Concentration is a value characterizing the quantitative composition of a solution.

The concentration of a solute is the ratio of the amount of a solute or its mass to the volume of a solution (mol/l, g/l), that is, it is a ratio of heterogeneous quantities.

Those quantities that are the ratio of similar quantities (the ratio of the mass of a dissolved substance to the mass of a solution, the ratio of the volume of a dissolved substance to the volume of a solution) are correctly called fractions. However, in practice, for both types of expression of composition, the term concentration is used and they speak of the concentration of solutions.

A solution is a homogeneous system consisting of two or more substances, one of which is a solvent and the other a solute. The solution can be saturated, i.e. contain a limiting amount of dissolved substance and be in a state of mobile equilibrium.

Mass fractions – the ratio of the mass of the dissolved substance to the mass of the solution (if in percentage, then ∙ 100%).

Molar concentration – С М – number of moles of dissolved substance in 1 liter of solution.

where V is the volume (if the volume is not indicated in the problem, then it is meant that it is equal to 1 liter), M is the molar mass.

Normality (normal concentration) - C n - the number of equivalents of a dissolved substance contained in 1 liter of solution.

C n = , where 1 eq. – substance equivalent (m e)

The equivalent of a substance is the amount of it that combines with 1 mole of a hydrogen atom or replaces such an amount in chemical reactions; a certain number of grams of a substance, numerically equal to its equivalent.

Equivalent mass = mass of one equivalent.

The equivalent is calculated:

a) the equivalent of an acid is equal to its molar mass divided by the basicity (number of hydrogen ions) of the acid.

b) the equivalent of a base is equal to its molar mass divided by the acidity (number of hydroxyl groups) of the base.

c) the equivalent of a salt is equal to its molar mass divided by the sum of the charges of the cations or anions that form it.

Law of equivalents: all substances interact with each other in equivalent quantities.

For substances;

C m1 ∙ V 1 = C n 2 ∙ V 2 for solutions;

Titer is the mass of a substance in 1 liter of solution.

Title = =

Let us also recall the formula:

m solution = ρ ∙ V, where ρ is the density of the substance.

ρ(solution)=1.33g/ml

() =49%, or 0.49

Find: C() Solution:

1. To move from mass fraction to molar concentration, you need to calculate what mass 1000 ml of solution has:

2. Calculate the mass in this solution:

3. Let’s find how many moles are contained in 651.7g:

4. Find the molar concentration of orthophosphoric acid in solution:

5. Let’s find the equivalent concentration of orthophosphoric acid in solution:

According to the formula:

≈ 20 mol/l Answer: () = 6.65 mol/l

() ≈ 20mol/l 4) Colloidal solutions.

Colloidal solutions are highly dispersed systems, where solid particles of the dispersed phase are evenly distributed in a liquid dispersion medium.

The structure of colloidal particles (using the example of AgI) - it explains the fact that an insoluble substance, i.e. the sediment is evenly distributed throughout the entire volume.

A prerequisite for obtaining a colloidal solution is an excess of one of the reactants.

– the core of a colloidal particle – the core of a micelle is always an insoluble compound.

Potential-determining ions are adsorbed on the surface of the nucleus (ions of the substance that are in excess).

(Ag + + NO 3 -) – adsorption layer is a change in the concentration of a substance at the phase boundary.

NO 3 - – counter-ions – fill the counterionic and diffuse (mobile) layers.

The micelle is electrically neutral, and the solid phase is always charged (its charge is determined by the charge of potential-determining ions).

Sols (German sole from Latin solutio - solution) are ultramicroheterogeneous disperse systems, the particle size of which ranges from 1 to 100 nm (10 −9 -10 −7 m).

Depending on the dispersion medium, sols are solid, aerosol (gaseous dispersion medium) and lyosols (liquid dispersion medium). Depending on the nature of the medium, lyosols are called hydrosols (water), organosols (organic medium) or, more specifically, alcosols (alcohols), etherosols, fats, etc. 3ols occupy an intermediate position between true solutions and coarse systems (suspensions, emulsions). Sols diffuse more slowly than inorganic salts and have a light scattering effect (Tyndall effect).

Ni(OH – core of colloidal particle

Potential-determining ions

( + ) – adsorption layer

– counter-ions

Ni(OH + + - solid phase

– diffuse layer 5) Abstract. Chemistry in construction.

§ 1. Solutions (definition). Concentration.

Solutions are called phases whose composition can be changed continuously (within certain limits), i.e. phases of variable composition 2. Thus, solutions are homogeneous mixtures of molecules (in particular cases, also atoms, ions) of two or more substances, between which there are physical and, often, chemical interactions.

Association of molecules of any compound and solvation (combination of molecules of a solute and molecules of a solvent into fragile complexes), which do not lead to the formation of particularly large associates, do not violate the homogeneity of the solution.

Mixtures in which the particles of one of the components of the mixture consist of a large number of molecules and are, as a rule, microcrystals with a complex structure of the surface layer have a different character. Such mixtures are heterogeneous, although at first glance they may seem homogeneous. They are microheterogeneous. These mixtures are called dispersed systems. Continuous transitions are possible between both classes of mixtures. However, the second part of our course is devoted to a detailed discussion of the properties of disperse systems.

Solutions, as a rule, are thermodynamically stable, and their properties do not depend on previous history, while dispersed systems are very often unstable and tend to change spontaneously.

The simplest components of a solution, which can be isolated in pure form and by mixing which one can obtain solutions of any possible composition, will be called components of the solution.

In many cases, their division into solvent and solutes is arbitrary. Typically, a component that is in excess compared to others is called a solvent, while the remaining components are called solutes. Thus, you can have solutions of alcohol or sulfuric acid in water and solutions of water in alcohol or sulfuric acid. If one of the components of a solution is a liquid and the others are gases or solids, then the solvent is considered to be a liquid.

The main parameters of the solution state, along with pressure and temperature, are concentrations, i.e., the relative amounts of components in a solution. Concentrations can be expressed in different ways in different units: quantities of components can be related to a known amount of solution or solvent, quantities of solutes can be expressed in units of weight and in moles; the amount of solvent or solution - in weight units, in moles and in volume units.

Let's look at some of the most commonly used methods and units for measuring solution concentrations. Let us denote the masses of the components, expressed in grams (“weight” quantities), by m 1 , m 2, ..., m i , and the sum of the masses of the components – through  m i; number of gram molecules or moles of components - through n 1 , n 2 , ..., n i , and their sum is  n i; volume of solution – through V, volumes of pure components – through V 1 , V 2 ... V i . Index 1 refers to a solvent in cases where one of the components of the solution can be unambiguously named as such.

Amounts of substances refer to a known amount of solution.

1. Mass fractionW i – mass of the component per unit mass of the solution:

(IV, 1a)

Mass percentage P i– mass of the component in one hundred mass units of the solution:

P i = 100W i. (IV, 1b)

2. Mole fractionx – number of moles of a component in one mole of solution:

(IV, 1c)

Mole fractions are most convenient in the theoretical (thermodynamic) study of solutions. From expression (IV, 1c) it is clear that

x i = 1

3. Volume fraction i– volume of pure component per unit volume of solution:

(IV,1g)

4. Mole-volume concentration – molarity C i– number of moles of a component per unit volume of solution:

(IV,1d)

In the case where the unit volume of a solution is a liter, the mole-volume concentration is called molarity.

5. Mole-weight ratio – the number of moles of a component per known weight of another component, usually a solvent. The mole-weight ratio, expressed by the number of moles of a component per 1000 g of solvent, is called molalityM i :

(IV, 1e)

Concentrations can also be expressed in other units.

You can move from one concentration unit to another by drawing up an equation for the relationship between these units. In the case of converting volumetric units of concentration to weight or molar units and vice versa, it is necessary to know the density of the solution. It should be remembered that only in very dilute solutions (i.e. for a component whose amount is small compared to others) the concentrations expressed in different units are proportional to each other.

§ 2. On the molecular structure of solutions

The idea of ​​a liquid as a completely amorphous phase in which the molecules are randomly located, like gas molecules, has now been abandoned. Studies on light and X-ray scattering have shown that liquids have elements of a crystalline structure (the presence of so-called short-range order in the arrangement of molecules) and in this regard are an intermediate formation between solid crystals and gases. As a liquid heats up, the similarity of its structure to crystals decreases and its similarity to gases increases.

The interaction between molecules in individual liquids is mainly van der Waals interaction. This name combines several types of intermolecular attraction, which are special cases of electrostatic interaction. These include: orientational attraction between molecules with a permanent dipole, inductive attraction between molecules with a permanent dipole and molecules with an induced dipole, and dispersive attraction between instantaneous dipoles of molecules whose moment oscillates around zero.

The energy of mutual attraction of molecules for all specified types of interaction is inversely proportional to the sixth power of the distance between them. These interactions in some cases lead to the association of liquid molecules (so-called associated liquids). Unstable bonds are formed between the molecules of the associated liquid. Such bonds include a hydrogen bond, which is created due to the electrostatic attraction of a proton of one molecule to an anion or electronegative atom (mainly fluorine, oxygen, nitrogen, chlorine atoms) of another molecule.

The attraction of molecules is counteracted by repulsion, which is important at short distances and is mainly due to the interaction of electron shells. This repulsion, combined with thermal motion, balances the attraction. Thus, average equilibrium distances are established between moving (oscillating, rotating and occasionally moving) liquid molecules.

A thermodynamic measure of molecular interaction in a liquid within known boundaries can be the quantity ( U/  V) P .

In a solution, along with the interaction between the molecules of one of the components (homogeneous molecules), there is interaction between the molecules of different components (dissimilar molecules). These interactions, in the absence of a chemical reaction, as in a pure liquid, are van der Waals. However, the molecules of the dissolved substance (the second component), changing the environment of the solvent molecule (the first component), can significantly change the intensity of interaction between the molecules of the latter and themselves interact with each other differently than in the pure second component. The interaction between dissimilar molecules may follow different patterns than the interaction between homogeneous molecules.

The tendencies toward association (the joining of similar molecules) and solvation (the joining of dissimilar molecules) are competing.

Let us consider here, as examples, diagrams depicting the dependence of some properties of binary liquid systems on their composition, from which we can establish the presence of a chemical compound between the components of the solution. Figure 4 shows the heat of mixing isotherms ( Q) components, volumetric compression (  V) during mixing and viscosity (  ) solutions of piperidine - allyl mustard oil (C 3 N 5 NCS). All properties exhibit a more or less sharp break at the maximum at a component ratio of 1:1. The break point at the maximum, called singular point, indicates the formation of a strong chemical compound containing components in the given ratio.

Fig.4. Dependence of some properties of the C 3 H 5 NCS – C 5 H 10 NH solution on the composition.

Thus, physicochemical analysis of single-phase liquid systems gives in some cases clear indications of the existence of certain chemical compounds. For the most part, the existence of certain compounds in solution cannot be established.

§ 3. On theories of solutions

For a long time, dissolution was viewed primarily as a chemical process. This view was also shared by D.I. Mendeleev, who excluded from consideration mixtures of liquids that were similar in nature (for example, mixtures of hydrocarbons). A different view of the dissolution process was developed by one of the prominent representatives of the “physical” theory of solutions, V.F. Alekseev, who outlined (1870–1880) a clear point of view on dissolution as the total result of molecular motion and mutual adhesion of molecules. Alekseev considered chemical interaction to be an important, but not obligatory factor in dissolution and argued with Mendeleev.

Subsequently, Mendeleev recognized the important role of the physical factor in the formation of solutions, but spoke out against the extreme, purely physical view of the nature of solutions. The physical theory of solutions received particular development after the 80s of the last century in connection with advances in the study of dilute solutions (Van't Hoff, Arrhenius, Ostwald). The first quantitative theory of solutions was created, associated with the idea of ​​a solute as a gas spreading in an inert solvent. However, it was soon discovered that the van't Hoff–Arrhenius quantitative theory was valid only for very dilute solutions. Many facts pointed to the interaction of the components of the solution. All attempts to consider solutions of any concentration from a single point of view led to the need to take into account the chemical factor.

In recent decades, the struggle between the two points of view has given way to recognition of the importance of both factors and the impossibility of opposing them. However, the complexity and variety of laws covering the properties of solutions of various substances make the theory of solutions the most difficult problem in molecular physics and the study of chemical bonds.

Deviations from the simplest properties are caused, for example, by the polarity of molecules. In solutions of polar molecules, association and solvation phenomena occur, as a result of which the properties of the solution become more complex. Deviations in the properties of a solution from the simplest ones are also caused by the chemical interaction of the components of the solution. It is usually accompanied by the release of heat and a decrease in the probability of transition of the component molecules, partially bound into more complex compounds, into the gas phase.

CHAPTER V. EQUILIBRIUM: LIQUID SOLUTION – SATURATED VAPOR

§ 1. Saturated vapor pressure of binary liquid solutions

The gaseous phase in equilibrium with a liquid solution (saturated vapor) contains, in general, all the components of the solution, and the saturated vapor pressure, which is also often called vapor pressure, is the sum of the partial pressures of the components 3. However, often the individual components are non-volatile at a given temperature and are practically absent in the gaseous phase.

The total vapor pressure (total or total pressure) and partial pressures are functions of temperature and solution composition. At a constant temperature, the state of a binary solution of components A and B is determined by one variable - the concentration of one of the components.

A convenient measure of concentration is the mole fraction. We will denote the mole fraction x 2 second component in solution through X. Obviously, the mole fraction of the first component X 1 = 1 – X. Boundaries of change X 1 And x 2 are zero and one; therefore, the diagram showing the dependence of the vapor pressure of a solution on its composition (pressure-composition diagram) has a finite extension. One of the possible types of diagram P – x for a solution of two liquids that are miscible in all respects (mole fraction X takes any value - from zero to one) is shown in Fig. 5. Extreme points of the curve P = f(x) are the saturated vapor pressures of pure liquids And . Total vapor pressure at any value X equal to the sum of the partial pressures of the components: P = P 1 + P 2 .

The composition of saturated steam is determined by the mole fractions of components in the vapor phase X" 1 and X" 2 ,. By definition of partial quantities (Dalton equation):

x" 1 =x" 2 =

§ 2. Raoult's law. Ideal solutions. Extremely diluted solutions

In the simplest case, the dependence of the partial vapor pressure of the solvent on the composition of the binary solution has the following form:

The partial pressure of a solvent in the vapor phase is proportional to its mole fraction in the solution.

Fig.5. Total and partial vapor pressures of a binary solution: dibromopropane - dibromoethane. Partial pressures on the diagram P – x are depicted by straight lines.

Equation (V, 1) can be given a different form:

(V, 2)

The relative decrease in the partial pressure of the solvent in the vapor phase is equal to the mole fraction of the solute (second component). Equations (V, 1) and (V, 2) are expressions Raoult's law (1886). Raoult's law, expressed in the form of equation (V, 1), is applicable to such solutions, the saturated vapor of which behaves like an ideal gas, and only a few solutions obey this law with sufficient accuracy, at any concentration (i.e., at values x, varying in the range from 0 to 1).

Typically, as the temperature increases (while the saturated vapor pressure is relatively low), deviations from Raoult's law in the form (V, 1) decrease. But at sufficiently high temperatures, when the saturated vapor pressure of the solution is very high, equation (V, 1) becomes inaccurate, since the deviations of the vapor from the ideal gas law increase.

Solutions that follow Raoult's law in the form of equation (V, 1) at all concentrations and all temperatures are called ideal (perfect) solutions, they are the ultimate, simplest type of liquid solutions.

It is easy to show that if equation (V, 1) is satisfied for the solvent vapor, then a similar equation must be observed for the vapor of the second, dissolved component

(V, 3)

Equations (V, 1) and (V, 3) reflect the properties of partial pressures of ideal solutions at low pressures. The set of these equations is called combined Raoult's law – Henry. In general, for a multicomponent ideal solution at low pressures we obtain:

(V, 4)

Equations (V, 1), (V, 3) and (V, 4) will further serve as starting points for studying the thermodynamic properties of ideal solutions at low pressures.

The total vapor pressure of an ideal binary solution is equal to

is also a linear function of the mole fraction.

Examples of ideal solutions (see Fig. 5) are mixtures: benzene - toluene, benzene - dichloroethane, hexane - octane and others.

The compositions of an ideal solution and its saturated vapor are different, i.e. X. In this case it is easy to find a connection between And X. In fact, the concentration of the second component in the vapor
. Substituting the value into this expression P 2 from Raoult's law (equation (V, 3)) and the value P from equation (V, 5), we obtain:

(V, 6)

From this it is clear that = x only when the saturated vapor pressures of both pure components are equal, i.e., when
.

§ 3. Real solutions. Positive and negative deviations from Raoult's law

Raoult's law does not hold for real solutions. The partial pressures of these solutions are greater or less than the vapor pressures of ideal solutions. Deviations from Raoult's law in the first case are called positive(total vapor pressure is greater than the additive value), and in the second case – negative(total vapor pressure is less than the additive value).

Examples of solutions with positive deviations from Raoult’s laws are the following solutions: acetone - ethyl alcohol, benzene - acetone, water - methyl alcohol.

Fig.6. Diagram of vapor pressure over a solution of C 6 H 6 - (CH 3) 2 CO.

Figure 6 shows a diagram P – X for one of these solutions (benzene - acetone).

Solutions with negative deviations from Raoult's laws include, for example, solutions: chloroform - benzene, chloroform - diethyl ether.

The diagram of vapor pressure over a solution of chloroform - diethyl ether is shown in Fig. 7.

Fig.7. Diagram of vapor pressure over a solution (C 2 H 5) 2 O – CHCl 3.

Total pressure values P in these systems change monotonically with a change in value X. If deviations from the law of ideal solutions are large, then the total vapor pressure curve passes through a maximum or minimum.

Positive and negative deviations of real solutions from Raoult’s law are caused by various factors. If dissimilar molecules in a solution attract each other with less force than homogeneous ones, this will facilitate the transition of molecules from the liquid phase to the gas phase (compared to pure liquids) and positive deviations from Raoult’s law will be observed. Strengthening the mutual attraction of dissimilar molecules in solution (solvation, formation of a hydrogen bond, formation of a chemical compound) makes it difficult for molecules to transition into the gas phase, so negative deviations from Raoult’s law will be observed.

It should be borne in mind that factors causing positive and negative deviations can act simultaneously in a solution, so the observed deviations are often the result of the superposition of deviations of opposite sign. The simultaneous action of opposing factors is especially clearly manifested in solutions in which the sign of deviations from the Raoult–Henry law changes with changes in concentration.

§ 4. Equilibrium diagrams of liquid – vapor in binary systems. Konovalov's first law. Fractional distillation

In Fig. 5, 6, 7, the total vapor pressure of a binary solution was presented as a function of the solution composition. As an argument, one can also use the composition of the vapor, which is determined by the partial pressure curves and differs from the composition of the liquid solution. In this way, you can obtain a second curve of the same property of the system - the total pressure of the saturated vapor of the solution, depending on another argument - the composition of the vapor.

Figure 8 shows a schematic diagram - the equilibrium isotherm of a binary solution - steam. Any point on the plane of the diagram characterizes the gross composition of the system (coordinate X) And pressure (coordinate P) and is called figurative point. The upper curve displays the dependence of saturated vapor pressure on the composition of the liquid, and the lower curve displays the dependence of saturated vapor pressure on the composition of the vapor. These curves divide the diagram plane into three fields. The top margin covers the values x andP, in which there is only one liquid phase - a solution of variable composition. The lower field corresponds to a gas mixture of variable composition. Any figurative point in the upper and lower fields depicts the state of one really existing phase. Field , enclosed between the two curves corresponds to a two-phase system. The system, the pressure and composition of which is reflected by the figurative point located in this field, consists of two phases - solution and saturated vapor. The composition of these phases is determined by the coordinates of the points lying at the intersection of the isobar passing through the figurative point of the system with the upper and lower curves. For example, a system characterized by a figurative point k, consists of two equilibrium phases, the composition of which is determined by the points A And b. Dot A, lying on the lower curve characterizes the composition of saturated steam, and the point b, lying on the upper curve is the composition of the solution. The lower curve is called the vapor branch, the upper curve is called the liquid branch.

Fig.8. Composition-pressure diagram of a binary system.

During isothermal compression of unsaturated steam of composition X 1 the figurative point of the system moves up vertically, steam condensation begins at the point A(Fig. 8) at a known pressure value P. The first drops of liquid have the composition X 2 ; the resulting liquid contains less component A than the condensing vapor.

With an isothermal decrease in pressure, the liquid composition is X 3 will begin to evaporate at a point d, giving steam composition x 4 (dot e); the resulting vapor contains more component A than the evaporating liquid. Consequently, component A always predominates in the vapor compared to the liquid in equilibrium with it, the addition of which to the system, as can be seen from the diagram, increases the total vapor pressure.

Based on the above, one can easily draw the following conclusion: saturated steam, compared to an equilibrium solution, is relatively richer in that component, the addition of which to the system increases the total steam pressure. This - Konovalov's first law (1881), which is valid for all stable solutions.

Let us also consider the phenomena of evaporation and condensation of solutions using the isobar diagram boiling point – solution composition.

Diagrams t kip. – X can be constructed from experimental data, or from a series of isothermal diagrams P – X. On every diagram P- X, constructed at a certain temperature, the compositions of the coexisting solution and steam at a given pressure are found. Based on the obtained from all isotherms P- X data for a certain pressure, one isobaric diagram is built t kip. – X.

Diagram t kip. – X is shown schematically in Fig. 9. Since component A with a higher saturated vapor pressure (Fig. 8) has a lower boiling point at a given pressure ( ), then the diagram t kip. - X has a mirror-like appearance in relation to the diagram P – X(there is only a qualitative similarity).

Top field on the chart t kip. – X corresponds to steam, and the bottom corresponds to liquid. The upper curve is the vapor branch, and the lower curve is the liquid branch.

Fig.9. Diagram boiling point - composition of a binary system.

In distillation columns, successive distillations are combined into one automated process, leading to the separation of the components of the liquid solution (rectification). The industrial separation of oil into fractions (primary oil refining) is based on the specified process.

In conclusion, we give another formulation of Konovalov’s first law:

In saturated steam, which is in equilibrium with a liquid binary solution, the relative content of the component that is at T = const has a higher vapor pressure value compared to another component or whenP = const has a lower boiling point in comparison, in other words, steam, compared to liquid, is relatively richer in the more volatile component.

I . Introduction.

In the last lesson, you wrote a paper on the topic “Main classes of inorganic substances.” The results of this work are as follows...

Until now we have talked about separate, individual substances; now we will move on to consider multicomponent systems consisting of 2 or more substances.

The topic of our lesson today is “Concentration of solutions.” writing a topic in a notebook

II . Testing knowledge and generalization.

Let's start by answering the following questions:

1 .What is a solution called?

(Solution- this is a homogeneous physicochemical system consisting of 2 or more parts (components), the relative quantities of which can vary within fairly wide limits.)

2. Why do we call a solution a homogeneous and physicochemical system?

(Homogeneous, because all components are in the same state of aggregation.)

3 .Which one?

(Solutions can be solid, liquid, gaseous. Examples of solid solutions are metal alloys, gaseous solutions are air).

A solution is a physicochemical system because

A) the homogeneity of solutions makes them very similar to chemical compounds;

B) The release of heat during the dissolution of some substances also indicates a chemical interaction between the solvent and the solute.

IN) The difference between solutions and chemical compounds is that the composition of the solution can vary within wide limits. When dissolved, substances can acquire chemical properties that do not appear in individual form.

G) In addition, in the properties of a solution one can detect many properties of its individual components, which is not observed in the case of a chemical compound.

Schema entry:

Solutions

Signs Signs

Physical system chemical system

(mechanical mixtures) (chemical compounds)

Inconsistency of composition - uniformity of composition

Manifestation of individual - highlighting Q due to

properties of components of destruction and formation of bonds

The variability of the composition of solutions brings them closer to mechanical ones. Thus, solutions occupy an intermediate position between mechanical mixtures and chemical compounds.

4. What substances are called solvents? Or how to determine which substance in a solution is a solvent and which is a solute?

(Solvent consider that component of the solution that, in its pure form, was in the same state of aggregation as the resulting solution. If both components were in identical states of aggregation before preparing the solution, then the solvent is considered to be the substance taken in larger quantities (for example, alcohol and water).

5. We found out what is a solution, a solute and a solvent. What is the dissolution process?

(Dissolution is the process of uniform distribution of particles (molecules, ions) of a solute between solvent molecules.

6. And only that? Do bonds break and form during dissolution? Explain.

(Yes, for example, when a crystal dissolves in a liquid, the following happens: [page 210 Glinka textbook]).

When a crystal is introduced into a liquid, individual molecules break off from its surface. The latter, due to diffusion, are evenly distributed throughout the entire volume of the solvent. The separation of molecules from the surface of a solid is caused, on the one hand, by their own vibrational movements, and on the other, by attraction from the solvent molecules. This process would have to continue until any number of crystals are completely dissolved, if the reverse process, crystallization, did not occur simultaneously. Molecules that have passed into solution, hitting the surface of a substance that has not yet dissolved, are again attracted to it and become part of its crystals.

7. Do you think the amount of a substance that can dissolve at a given temperature is unlimited?

(No. The ability of a substance to dissolve is always limited.)

8 . How can we determine how much of a substance can dissolve at given T, V?

(The amount of a substance that can dissolve at a given temperature in a certain amount of solvent is called solubility).

There are several ways to express solubility:

g or mole of dissolved substance in 1 dm 3 (1 l) solution – g/dm 3 solution (g/l solution) or g of substance in 100 g of solvent – ​​g/100 g solution, etc. .

The solubility of different substances in water can vary greatly. For example, at T = 25 C, the solubility of NaCl in H 2 O is 36 g per 100 g of H 2 O. Under the same conditions, the solubility of AgCl is only 0.00014 g in 100 g of H 2 O.

9 . What if we try to dissolve 40 g of NaCl in 100 g of H 2 O at T = 25 C?

(No, we won’t succeed, because each substance has its own solubility limit in H 2 O. Under these conditions, only 36 g of NaCl dissolves in 100 g of H 2 O, and 4 g will remain in the form of crystals (precipitate).

10. What is the name of a solution in which at a given temperature a substance no longer dissolves?

(Saturated. If a substance can still dissolve in a given solution, then it is called unsaturated).

From the above it follows that

III . Explanation of new material.

The concentration of a substance in a saturated solution is equal to its solubility; the concentration of a substance in an unsaturated solution is always less than the solubility value.

Do you think the solubility of substances depends on any conditions? Explain.

(Yes. The solubility of most solids increases with increasing temperature. For gases and liquids it is usually the other way around.

What solutions are called diluted and concentrated?

(The division of solutions into dilute and concentrated is conditional. For example, concentrated sulfuric acid contains 98% H 2 SO 4 and 2% H 2 O. A 40% solution of this acid is already considered dilute. At the same time, a concentrated solution of HCl contains only 36% HCl. Those solutions that contain no more than 1 mole of dissolved substance in 1 dm 3 (l) can be considered conditionally diluted.

Methods for quantitatively expressing the composition of solutions are of great importance for chemistry and technology. Which we now turn to.

Methods of expressing the quantitative composition of solutions.

The quantities of solute and solvent can be measured in different units. Because of this, there are several ways to express composition.

1. Mass fraction of solute

Mass fraction w shows what mass of solute is contained in 100 g of solution.

It must be remembered that m r-ra = m r.v-va + m r-la

Example: How many g of sugar must be dissolved in 500 g of water to prepare a solution with a mass fraction of 5%?

Solution: Let us denote the required amount of sugar by x, then

m r.v-va =x

m solution =x+ m H 2 O =x+500

2. Molar concentration of solution

Molar concentration C m shows the amount of solute? per mole, which is contained in 1 dm 3 (1 l) solution.

The calculation is carried out according to the formula:

Example: What is the molar concentration of a glucose solution in 500 cm 3 containing 2 g of glucose C 6 H 12 O 6?

Solution: Recall that the amount of substance is calculated as

Then C m = = mol/dm 3.

Example: The mass fraction of sucrose in the solution is 10%. What is the molar concentration of sucrose if the density of the solution is 1.1 g/cm3.

Solution: Let V solution = 1 dm 3 or 1 l, then 1 dm 3 has a mass of 1100 g, and m of sucrose is equal to:

x=110g; M(C 12 H 22 O 11) = 342 g/mol,

then C m = 0.322 mol/dm 3

Answer: 0.322 mol/dm 3

3. Molar concentration equivalent (normality).

To better understand the essence of this method, let's consider some basic concepts.

Equivalent called a real or conditional particle that can replace, attach, release or be in some other way equivalent to one hydrogen ion in acid-base or ion exchange reactions or one e in oxidation-reduction. Reactions.

For example, the equivalent of potassium hydroxide and hydrochloric acid will be a molecule of KOH and a molecule of HCl, sulfuric acid, respectively? H2SO4 molecules

HCl+NaOH=NaCl+H 2 O

2HCl+Ca(OH) 2 =CaCl 2 +2H 2 O

3HCl+Al(OH) 3 =AlCl 3 +3H 2 O

In the first reaction, one hydrogen ion is equivalent to a NaOH molecule; in the second reaction, one hydrogen yl is equivalent to a conventional particle-half of a Ca(OH) 2 molecule; in the third reaction the equivalent is one third of the Al(OH) 3 molecule.

Equivalence factor – a number indicating what fraction of an actual particle X is equivalent to one hydrogen ion in a given acid-base reaction or one electron in a given redox reaction.

The equivalence factor of a substance X is denoted as f equiv (X).

In the following reactions:

f eq (NaOH)=1; f eq (Ca(OH) 2 )=1/2; f eq (Al(OH) 3 )=1/3.

In all cases:

f eq (acid) = 1/basicity;

f eq (base) = 1/acidity;

f eq ( salts) = 1/number of cations * valency Me);

f eq (oxidizing agent) = 1/number of e taken

f eq (reducing agent) = 1/number of given e.

The molar mass of the equivalent of substance X is:

The product of the equivalence factor of substance X and its molar mass.

M eq(X)=f eq(X)*M(f) . Means,

M equiv(NaOH)=1*40=40g/mol;

M equiv(Ca(OH) 2 )=1/2*74=37g/mol;

M equiv(Al(OH) 3 )=1/3*78=g/mol.

The amount of equivalent substance is equal to the mass of substance X divided by the molecular (molar) mass of the equivalent.

For example,

Molar concentration of the equivalent of substance X (normality) Сн and is expressed in mol-eq/dm3.

Molar concentration equivalent– the ratio of the amount of equivalent substance to the volume of solution.

СН=, where V is the volume in dm 3

Example: 150 g of water was dissolved in 1 dm 3 of water. The density of the resulting solution is 1.1 g/cm 3 . Determine the molar concentration of the equivalent (normality).

Solution: Mass of the resulting solution: 1000+150=1150g

V size = 1150/1.1 = 1045 cm 3 = 10.45 dm 3

Eq. (H 2 SO 4 )=m. (H 2 SO 4 )/M eq. (H2SO4)

Eq. (H 2 SO 4 ) = 150/49 = 3.06 mol-eq.

With eq. . (H2SO4)=? eq. (H 2 SO 4 )/V = 3.06/1.045 = 2.93 mol-eq/dm 3

Answer: CH = 2.93 mol-equiv/dm 3

3. Molality of solution

Molality of solution – the amount of dissolved substance per 1 kg of solvent.

Сm – designation; expressed in mol/kg.

Сm=

Example: The mass fraction of potassium hydroxide in the solution is 10%. Calculate the molality of the solution.

Solution: A mass fraction of 10% indicates that for every 90 g of water there is 10 g of KOH. Let's calculate the mass of potassium hydroxide per 1 kg of water.

10g-----90g water,

X g-------1000g of water;

X= 10000/90=111g.

R.v-va = 111/59 g/mol = 2 mol.

Сm =2 mol/1kg=2mol/kg.

4. Mole fraction.

Mole fraction is the ratio of the amount of solute to the total amount of solute and solvent.

Denoted as N r.v.=

Before the task, you need to say that N r.v. + N r-la = 1

Solution: Find the amount of iodine and carbon tetrachloride in this solution:

?(J 2 )=20g/254g/mol=0.079mol

? (CCl 4 )=500g/154g/mol=3.25mol

N (J 2 )=0.079/(0.079+3.25)=0.024

N (СCl 4 )=1-0.024=0.976

Answer: 0.024; 0.976.

6. Title.

Titer is the number of grams of solute contained in 1 ml of solution, i.e.

D/s: Solve problems:

1.1.33 g of aluminum chloride was dissolved in 200 cm 3 of water. The density of the resulting solution is 1.05 g/cm3. Calculate the mass fraction of solute, molar concentration, molality, molar equivalent concentration, mole fractions of aluminum chloride and water.

2. Glinka’s problems “Problems and exercises in general chemistry” p. 103 No. 391-398,408,413,418,414,424,428.

The lesson is devoted to the topic “Phase states of substances. Dispersed systems. Ways of expressing concentration." You will become familiar with the definitions of phase or phase state of a substance, and become familiar with homogeneous and heterogeneous systems. Learn how mixtures of substances are classified. Get acquainted in detail with the dispersed system and its types, the colloidal system and its types and true solutions, with the concept of solution saturation and the solubility of a substance.

Topic: Solutions and their concentration, dispersed systems, electrolytic dissociation, hydrolysis

Lesson: Phase states of substances. Dispersed systems. Ways to Express Concentration

1. Classification of dispersion systems by state of aggregation

It is necessary to distinguish between the concepts of aggregate and phase states of substances.

Phase- a part of the system under consideration that is homogeneous in composition and properties, separated from other phases by interfaces on which some properties of the system change abruptly - for example, density, electrical conductivity, viscosity.

Phase is a homogeneous part of a heterogeneous system.

For example, if we pour sunflower oil into water, we will get a system that is in one state of aggregation - liquid. But the substances in it will be in two different phases: one is water, the other is vegetable oil, and there will be a distinct boundary between them, the so-called interface. This means that the system will be heterogeneous.

Another similar example: if we mix flour and granulated sugar, we get a system where the substances are in the same state of aggregation, but in two different phases, and the system is heterogeneous.

Rice. 1. Classification of disperse systems

It is not always possible to clearly draw the line between the concepts of “homogeneous” and “heterogeneous” systems. As the particle size increases, the mixture of substances is divided into coarse, colloidal solutions and true solutions. See fig. 1.

A system is called dispersed in which one substance in the form of small particles is distributed in the volume of another.

The dispersed phase is a substance that is present in a dispersion system in smaller quantities. It can also consist of several substances.

A dispersion medium is a substance that is present in a dispersion system in larger quantities, and in the volume of which the dispersed phase is distributed. Rice. 2.

Coarse systems

The dispersion medium and the dispersed phase can be composed of substances in different states of aggregation. Depending on the combination of dispersion medium and dispersed phase, 8 types of such systems are distinguished. See Table. 1.

Classification of disperse systems by state of aggregation

2. Some properties of coarse systems

Coarsely dispersed systems with a liquid or gaseous dispersion medium are gradually separated into their constituent components. The duration of such processes may vary. They determine the expiration dates and terms of possible use of both food products and other substances containing dispersed systems. In the case of substances with a solid dispersion medium, their properties are determined by the particle sizes of the dispersed phase. For example, pumice does not sink in water. Because water cannot penetrate the pores inside the stone, and the average density of the object becomes less than the density of water.

Coarsely dispersed systems can be easily viewed through a microscope. It was this property that the Scottish botanist John Brown took advantage of in 1827, who discovered the chaotic movement of tiny particles of pollen and other substances suspended in water. Later, this phenomenon was called Brownian motion and became one of the foundations of molecular kinetic theory.

3. Colloidal solutions

Colloidal systems

Colloidal systems are divided into:

· ZOLI- the dispersed phase does not form continuous rigid structures.

· GELS- particles of the dispersed phase form rigid spatial structures. Examples: cheese, bread, marmalade, marshmallows, jelly, jellied meat.

A solution of protein in water is a colloidal solution. Colloidal solutions are transparent but scatter light.

When light is passed through a transparent vessel containing a solution, a luminous cone can be observed.

Using a special microscope, individual particles can be detected in colloidal solutions.

Substances in a colloidal state take part in the formation of many minerals, such as agate, malachite, opal, carnelian, chalcedony, and pearls. See fig. 3.

There are gels in the human body. These are hair, cartilage, tendons. There are many sols and gels in the human body, so one of the scientists of Russian chemical science, I. I. Zhukov, said that a person is a walking colloid.

Coagulation

Coagulation is the sticking together of colloidal particles and their precipitation from solution.

Why does nature so often prefer the colloidal state of matter? The fact is that substances in a colloidal state have a large interface between phases. This facilitates the metabolic reaction that occurs precisely at the interface. For colloidal particles, their surface area is of great importance. Colloidal particles easily adsorb various substances on their surface. For example, ions seem to stick to their surface. In this case, colloidal particles acquire a positive or negative charge. Particles with the same charge will repel each other. If an electrolyte is added to a colloidal system, particles deprived of surface charge will begin to stick together into larger formations. Coagulation of the colloid occurs, which is accompanied by precipitation. Coagulation can also be caused by other influences, for example, heating. Such phenomena are of great importance both in nature and in industry.

4. True solutions, ways of expressing the composition of a solution

True solutions

The state of solutions is determined by the concentration of dissolved substances, temperature and pressure. When a substance is dissolved in a solvent at constant temperature and pressure, the concentration of the solute will not increase indefinitely. At some point it will stop dissolving and its maximum possible concentration will be reached. A dynamic equilibrium will occur, which consists in the fact that part of the substance constantly dissolves, and part passes from solution to precipitate. But the concentration of the dissolved substance will no longer change.

A saturated solution is a solution that is in phase equilibrium with the solute.

A supersaturated solution is a solution that contains more solute than a saturated solution at the same temperature and pressure.

Since the solubility of most substances increases with increasing temperature, a supersaturated solution can be obtained by cooling a solution saturated at a higher temperature. A supersaturated solution is unstable, and the introduction of a small crystal, dust, or even a sharp shock can cause rapid crystallization of the solute.

An unsaturated solution is a solution containing less solute than can be contained in a saturated solution of the same substance at the same temperature and pressure.

Solubility is the mass of a portion of a solute that, at a given temperature and pressure, must be dissolved in a certain amount of solvent to prepare a saturated solution. Most often, solubility is considered in 100 g, 1 kg or 1 liter of solvent.

To express the quantitative composition of a solution, the concept is used mass fraction of solute. This is the ratio of the solute to the total mass of the solution.

Molar concentration is the ratio of the amount of substance in moles to the volume of solution.

The areas of application of true solutions are very extensive. Therefore, it is very important to be able to prepare solutions of the corresponding substances.

Summing up the lesson

The lesson covered the topic “Phase states of substances. Dispersed systems. Ways of expressing concentration." You became familiar with the definitions of the phase or phase state of a substance, learned what homogeneous and heterogeneous systems are, and the classification of a mixture of substances. We got acquainted in detail with the dispersed system and its types, the colloidal system and its types and true solutions, with the concept of solution saturation and the solubility of a substance.

References

1. Rudzitis G. E. Chemistry. Fundamentals of general chemistry. 11th grade: textbook for general education institutions: basic level / G. E. Rudzitis, F. G. Feldman. - 14th ed. - M.: Education, 2012.

2. Popel P. P. Chemistry: 8th grade: textbook for general education institutions / P. P. Popel, L. S. Krivlya. - K.: IC "Academy", 2008. - 240 pp.: ill.

3. Gabrielyan O. S. Chemistry. 11th grade. Basic level. 2nd ed., erased. - M.: Bustard, 2007. - 220 p.

1. Internerok. ru.

2. Hemi. nsu. ru.

3. Chemport. ru.

Homework

1. No. 10-13 (p. 41) Rudzitis G. E. Chemistry. Fundamentals of general chemistry. 11th grade: textbook for general education institutions: basic level / G. E. Rudzitis, F. G. Feldman. - 14th ed. - M.: Education, 2012.

2. What determines the shelf life of cosmetic, medical and food gels?

3. How do colloidal systems illustrate the thesis of the relativity of truth?

7. Calculate the thermal effect of the reaction under standard conditions: Fe 2 O 3 (t) + 3 CO (g) = 2 Fe (t) + 3 CO 2 (g), if the heat of formation: Fe 2 O 3 (t) = – 821.3 kJ/mol; CO (g ) = – 110.5 kJ/mol;

CO 2 (g) = – 393.5 kJ/mol.

Fe 2 O 3 (t) + 3 CO (g) = 2 Fe (t) + 3 CO 2 (g),

Knowing the standard thermal effects of combustion of starting materials and reaction products, we calculate the thermal effect of the reaction under standard conditions:

16. Dependence of the rate of a chemical reaction on temperature. Van't Hoff's rule. Temperature coefficient of reaction.

Reactions result only from collisions between active molecules whose average energy exceeds the average energy of the participants in the reaction.

When molecules are given some activation energy E (excess energy above the average), the potential energy of interaction between atoms in molecules decreases, the bonds inside the molecules weaken, and the molecules become reactive.

The activation energy is not necessarily supplied from the outside; it can be imparted to some part of the molecules by redistributing energy during their collisions. According to Boltzmann, among N molecules there is the following number of active molecules N   possessing increased energy :

N  N·e – E / RT (1)

where E is the activation energy, which shows the necessary excess of energy, compared to the average level, that molecules must have in order for the reaction to become possible; the remaining designations are well known.

With thermal activation for two temperatures T 1 and T 2, the ratio of rate constants will be:

, (2)

where , (3)

which makes it possible to determine the activation energy by measuring the reaction rate at two different temperatures T 1 and T 2.

An increase in temperature by 10 0 increases the reaction rate by 2–4 times (approximate Van't Hoff rule). The number showing how many times the reaction rate (and therefore the rate constant) increases when the temperature increases by 10 0 is called the temperature coefficient of the reaction:

 (4)

Or .(5)

This means, for example, that with an increase in temperature by 100 0 for a conventionally accepted increase in the average rate by 2 times ( = 2), the reaction rate increases by 2 10, i.e. approximately 1000 times, and when  = 4 – 4 10, i.e. 1000000 times. Van't Hoff's rule is applicable for reactions occurring at relatively low temperatures in a narrow temperature range. The sharp increase in the reaction rate with increasing temperature is explained by the fact that the number of active molecules increases exponentially.

25. Van't Hoff chemical reaction isotherm equation.

In accordance with the law of mass action for an arbitrary reaction

and A + bB = cC + dD

the rate equation for the forward reaction can be written:

,

and for the rate of reverse reaction: .

As the reaction proceeds from left to right, the concentrations of substances A and B will decrease and the rate of the forward reaction will decrease. On the other hand, as reaction products C and D accumulate, the rate of the reaction from right to left will increase. There comes a moment when the speeds υ 1 and υ 2 become the same, the concentrations of all substances remain unchanged, therefore,

WhereK c = k 1 / k 2 = .

The constant value Kc, equal to the ratio of the rate constants of the forward and reverse reactions, quantitatively describes the state of equilibrium through the equilibrium concentrations of the starting substances and the products of their interaction (to the extent of their stoichiometric coefficients) and is called the equilibrium constant. The equilibrium constant is constant only for a given temperature, i.e.

K c = f (T). The equilibrium constant of a chemical reaction is usually expressed as a ratio, the numerator of which is the product of the equilibrium molar concentrations of the reaction products, and the denominator is the product of the concentrations of the starting substances.

If the reaction components are a mixture of ideal gases, then the equilibrium constant (K p) is expressed in terms of the partial pressures of the components:

K p = .

To move from K p to K c, we use the equation of state P · V = n · R · T. Because

Then P = C R T.

From the equation it follows that K p = K c provided that the reaction proceeds without changing the number of moles in the gas phase, i.e. when (c + d) = (a + b).

If the reaction proceeds spontaneously at constant P and T or V and T, then the values ​​of G and F of this reaction can be obtained from the equations:

,

where С А, С В, С С, С D are nonequilibrium concentrations of starting substances and reaction products.

where Р А, Р В, Р С, Р D are the partial pressures of the starting substances and reaction products.

The last two equations are called van't Hoff isotherm equations for a chemical reaction. This relationship makes it possible to calculate the values ​​of G and F of the reaction and determine its direction at different concentrations of the starting substances.

It should be noted that for both gas systems and solutions, when solids participate in the reaction (i.e. for heterogeneous systems), the concentration of the solid phase is not included in the expression for the equilibrium constant, since this concentration is almost constant. Yes, for reaction

2 CO (g) = CO 2 (g) + C (t)

the equilibrium constant is written as

The dependence of the equilibrium constant on temperature (for temperature T 2 relative to temperature T 1) is expressed by the following van't Hoff equation:

,

where Н 0 is the thermal effect of the reaction.

For an endothermic reaction (the reaction occurs with the absorption of heat), the equilibrium constant increases with increasing temperature, the system seems to resist heating.

34. Osmosis, osmotic pressure. Van't Hoff equation and osmotic coefficient.

Osmosis is the spontaneous movement of solvent molecules through a semi-permeable membrane that separates solutions of different concentrations, from a solution of lower concentration to a solution of higher concentration, which leads to the dilution of the latter. A cellophane film is often used as a semi-permeable membrane, through the small holes of which only small-volume solvent molecules can selectively pass through and large or solvated molecules or ions are retained - for high-molecular substances, and a copper ferrocyanide film for low-molecular substances. The process of solvent transfer (osmosis) can be prevented if external hydrostatic pressure is applied to a solution with a higher concentration (under equilibrium conditions this will be the so-called osmotic pressure, denoted by the letter ). To calculate the value of  in solutions of non-electrolytes, the empirical Van't Hoff equation is used:

where C is the molal concentration of the substance, mol/kg;

R – universal gas constant, J/mol K.

The magnitude of osmotic pressure is proportional to the number of molecules (in general, the number of particles) of one or more substances dissolved in a given volume of solution, and does not depend on their nature and the nature of the solvent. In solutions of strong or weak electrolytes, the total number of individual particles increases due to the dissociation of molecules, therefore, an appropriate proportionality coefficient, called the isotonic coefficient, must be introduced into the equation for calculating osmotic pressure.

i C R T,

where i is the isotonic coefficient, calculated as the ratio of the sum of the numbers of ions and undissociated electrolyte molecules to the initial number of molecules of this substance.

So, if the degree of dissociation of the electrolyte, i.e. the ratio of the number of molecules disintegrated into ions to the total number of molecules of the dissolved substance is equal to  and the electrolyte molecule disintegrates into n ions, then the isotonic coefficient is calculated as follows:

i = 1 + (n – 1) · ,(i > 1).

For strong electrolytes, we can take  = 1, then i = n, and the coefficient i (also greater than 1) is called the osmotic coefficient.

The phenomenon of osmosis is of great importance for plant and animal organisms, since the membranes of their cells in relation to solutions of many substances have the properties of a semi-permeable membrane. In pure water, the cell swells greatly, in some cases to the point of rupture of the membrane, and in solutions with a high concentration of salts, on the contrary, it decreases in size and wrinkles due to large loss of water. Therefore, when preserving foods, large amounts of salt or sugar are added to them. Microbial cells under such conditions lose a significant amount of water and die.

Osmotic pressure ensures the movement of water in plants due to the difference in osmotic pressure between the cell sap of plant roots (5-20 bar) and the soil solution, which is further diluted during watering. Osmotic pressure causes water to rise in the plant from the roots to the top. Thus, leaf cells, losing water, osmotically absorb it from stem cells, and the latter take it from root cells.

Equivalent electrical conductivity is the value of electrical conductivity per one mole equivalent of an electrolyte:

,

where λ is the equivalent electrical conductivity, Ohm – 1 cm 2 mol – 1;

C eq – molar concentration of equivalents of electrolyte solution, mol/l.

Electrical conductivity is the reciprocal of resistivity:

where cspecific electrical conductivity, Ohm – 1. cm – 1;

ρ – electrical resistivity, Ohm cm.

52. Distinctive features of dispersed systems from true solutions. The mechanism of manifestation of each distinctive feature.

A true solution is a homogeneous mixture consisting of solutes and a solvent. In true solutions, the dissolved substances are either in a molecularly dispersed or ionically dispersed state.

Truly soluble particles determine, in particular, osmotic pressure, the osmotic phenomena of lowering the freezing point and increasing the boiling point.

Highly dispersed heterogeneous systems, unlike solutions, most often contain 2 phases. One phase represents highly dispersed tiny particles of a substance or macromolecule of the Navy and is called the dispersed phase. The other phase, in which aggregates of dispersed particles or macromolecules are distributed, is called a dispersion medium. The condition for the formation of such disperse systems (colloidal state of matter) is the insolubility of one phase in another.

The dispersed phase, consisting of many tiny particles, has a very large specific interface with the dispersion medium. The special properties of the interface determine the specific characteristics of disperse systems, which is the reason for the separation of this field of knowledge into a separate science - colloidal chemistry.

The main distinguishing features of dispersed systems from true solutions are:

a) the ability to scatter light;

b) slow diffusion of dispersed phase particles in a dispersion medium;

c) ability to dialysis;

d) aggregative instability of the dispersed phase, which is determined by the release of particles from the dispersion medium when electrolytes are added to the system or under the influence of other external influences.

61. Calculate the average shift of aerosol particles with a particle radius of 10 -7 m over a time of 10 s at a temperature of 273 K and an air viscosity of 1.7 10 -5 n s/m 2. How will the average particle shift change if the radius of the aerosol particles increases to 10 -6 m?

Average shear of aerosol particles

time during which the particle is displaced (diffusion duration), s;

D  diffusion coefficient, m 2. s -1 .

The diffusion coefficient for a spherical particle is calculated using the Einstein equation:

,

where N А is Avogadro’s number, 6 10 23 molecules/mol;

h – viscosity of the dispersion medium, N s/m 2 (Pa s);

r – particle radius, m;

R – universal gas constant, 8.314 J/mol K;

T – absolute temperature, K;

number 3.14.

answer to the second question of the assignment:

Thus, the average particle shift will decrease by 10 times.

Answer: , will decrease by 10 times.

80. Adsorption of ions on a solid surface. The concept of ionites. Reversible ion exchange adsorption is the basis of ion exchange chromatography.

The physical processes of molecular adsorption on a solid surface are described by the Langmuir and Freundlich equations.

Langmuir equation:

,

where G is the adsorption value, kmol/kg or kmol/m2;

Г max – value of limiting adsorption, kmol/kg (kmol/m2);

C – solution concentration, kmol/l;

a is the adsorption equilibrium constant.

This equation describes adsorption well for low and high concentrations of solutions (or gas pressures).

Empirical Freundlich equation:

,

where G is the adsorption value, kmol/kg (kmol/m2);

n – amount of adsorbent substance, kmol;

m – adsorbent mass, kg;

K – constant (at C = 1 mol/l K = G);

1/a – constant (adsorption indicator); depends on the nature of the adsorbent and temperature. 1/a = 0.1–1.

Adsorption chromatography is based on the difference in the sorption of the separated substances by the adsorbent (a solid with a developed surface); partition chromatography - on different solubilities of mixture components in the stationary phase (high-boiling liquid deposited on a solid macroporous carrier) and eluent (it should be borne in mind that with the distribution mechanism of separation, the movement of component zones is partially influenced by the adsorption interaction of the analyzed components with the solid sorbent) .

Ion exchange chromatography is based on the difference in ion exchange equilibrium constants between the stationary phase (ion exchanger) and the components of the mixture being separated.

If an electrolyte is already adsorbed on the surface of the adsorbent, then when this adsorbent comes into contact with another electrolyte, ion-exchange adsorption almost always occurs to one degree or another. It is observed on a surface with a fairly pronounced electrical double layer. Mobile counterions of the electrical layer are able to exchange for other ions of the same sign in the solution.

A quantitative description of the ion exchange process (reversibility of the process, equivalence of exchange, order of ion exchange) was made by Giedroyc already at the beginning of the 20th century. Substances that exhibit the ability to ion exchange and are used for the adsorption of ions are called ion exchangers or ion exchangers.

Ion exchangers have a framework structure “cross-linked” by covalent bonds. The frame (matrix) has a positive or negative charge, which is compensated by the opposite charge of mobile ions - counterions located in the adsorption and diffuse parts of the double electrical layer. Counterions can be replaced by other ions from the solution with a charge of the same sign, and the framework acts as a polyion and determines the insolubility of the ion exchanger in the solvent.

Ion exchangers are divided by composition into organic and inorganic, by origin - into natural and synthetic, by the nature of the exchanged ions - into cation exchangers, anion exchangers and ampholytes.

Of the natural inorganic cation exchangers, crystalline silicates of the zeolite type are most often used: chabazite, glacuonite, etc. Their frame consists of a network structure of aluminosilicates, in the pores of which ions of alkali or alkaline earth metals, which are counter-ions, are located. Natural ion exchangers include apatites.

Natural organic ion exchangers are humic substances in soils containing a carboxyl group capable of ion exchange. The substances that make up the soil have amphoteric properties and therefore, depending on conditions, can exchange both cations and anions. However, natural ion exchangers are not widely used due to chemical instability and low exchange capacity.

Synthetic ion exchangers are used industrially, and among them the most widely used are ion exchange resins, which have a network structure and contain ionogenic groups: - OH, COOH, SO 3 H, - COONa, etc.

89. Write the formula for the structure of a sol micelle formed as a result of the interaction of the indicated substances (an excess of one, then another substance): CdCl 2 + Na 2 S ; FeCl 3 + NaOH . Name the constituent components micelles.

1) CdCl 2 + Na 2 S

Excess CdCl 2 gives a micelle:

[ (CdCl 2) Cd 2+ Cl – ] + x Cl –

germ: (CdCl 2)

core: [ (CdCl 2) Cd 2+

Excess Na 2 S gives a micelle:

–xNa+

germ: (NaCl)

core: (NaCl)2Cl -

granule: [ (CdCl 2) Cd 2+ Cl – ] +

2) FeCl 3 + NaOH

Excess FeCl 3 gives a micelle:

[ (FeCl 3) Fe 3+ 2Cl – ] + x Cl –

germ: (FeCl 3)

core: (FeCl 3) Fe 3+

granule: [ (FeCl 3) Fe 3+ 2Cl – ] +

Excess NaOH gives a micelle:

–xNa+

germ: (NaCl)

core: 3 (NaCl) 3 Cl –

granule: –

98. Coagulation of a dispersed system. Coagulation rate. The reasons causing the process of spontaneous coagulation.

Coagulation is a process of spontaneous enlargement (sticking together) of dispersed particles, which can occur when a dispersed system is exposed to various factors: intense stirring or shaking, heating or cooling, irradiation with light or passing an electric current, when adding electrolytes or non-electrolytes to the system, etc. When In different ways of influencing the system, the binding energy of dispersed particles with the surrounding dispersion medium decreases. Thus, the addition of an electrolyte causes compression of the diffuse layer in the colloidal particle, hence a decrease in the value of the electrokinetic potential. This leads to a decrease in the electrostatic repulsion of colloidal particles and, as a consequence, to a greater likelihood of their sticking together.

The minimum concentration of electrolyte added to the dispersed system, at which obvious coagulation occurs over a certain period of time, is called the coagulation threshold. The coagulation threshold is determined by temperature, the nature of the added electrolyte, the sign of the charge of the added ion (primarily the ion that has a charge opposite to the charge of the colloidal particles acts) and the magnitude of the charge of this ion. Thus, for three-, doubly and singly charged ions, obvious coagulation occurs at a concentration of electrolytes in the ratio 1:10 - 50: 500-1000 (approximate Schulze-Hardy rule).

The coagulation threshold is calculated as follows:

, (1)

where coagulation threshold, kmol/m 3 ;

C – molar concentration of electrolyte solution, kmol/m3;

V el – volume of electrolyte solution, m 3 ;

V sol – volume of sol, m3.

The theory of rapid coagulation speed was developed by Smoluchowski. He is the author of the equation for calculating the coagulation rate constant K:

, (2)

where n 0 and n are the number of particles per unit volume of the system before the start of coagulation and at the moment of time, respectively

 coagulation time, s.

The coagulation rate constant depends on the diffusion coefficient for the particles and their radius as follows:

К = 8 · ·D · r.(3)

Taking into account equation (2) and Einstein's equation, the final equation for the coagulation rate constant takes the form:

, (4)

where K is the coagulation rate constant, m 3 /s;

– viscosity of the medium, Pa s;

N A – Avogadro's number.

Smoluchowski also introduced the concept of half coagulation time, according to which the time required to reduce the initial number of particles by 2 times is related to their initial number as follows:

where  – time of half coagulation, s;

 – time from the beginning of coagulation, s.

From the equation converted to:

it follows that if the graph constructed in coordinates n o /n = f ( is a straight line, then this serves as an indicator of the correspondence of the experimental data to Smoluchowski’s theory.

108. Suspensions. Conditions of their formation and properties. Pastes are concentrated suspensions. Examples of suspensions among food products.

Suspensions are suspensions of powders in liquid (type T/F). The dispersed phase in suspensions contains particles of relatively large sizes (more than 10–4 cm), therefore suspensions are sedimentationally (i.e., in terms of their ability to settle) unstable. They are not characterized by osmotic pressure, Brownian motion and diffusion. Particles can carry a double electrical layer on their surface, which contributes to their stabilization, but under the influence of electrolytes, suspensions coagulate or form aggregates, and a coagulated suspension is usually easily peptized. Suspensions of hydrophilic particles are stable in water, but unstable in hydrocarbons. Their resistance increases in the presence of surfactants. The formation of a charge on the surface of particles (micellization) also contributes to increasing the stability of suspensions.

Pastes are highly concentrated stabilized suspensions (T/L type), in which particles of the dispersed phase are bound due to molecular forces and for this reason are not capable of mutual movement. In such highly viscous (plastic-viscous) systems, almost the entire dispersion medium is solvately associated with the dispersed phase. Thus, pastes occupy an intermediate position between powders and suspensions. Processes characteristic of colloidal systems with an internal structure (syneresis, etc.) can occur in them. The great practical importance of such concentrated systems is due to their plasticity.

Examples of suspensions among food products are all food spreads: tomato, chocolate, cheese, etc.

LITERATURE

1. Akhmetov B.V. Problems and exercises in physical and colloidal chemistry. – L.: Chemistry, 1989.

2. Gameeva O. S. Physical and colloidal chemistry. – M.: Higher School, 1983.

3. Evstratova K. I., Kupina N. A., Malakhova E. M. Physical and colloidal chemistry. – M.: Higher School, 1990.

4. Zimon A. D., Leshchenko N. F. Colloid chemistry. – M.: Chemistry, 2001.

5. Zimon A. D., Leshchenko N. F. Physical chemistry. – M.: Chemistry, 2000.

6. Kiselev E.V. Collection of examples and problems in physical chemistry. – M.: Higher School, 1983.

7. Knorre D. G. Physical chemistry. – M.: Higher School, 1990.

8. Stromberg A.G. Physical chemistry. – M.: Higher School, 2001.

9. Stepin B. D. International systems of units of physical quantities in chemistry. – M.: Higher School, 1990.

10. Friedrichsberg D. A. Course in colloid chemistry. – L.: Chemistry, 1995.

11. Khmelnitsky R. A. Physical and colloidal chemistry. – M.: Higher School, 1988.