cl -) \u003d 0.1 mol / l.

For a strong electrolyte of a more complex composition, for example, aluminum sulfate Al 2 (SO 4) 3, the concentrations of the cation and anion are also easily calculated, taking into account the stoichiometry of the dissociation process:

Al 2 (SO 4) 3 → 2 Al 3+ + 3 SO 4 2-

If the initial concentration of aluminum sulfate from ref\u003d 0.1 mol / l, then c (A1 3+) \u003d 2 0.1 \u003d 0.2 mol / l and with( SO 4 2-) \u003d 3 0.1 \u003d \u003d 0.3 mol / l.

Activity a related to total concentration with formal relation

where f ˗ activity factor.

At with→ 0 value a → c, so f→1, i.e. for extremely diluted solutions, the activity coincides in numerical value with the concentration, and the activity coefficient is equal to one.

Lewis and Randall introduced some mathematical corrections to the ratios proposed by Arrhenius.

G. Lewis and M. Randall proposed a method of using activities instead of concentrations, which made it possible to formally take into account the whole variety of interactions in solutions without taking into account their physical nature.

In electrolyte solutions, both cations and anions of the solute are simultaneously present. It is physically impossible to introduce only one kind of ions into the solution. Even if such a process were feasible, it would cause a significant increase in the energy of the solution due to the introduced electric charge.

The relationship between the activities of individual ions and the activity of the electrolyte as a whole is established based on the condition of electrical neutrality. For this, the concepts average ionic activity and average ionic activity coefficient.

If an electrolyte molecule dissociates into n + cations and n - anions, then the average ionic activity of the electrolyte a ± is:

,

where and is the activity of cations and anions, respectively, n is the total number of ions (n=n + + n -).

Similarly, the average ionic activity coefficient of the electrolyte is written, which characterizes the deviations of the real solution from the ideal

.

Activity can be represented as the product of concentration and activity coefficient. There are three scales for expressing activities and concentrations: molality (molal or practical scale), molarity with(molar scale) and mole fraction X(rational scale).

In the thermodynamics of electrolyte solutions, the molar concentration scale is commonly used.

DEBYE-HUKKEL ELECTROLYTES.

One of the theories that quantitatively takes into account ion-ion interactions is Debye-Hückel theory, which explains quite well the properties of dilute solutions of strong electrolytes. The degree of dissociation for strong electrolytes is equal to one. Therefore, the dependence of electrical conductivity, osmotic pressure, and other properties of solutions on concentration is determined mainly by the action interionic forces and solvation effects. Solvation is understood as a set of energy and structural changes that occur in a solution during the interaction of solute particles with solvent molecules.

The Debye-Hückel theory is based on the following provisions: the electrostatic interaction of oppositely charged ions leads to the fact that around positive ions the probability of finding negative ions will be greater than positive ones. Thus, around each ion, as it were, there is an ionic atmosphere of oppositely charged ions. (The sphere in which the charge opposite in sign to the central ion predominates is called ionic atmosphere). The ionic atmosphere around the ion contains both positive and negative ions, however, on average, there is an excess of negative ions around each positive ion, and an excess of positive ions around a negative ion. The solution as a whole remains electrically neutral.

Chemical potential i th component in an ideal solution is equal to:

where with i– concentration i th ion in solution. For a real solution:

where a i = c i · f i- activity of the i-th ion in solution, f i– activity coefficient. Then the interaction energy of the central ion with the ionic atmosphere per 1 mole of ions is equal to

Thus, the value of the activity coefficient, which depends on the strength of the electrostatic interaction of ions, the degree of their solvation, and a number of other effects, characterizes the degree of deviation of the properties of real electrolyte solutions from the laws of ideal solutions.

1.3. Activity and activity coefficient of electrolytes.

AVERAGE IONIC ACTIVITY AND AVERAGE IONIC COEFFICIENT

ACTIVITIES. IONIC POWER. IONIC STRENGTH RULE.

Distinguish electrolyte activity and ion activity. For any electrolyte, the dissociation process can be written as follows:

where  + and  - - the number of ions BUT with charge z+ and ions B with charge z– into which the original particle decays. For example, during the dissociation of barium chloride:

.

The relationship between the activity of the electrolyte and the activities of the ions is expressed by the following relationship:

, (1.11)

where a- electrolyte activity, a+ and a– - activities of positive and negative ions. For example, for binary electrolytes it is true:

.

Experimental methods for determining the activity of individual ions ( a+ and a-) does not exist. Therefore, the concept was introduced average ionic activity(), which is the geometric mean of the activity of individual ions:

, (1.12)

where
.

The cryoscopic method and the method based on the determination of vapor pressure make it possible to determine the activity of the electrolyte as a whole ( a) and using equation (7.13) find the average ionic activity.

Average ionic activity coefficient() is determined by the expression

. (1.14)

Values mainly determined by the cryoscopic method and the EDS method.

Average ionic molality(
) is defined as

. (1.15)

If the concentration of a solution is expressed in terms of molality, then

Example 1.1. Find the relationship between the activity of the electrolyte, its molar concentration and the average ionic activity coefficient for solutions NaCl and Na 2 CO 3 molality m.

a) Concentrations of ions formed during complete dissociation NaCl, are equal m:

.

Since  + =  – = 1, then

.

For equal-valent electrolytes, the average molality will be equal to the total molality of the electrolyte:

,

b) Concentrations of ions formed during complete dissociation Na 2 CO 3 , are equal

.

Since  + = 2,  – = 1, then

.

With

the average ionic activity coefficient depends on the concentration of the solution (Fig. 1). In the region of extremely dilute electrolyte solutions, this dependence is linear in the coordinates
.

Rice. Fig. 1. Dependence of the average ionic 2. Dependence of the average coefficient

activity coefficient of ion activity on the ionic strength of the solution.

on electrolyte concentration. Curve 1 describes the experimental

dependence, curve 2 describes dependence

according to the Debye-Hückel limit law.

The presence of other salts in the solution changes the activity coefficient of the given salt and the stronger, the greater the charge of the added ions. The total concentration of all ions in a solution is expressed through ionic strength of the solution , defined as half the sum of the products of the molality of all ions and the square of their charges :

, (1.16)

where m i– concentration i-th ion; z i- charge i-th ion.

The dependence of the average activity coefficient of ions on the ionic strength of the solution has a complex character and is shown in fig. 2.

Example 1.2. Determine the ionic strength of a solution containing 0.01 mol per 1000 g of water
and 0.1 mol
.

Decision. The ionic strength of such a solution is

Example 1.3. Determine the ionic strength of the solution
with molality m = 0,5.

Decision. By equation (7.16) we get

For solutions of strong electrolytes, ionic strength rule : in solutions with the same ionic strength, the average activity coefficients of the ions are equal. The theory of strong electrolytes leads to the following relationship relating the average activity coefficients of ions to the ionic strength of the solution in the region of highly dilute electrolytes:

, (1.17)

where A = f (D, T) is a constant depending on the permittivity of the solvent ( D) and temperature ( T).

Equation (1.17) is applicable only for very large dilutions ( I≤ 0.01, fig. 2) why it got the name limiting Debye-Hückel law. In weakly mineralized waters for calculation at 25°C the following equation is used:

. (1.18)

For aqueous solutions of binary electrolytes at 25 o C, the following is true:

. (1.19)

It is known that in highly dilute solutions of electrolytes, the activity coefficients of ions mainly take into account corrections to their concentrations (molalities) due to electrostatic (ion–ion) interaction. At the same time, according to the Coulomb law, these interactions also depend on the magnitude of the charges and radii of the ions. Therefore, it is natural to accept, as was first done by D. McInnes, that the activity coefficients of ions with the same charges and radii in solutions with the same ionic strength will be the same. This assumption has been called McInnes rule.

McInnes suggested taking potassium and chlorine ions as standards, as having the same charges and radii of hydrated ions. Having defined the values
and
, one can then calculate the activity coefficients of all other ions based on the law of ionic strength.

SUBJECT2

Specific and equivalent electrical conductivity, their dependence on concentration for strong and weak electrolytes. Ion mobility. Kohlrausch's law of independence of movement of ions, limiting ionic electrical conductivity. Abnormal mobility of hydroxyl and hydroxonium ions. Experimental applications of the electrical conductivity method.

2.1. SPECIFIC ELECTRICAL CONDUCTIVITY OF ELECTROLYTE SOLUTIONS.

When an electric field is applied to an electrolyte solution, the solvated ions, which were previously in random thermal motion, begin an ordered movement (migration) to oppositely charged electrodes. With an increase in the speed of movement of ions, the resistance of the medium increases and after a while the speed of movement of the ions becomes constant.

Ion movement speed i-th type is determined by the gradient of the potential (strength) of the electric field E(V / cm) and the resistance of the medium, depending on the temperature, the nature of the ion and the solvent:

, (2.1)

where U(B) - potential difference between the electrodes, l(cm) is the distance between them, u i(cm 2 V -1 s -1) - the absolute speed of movement of ions under these conditions (i.e., the speed of movement of ions at E= 1 V/cm).

A measure of the ability of a substance to conduct an electric current when an external electric field is applied is electrical conductivity (electrical conductivity)L. In practice, this ability is often characterized by the reciprocal - conductor resistance. So, the total resistance of the conductor R(ohm) length l(cm) and cross section S(cm 2) equals

, (2.2)

where ρ is the coefficient of proportionality, called resistivity. From (8.2) it follows that the resistivity is the resistance of a conductor 1 cm long and 1 cm 2 in cross section, its dimension is:

. (2.2)

Electrical conductivity electrolyte æ - the reciprocal of the resistivity:

æ
[Ohm -1 cm -1]. (2.3)

It characterizes the electrical conductivity of an electrolyte layer 1 cm thick with a cross-sectional area of ​​1 cm 2 . Then

æ . (2.4)

The electrical conductivity of an electrolyte solution is determined by the number of ions that carry electricity and the rate of their migration.

Let between electrodes located at a distance l(cm) and to which the potential difference is applied U(B), there is an electrolyte solution (Fig. 3). For ions i-th type: concentration C i(mol-eq / cm 3) and migration rate υ i(cm/s).

Rice. 3. Scheme of charge transfer through the electrolyte solution.

H
through the cross section S solution (Fig. 3) migrates in 1 s ( C i υ i S) mole equivalents of ions i-th species that will transfer (
) to the amount of electricity where F– Faraday number(96485 C/mol-eq). The amount of electricity (C) transferred by all ions in 1 s (i.e., the current strength I in A) is equal to:

(2.5)

Or, taking into account (8.1),

. (2.6)

Ohm's law

(æ S), (2.7)

æ. (2.8)

Then, from equations (8.6) and (8.8), for the electrical conductivity we obtain

æ
. (2.9),

i.e., the specific electrical conductivity of the electrolyte is proportional to the concentrations of ions and their absolute velocities. For a binary electrolyte solution of concentration With(mol-equiv / cm 3) with the degree of dissociation α we have

æ
, (2.10)

where u+ and u‑ ‑ absolute velocities of cations and anions.

With an increase in the temperature of the electrolyte, the velocities of the movement of ions and the electrical conductivity increase:

æ 2 = æ 1
, (2.11)

where B- temperature coefficient (for strong acids 0.016; for strong bases 0.019; for salts 0.022).

2.2. EQUIVALENT ELECTRICAL CONDUCTIVITY.

Specific conductivity of solutions depends on the nature of the electrolyte, the nature of the solvent, temperature, concentration of ions in the solution, etc. Although the electrical conductivity is an inconvenient quantity for understanding the properties of electrolytes, it can be measured directly and then converted to equivalent electrical conductivity λ. The equivalent electrical conductivity is the electrical conductivity of such a volume of solution V (cm 3 ), which contains 1 mole equivalent of a solute and is enclosed between two parallel electrodes of the corresponding area, located at a distance of 1 cm from each other:

æ V = æ / C, (2.12)

where With- concentration of the solution (mol-equiv / cm 3).

Equivalent electrical conductivity (Ohm -1 cm 2 (mol-equiv) -1) is easy to calculate if the specific electrical conductivity and concentration of the solution are known.

The following equation is used to describe the temperature dependence of the equivalent electrical conductivity:

, (2.13)

where  and  are empirical coefficients. The increase in electrical conductivity with increasing temperature is mainly due to a decrease in the viscosity of the electrolyte solution. Typically, with an increase in temperature by 1 K, the electrical conductivity increases by 1.5 - 2%.

The equivalent electrical conductivity of electrolyte solutions increases with dilution and in the range of limiting dilutions reaches the limiting value λ ∞ , called electrical conductivity at infinite dilution or ultimate electrical conductivity. This value corresponds to the electrical conductivity of a hypothetically infinitely dilute solution characterized by complete dissociation of the electrolyte and the absence of electrostatic interaction forces between ions.

Equations (2.10) and (2.11) imply that

The product of the Faraday number and the absolute speed of the ion is called mobility and she:

. (2.15)

where λ + and λ - are the cation and anion mobilities, respectively. Ion mobilities are measured in the same units as the equivalent electrical conductivity (cm 2 Ohm -1 mol-eq -1), so they are sometimes called ionic conductivities or electrical conductivities of ions.

With an infinite dilution (α = 1), we obtain

, (8.17)

where
and
- limiting mobility of ions.

The value of the limiting electrical conductivity of an infinitely dilute electrolyte solution is the sum of two independent terms, each of which corresponds to a certain type of ion. This relation was established by Kohlrausch and is called the law of independent motion of ions (Kohlrausch law): the equivalent electrical conductivity at infinite dilution is equal to the sum of the limiting ion mobilities. The essence of this law is as follows: In an extremely dilute electrolyte solution, cations and anions carry current independently of each other.

Kohlrausch's law helped to calculate the values ​​of λ ∞ for many weak electrolytes, for which it was impossible to determine these values ​​from experimental data by extrapolating them to zero concentration (or to infinite dilution) as is done in the case of strong (and average) electrolytes. The limiting ion mobilities, as well as the equivalent electrical conductivity, increase with temperature. Their values, for example, at 25 ° C lie in the range from 30 to 80 and from 40 to 80 (cm 2 Ohm -1 mol-eq -1) for singly charged cations and anions, respectively.

Ions IS HE- and H+ Abnormally high mobility is observed:

198 and
350 (cm 2  Ohm -1 mol-equiv -1) at 25 o C,

which is explained by a special - relay - mechanism of their movement (Fig. 4).

R
is. 4. Relay-race mechanism for ion movement IS HE- and H + .

Based on the equivalent electrical conductivity of the electrolyte solution and the limiting ion mobilities, the degree of dissociation of a weak electrolyte can be calculated:

, (2.18).

For strong electrolytes that dissociate completely, calculate to conductivity factor:

, (2.19)

which takes into account the influence of the electrostatic interaction of ions on the speed of their movement.

Taking into account the new concept - ion mobility - for electrical conductivity, we can write:

æ
, (2.20)

Note that modern scientific and educational literature also uses the concept molar electrical conductivity λ m, which is easily related to the value of λ, knowing the number of mole equivalents ( Z) in 1 mole of a substance:

. (2.22)

2.2. DEPENDENCE OF SPECIFIC AND EQUIVALENT ELECTRICAL CONDUCTIVITIES ON CONCENTRATION

FOR WEAK AND STRONG ELECTROLYTES.

E
equivalent electrical conductivity weak and strong electrolytes increases with dilution (Fig. 5 b). For weak electrolytes, this is mainly due to the fact that with increasing dilution, the degree of electrolyte dissociation increases and tends to 1 in the limit. The increase in the equivalent electrical conductivity of strong electrolytes is mainly due to a change in the ion mobilities. The mobility of ions is the less, the greater the concentration of the solution. In the region of highly dilute solutions, the ion mobilities reach their limiting value.

Rice. 5. Dependence of specific ( a) and equivalent ( b)

conductivity on the concentration of the electrolyte solution.

Electrical conductivity for strong electrolytes, the higher the concentration of ions and the greater their absolute velocities (mobilities). Acids have the highest electrical conductivity, then bases, followed by salts, the electrical conductivity of solutions of such weak electrolytes as acetic acid or ammonia is very low.

The curves of dependences of specific electrical conductivity on concentration have maxima (Fig. 5 a). In dilute solutions of weak and strong electrolytes, the increase in electrical conductivity with concentration is due to an increase in the number of ions that carry electricity. A further increase in concentration is accompanied by an increase in the viscosity of the solution, which reduces the ion velocity and electrical conductivity. Moreover, for weak electrolytes in concentrated solutions, the degree of dissociation and, consequently, the total number of ions decreases markedly. For weak electrolytes, the speed of ion movement is almost independent of concentration, and in the general case, their electrical conductivity varies slightly with concentration.

For strong electrolytes in the region of dilute solutions, there are practically no interionic interactions, but the number of ions is small - the electrical conductivity is low. With increasing concentration, the number of ions per unit volume increases, which leads to an increase in electrical conductivity. However, further increasing interaction between ions leads to a decrease in the mobility of ions and the increase in electrical conductivity slows down. Finally, the interaction between ions with increasing concentration begins to increase so strongly that it leads to a decrease in electrical conductivity.

From the standpoint of the Debye-Hückel theory, the decrease in the mobility of ions with increasing concentration is due to the effects of deceleration of the movement of ions due to the electrostatic interaction between the ion and the ionic atmosphere.

The effect of electrophoretic inhibition due to the deceleration of the movement of the central ion by the oncoming movement of the ionic atmosphere and has a hydrodynamic nature. Since the ions are hydrated, the movement of the central ion occurs not in a stationary medium, but in a medium moving towards it. A moving ion is under the influence of an additional retarding force (electrophoretic braking force), which leads to a decrease in the speed of its movement.

The effect of relaxation inhibition. The ionic atmosphere has spherical symmetry as long as there is no external electric field. As soon as the central ion begins to move under the action of an electric field, the symmetry of the ionic atmosphere is broken. The movement of the ion is accompanied by the destruction of the ionic atmosphere in the new position of the ion and its formation in another, new one. This process takes place at a finite rate over a period of time called relaxation time. As a result, the ionic atmosphere loses its central symmetry and behind the moving ion there will always be some excess charge of the opposite sign, which causes a decrease in its speed.

The density of the ionic atmosphere increases with an increase in the electrolyte concentration, which leads to an increase in the braking effects. The theory of electrophoretic and relaxation effects was developed by L. Onsager. It quantitatively allows taking into account the influence of these effects on the value of the equivalent electrical conductivity of the electrolyte solution:

where constants ( AT 1 λ∞) and AT 2 characterize the influence of relaxation and electrophoretic effects, respectively. In solutions with With→ 0, these effects practically do not appear and
.

2.4. EXPERIMENTAL APPLICATIONS OF THE ELECTRICAL CONDUCTIVITY METHOD.

2.4.1. Determination of the dissociation constant and degree of dissociation

weak electrolytes.

The degree of dissociation  of a weak electrolyte can be found from relation (8.18):

.

Dissociation constant To D weak electrolyte is related to the degree of dissociation  by the equation

. (2.24)

Taking into account (8.18), we obtain

. (2.25)

The value λ ∞ is calculated according to the Kohlrausch law (Equation 2.17).

2.4.2. Determination of the solubility product

insoluble compounds.

Electrolyte solubility (S) is its concentration in a saturated solution (mol/l), and solubility product (ETC) is the product of the activities of the cation and anion of a sparingly soluble salt.

A saturated solution of a sparingly soluble salt is a very dilute solution (α → 1 and λ → λ ∞). Then

(æ 1000) / C. (2.26)

By finding the value of λ ∞ from tabular data and measuring the electrical conductivity of the solution, we can calculate the concentration of the saturated solution (in mol-eq/l), which is the solubility of the salt

C= (æ 1000) / λ∞ = S (2.27).

Since æ of sparingly soluble solutions (æ R) is often commensurate with the electrical conductivity of water (æ B), then in the equations the specific electrical conductivity of the solution is often calculated as the difference: æ = æ R - æ B.

For sparingly soluble salts, the activities of the cation and anion practically coincide with their concentrations, therefore

ETC =
(2.28),

where  i is the stoichiometric coefficient of the ion in the dissociation equation; n is the number of types of ions into which the electrolyte dissociates; C i is the ion concentration related to the electrolyte concentration With ratio

.

Since  = 1, then

,

and the solubility product

. (2.29)

So, for a poorly soluble (binary) monovalent electrolyte dissociating according to the scheme

,

(mol/l) 2 .

THEME 3

Electrode processes. The concept of electromotive forces (EMF) and potential jumps. Electrochemical circuits, galvanic cells. Normal hydrogen electrode, standard electrode potential. Thermodynamics of a galvanic cell. Classification of electrochemical circuits and electrodes.

The science that explains chemical phenomena and establishes their patterns based on the general principles of physics. The name of the science Physical Chemistry was introduced by M.V. Lomonosov, who for the first time (1752 1753) formulated its subject and tasks and established one ... ... Big Encyclopedic Dictionary

PHYSICAL CHEMISTRY- PHYSICAL CHEMISTRY, “a science that explains, on the basis of provisions and experiments, the physical cause of what happens through chem. operations in complex bodies. This definition was given to her by the first physicochemist M.V. Lomonosov in a course read by ... Big Medical Encyclopedia

PHYSICAL CHEMISTRY, the science that studies the physical changes associated with CHEMICAL REACTIONS, as well as the relationship between physical properties and chemical composition. The main sections of physical chemistry THERMODYNAMICS, dealing with changes in energy in ... ... Scientific and technical encyclopedic dictionary

Physical chemistry- - a branch of chemistry in which the chemical properties of substances are studied on the basis of the physical properties of their constituent atoms and molecules. Modern physical chemistry is a broad interdisciplinary field bordering on various branches of physics… Encyclopedia of terms, definitions and explanations of building materials

PHYSICAL CHEMISTRY, explains chemical phenomena and establishes their laws on the basis of the general principles of physics. Includes chemical thermodynamics, chemical kinetics, the doctrine of catalysis, etc. The term physical chemistry was introduced by M.V. Lomonosov in 1753 ... Modern Encyclopedia

Physical chemistry- PHYSICAL CHEMISTRY, explains chemical phenomena and establishes their patterns based on the general principles of physics. Includes chemical thermodynamics, chemical kinetics, the doctrine of catalysis, etc. The term “physical chemistry” was introduced by M.V. Lomonosov in ... ... Illustrated Encyclopedic Dictionary

PHYSICAL CHEMISTRY- section of chem. science, studying chemistry. phenomena based on the principles of physics (see (1)) and physical. experimental methods. F. x. (like chemistry) includes the doctrine of the structure of matter, chem. thermodynamics and chemistry. kinetics, electrochemistry and colloidal chemistry, teaching ... ... Great Polytechnic Encyclopedia

Exist., number of synonyms: 1 physical (1) Dictionary of ASIS synonyms. V.N. Trishin. 2013 ... Synonym dictionary

physical chemistry- — EN physical chemistry A science dealing with the effects of physical phenomena on chemical properties. (Source: LEE) … … Technical Translator's Handbook

physical chemistry- - a science that explains chemical phenomena and establishes their patterns based on physical principles. Dictionary of Analytical Chemistry ... Chemical terms

Books

• Physical Chemistry, A. V. Artemov. The textbook was created in accordance with the Federal State Educational Standard in the areas of training of bachelors, providing for the study of the discipline `Physical Chemistry`.…
• Physical Chemistry, Yu. Ya. Kharitonov. The textbook outlines the basics of physical chemistry in accordance with the approximate program for the discipline "Physical and colloidal chemistry" for the specialty 060301 "Pharmacy". The publication is intended…