Solubility is the ability of a substance to form with various solvents. homogeneous mixtures. As we have already mentioned, the amount of solute required to obtain a saturated solution determines this substance. In this regard, solubility has the same measure as composition, such as the mass fraction of a solute in its saturated solution or the amount of solute in its saturated solution.

All substances from the point of view of their solubility can be classified into:

• Well soluble - more than 10 g of substance can dissolve in 100 g of water.
• Slightly soluble - less than 1 g of substance can dissolve in 100 g of water.
• Insoluble - less than 0.01 g of substance can dissolve in 100 g of water.

It is known that if polarity The polarity of the solute is similar to the polarity of the solvent, then it is likely to dissolve. If the polarities are different, then with a high degree of probability the solution will not work. Why is this happening?

Polar solvent – polar solute.

As an example, let us describe a solution of table salt in water. As we already know, water molecules are polar in nature with a partial positive charge on each hydrogen atom and a partial negative charge on the oxygen atom. And ionic solids, like sodium chloride, contain cations and anions. Therefore, when table salt is placed in water, partial positive charge on the hydrogen atoms of the water molecules is attracted by the negatively charged chlorine ion in NaCl. Likewise, partial negative charge on the oxygen atoms of the water molecules is attracted by the positively charged sodium ion in NaCl. And, since the attraction of water molecules for sodium and chlorine ions stronger interaction holding them together, the salt dissolves.

Non-polar solvent – a non-polar solute.

Let's try to dissolve a piece of carbon tetrabromide in carbon tetrachloride. In the solid state, carbon tetrabromide molecules are held together by very weak dispersion interactions. When placed in carbon tetrachloride, its molecules will be arranged more chaotically, i.e. the entropy of the system increases and the compound dissolves.

Dissolution equilibria

Consider a solution of a slightly soluble compound. In order for equilibrium to be established between a solid and its solution, the solution must be saturated and in contact with the undissolved part of the solid.

For example, let equilibrium be established in a saturated solution of silver chloride:

AgCl(s)=Ag + (aq) + Cl - (aq)

The compound in question is ionic and is present in the form of ions when dissolved. We already know that in heterogeneous reactions the concentration of the solid remains constant, which allows us to include it in the equilibrium constant. Therefore, the expression for will look like this:

K = [Cl - ]

This constant is called solubility product PR, provided that concentrations are expressed in mol/l.

PR = [Cl - ]

Solubility product is equal to the product of the molar concentrations of the ions participating in the equilibrium, in powers equal to the corresponding stoichiometric coefficients in the equilibrium equation.
It is necessary to distinguish between the concept of solubility and product of solubility. The solubility of a substance can change when another substance is added to the solution, and the solubility product does not depend on the presence of additional substances. Although these two quantities are interrelated, which allows knowing one quantity to calculate the other.

Dependence of solubility on temperature and pressure

Water plays important role in our lives, it is capable of dissolving a large number of substances that have great importance for us. Therefore, we will focus on aqueous solutions.

Solubility gases rise with increase in pressure gas over the solvent, and the solubility of solids and liquid substances depends on pressure insignificantly.

William Henry first came to the conclusion that the amount of gas that dissolves at a constant temperature in a given volume of liquid is directly proportional to its pressure. This statement known as Henry's law and it is expressed by the following relation:

С = k·P,

where C is the solubility of gas in the liquid phase

P – gas pressure above the solution

k – Henry’s constant

The following figure shows the solubility curves of some gases in water on temperature at constant gas pressure above the solution (1 atm)

As can be seen, the solubility of gases decreases with increasing temperature, in contrast to most ionic compounds, the solubility of which increases with increasing temperature.

Effect of temperature on solubility depends on the enthalpy change that occurs during the dissolution process. During an endothermic process, solubility increases with increasing temperature. This follows from what we already know : If you change one of the conditions under which the system is in a state of equilibrium - concentration, pressure or temperature - then the equilibrium will shift in the direction of the reaction that counteracts this change.

Let's imagine that we are dealing with a solution that is in equilibrium with a partially dissolved substance. And this process is endothermic, i.e. goes with the absorption of heat from outside, then:

Substance + solvent + heat = solution

According to Le Chatelier's principle at endothermic process, the equilibrium shifts in a direction that contributes to a decrease in heat input, i.e. to the right. Thus, solubility increases. If the process exothermic, then an increase in temperature leads to a decrease in solubility.

Dependence of solubility of ionic compounds on Temperature

It is known that there are solutions of liquids in liquids. Some of them can dissolve in each other in unlimited quantities, like water and ethanol, while others dissolve only partially. So, if you try to dissolve carbon tetrachloride in water, then two layers are formed: the upper one is a saturated solution of water in carbon tetrachloride and the lower one is a saturated solution of carbon tetrachloride in water. As the temperature increases, the mutual solubility of such liquids generally increases. This happens until it is reached critical temperature, in which both liquids are mixed in any proportions. The solubility of liquids is practically independent of pressure.

When a substance that can dissolve in either of these two liquids is introduced into a mixture consisting of two immiscible liquids, its distribution between these liquids will be proportional to its solubility in each of them. Those. according to distribution law a substance capable of dissolving in two immiscible solvents is distributed between them so that the ratio of its concentrations in these solvents at a constant temperature remains constant, regardless of the total amount of solute:

C 1 / C 2 = K,

where C 1 and C 2 are the concentrations of the substance in two liquids

K – distribution coefficient.

Categories ,

SOLUBILITY The ability of a substance to dissolve in a particular solvent is called. A measure of the solubility of a substance under given conditions is its content in a saturated solution . If more than 10 g of a substance dissolves in 100 g of water, then such a substance is called highly soluble. If less than 1 g of a substance dissolves, the substance slightly soluble. Finally, the substance is considered practically insoluble, if less than 0.01 g of substance goes into solution. There are no absolutely insoluble substances. Even when we pour water into a glass vessel, a very small part of the glass molecules inevitably goes into solution.

Solubility, expressed in terms of the mass of a substance that can dissolve in 100 g of water at a given temperature, is also called solubility coefficient.

The solubility of some substances in water at room temperature.

Solubility of most (but not all!) solids with increasing temperature it increases, and the solubility of gases, on the contrary, decreases. This is primarily due to the fact that gas molecules thermal movement are able to leave solution much more easily than solid molecules.

If we measure the solubility of substances at different temperatures, then it will be found that some substances noticeably change their solubility depending on temperature, others - not very much

When dissolving solids in water the volume of the system usually changes slightly. Therefore, the solubility of substances in the solid state is practically independent of pressure.

Liquids can also dissolve in liquids. Some of them are unlimitedly soluble in one another, that is, they mix with each other in any proportions, such as alcohol and water, while others dissolve mutually only to a certain limit. So, if you shake diethyl ether with water, two layers are formed: the upper one is a saturated solution of water in ether, and the lower one is a saturated solution of ether in water. In most such cases, with increasing temperature, the mutual solubility of liquids increases until a temperature is reached at which both liquids mix in any proportions.

Dissolution of gases in water is an exothermic process. Therefore, the solubility of gases decreases with increasing temperature. If you leave a glass with cold water, then its inner walls are covered with gas bubbles - this is air that was dissolved in water and is released from it due to heating. Boiling can remove all dissolved air from water.

A solution is a homogeneous system consisting of two or more substances, the content of which can be changed within certain limits without disturbing the homogeneity.

Water solutions consist of water(solvent) and dissolved substance. State of substances in aqueous solution if necessary, it is indicated by a subscript (p), for example, KNO 3 in solution - KNO 3 (p).

Solutions that contain a small amount of solute are often called diluted and solutions with high content solute – concentrated. A solution in which further dissolution of a substance is possible is called unsaturated and a solution in which a substance ceases to dissolve under given conditions is saturated. The latter solution is always in contact (in heterogeneous equilibrium) with an undissolved substance (one crystal or more).

IN special conditions, for example, when carefully (without stirring) cooling a hot unsaturated solution solid substances that can form oversaturated solution. When a crystal of a substance is introduced, such a solution is divided into a saturated solution and a precipitate of the substance.

In accordance with chemical theory solutions D.I. Mendeleev, the dissolution of a substance in water is accompanied, firstly, by destruction chemical bonds between molecules ( intermolecular bonds V covalent substances) or between ions (in ionic substances), and thus the particles of the substance mix with water (in which part of the hydrogen bonds between molecules). The breaking of chemical bonds occurs due to the thermal energy of movement of water molecules, and this occurs cost energy in the form of heat.

Secondly, once in water, particles (molecules or ions) of the substance are subjected to hydration. As a result, hydrates– compounds of uncertain composition between particles of matter and water molecules ( internal composition the particles of the substance themselves do not change upon dissolution). This process is accompanied highlighting energy in the form of heat due to the formation of new chemical bonds in hydrates.

In general, the solution is either cools down(if the heat consumption exceeds its release), or heats up (in otherwise); sometimes - if the heat input and its release are equal - the temperature of the solution remains unchanged.

Many hydrates turn out to be so stable that they do not collapse even when the solution is completely evaporated. Thus, solid crystalline hydrates of the salts CuSO 4 5H 2 O, Na 2 CO 3 10H 2 O, KAl(SO 4) 2 12H 2 O, etc. are known.

The content of a substance in a saturated solution at T= const quantitatively characterizes solubility of this substance. Solubility is usually expressed as the mass of solute per 100 g of water, for example 65.2 g KBr/100 g H 2 O at 20 °C. Therefore, if 70 g of solid potassium bromide is added to 100 g of water at 20 °C, then 65.2 g of salt will go into solution (which will be saturated), and 4.8 g of solid KBr (excess) will remain at the bottom of the glass.

It should be remembered that the solute content in rich solution equals, V unsaturated solution less and in oversaturated solution more its solubility at a given temperature. Thus, a solution prepared at 20 °C from 100 g of water and sodium sulfate Na 2 SO 4 (solubility 19.2 g/100 g H 2 O), containing

15.7 g salt – unsaturated;

19.2 g salt – saturated;

2O.3 g of salt – supersaturated.

The solubility of solid substances (Table 14) usually increases with increasing temperature (KBr, NaCl), and only for some substances (CaSO 4, Li 2 CO 3) the opposite is observed.

The solubility of gases decreases with increasing temperature, and increases with increasing pressure; for example, at a pressure of 1 atm, the solubility of ammonia is 52.6 (20 °C) and 15.4 g/100 g H 2 O (80 °C), and at 20 °C and 9 atm it is 93.5 g/100 g H 2 O.

In accordance with solubility values, substances are distinguished:

– highly soluble, the mass of which in a saturated solution is comparable to the mass of water (for example, KBr - at 20 °C solubility 65.2 g/100 g H 2 O; 4.6 M solution), they form saturated solutions with a molarity of more than 0.1 M;

– slightly soluble, whose mass in a saturated solution is significantly less than the mass of water (for example, CaSO 4 - at 20 °C solubility 0.206 g/100 g H 2 O; 0.015 M solution), they form saturated solutions with a molarity of 0.1–0.001 M;

– practically insoluble, whose mass in a saturated solution is negligible compared to the mass of the solvent (for example, AgCl - at 20 °C solubility 0.00019 g per 100 g H 2 O; 0.0000134 M solution), they form saturated solutions with a molarity of less than 0.001 M.

Compiled based on reference data solubility table common acids, bases and salts (Table 15), which indicates the type of solubility, substances that are not known to science(not obtained) or completely decomposed by water.

Legend, used in the table:

"r" - good soluble substance

“m” – slightly soluble substance

“n” – practically insoluble substance

“-” – substance not received (does not exist)

“” – the substance mixes with water unlimitedly

Note. This table corresponds to the preparation of a saturated solution at room temperature by adding a substance (in the appropriate state of aggregation) in water. It should be taken into account that obtaining precipitation of poorly soluble substances using ion exchange reactions is not always possible (for more details, see 13.4).

13.2. Electrolytic dissociation

The dissolution of any substance in water is accompanied by the formation of hydrates. If at the same time no formula changes occur in the particles of the dissolved substance in the solution, then such substances are classified as non-electrolytes. They are, for example, gas nitrogen N 2, liquid chloroform CHCl 3, solid sucrose C 12 H 22 O 11, which in aqueous solution exist in the form of hydrates of their molecules.

Many substances are known (in general view MA), which, after dissolving in water and forming hydrates of MA nH 2 O molecules, undergo significant formula changes. As a result, hydrated ions appear in the solution - cations M + nH 2 O and anions A nH 2 O:

Such substances are classified as electrolytes.

The process of appearance of hydrated ions in an aqueous solution called electrolytic dissociation (S. Arrhenius, 1887).

Electrolytic dissociation ionic crystalline substances(M +)(A -) in water is irreversible reaction:

Such substances belong to strong electrolytes these include many bases and salts, for example:

Electrolytic dissociation of MA substances consisting of polar covalent molecules, is reversible reaction:

Such substances are classified as weak electrolytes; they include many acids and some bases, for example:

In dilute aqueous solutions of weak electrolytes, we will always find both the original molecules and the products of their dissociation - hydrated ions.

The quantitative characteristic of electrolyte dissociation is called degree of dissociation and is indicated? , Always? > 0.

For strong electrolytes? = 1 by definition (the dissociation of such electrolytes is complete).

For weak of electrolytes, the degree of dissociation is the ratio of the molar concentration of the dissociated substance (c d) to the total concentration of the substance in solution (c):

The degree of dissociation is a fraction of unity or 100%. For weak electrolytes? « From 1 (100%).

For weak acids H n And the degree of dissociation at each next step decreases sharply compared to the previous one:

The degree of dissociation depends on the nature and concentration of the electrolyte, as well as on the temperature of the solution; it grows with decrease concentration of the substance in the solution (i.e. when the solution is diluted) and when heating.

IN diluted solutions strong acids H n A their hydroanions H n-1 A do not exist, for example:

B concentrated In solutions, the content of hydroanions (and even the original molecules) becomes noticeable:

(it is impossible to summarize the equations for the stages of reversible dissociation!). When heating the values? 1 and? 2 increase, which promotes the occurrence of reactions involving concentrated acids.

Acids are electrolytes that, upon dissociation, supply hydrogen cations to an aqueous solution and do not form any other positive ions:

Common strong acids:

In a dilute aqueous solution (conditionally up to 10% or 0.1 molar) these acids dissociate completely. For strong acids H n A, the list includes their hydroanions(anions acid salts), also dissociating completely under these conditions.

Common weak acids:

Bases are electrolytes that, when dissociated, supply hydroxide ions and no other ions to an aqueous solution. negative ions do not form:

Dissociation sparingly soluble bases Mg(OH) 2, Cu(OH) 2, Mn(OH) 2, Fe(OH) 2 and others practical significance does not have.

TO strong reasons ( alkalis) include NaOH, KOH, Ba(OH) 2 and some others. The most famous weak base is ammonia hydrate NH 3 H 2 O.

Medium salts are electrolytes that, upon dissociation, supply any cations except H + and any anions except OH - into an aqueous solution:

We are talking only about highly soluble salts. Dissociation sparingly soluble and practically insoluble salts don't matter.

Dissociate similarly double salts:

Acid salts(most of them are soluble in water) dissociate completely according to the type of medium salts:

The resulting hydroanions are, in turn, exposed to water:

a) if the hydroanion belongs to strong acid, then it itself also dissociates completely:

And complete equation dissociation will be written as:

(solutions of such salts will necessarily be acidic, as well as solutions of the corresponding acids);

b) if the hydroanion belongs to weak acid, then its behavior in water is dual – either incomplete dissociation like a weak acid:

or interaction with water (called reversible hydrolysis):

At? 1 > ? 2 dissociation predominates (and the salt solution will be acidic), and at? 1 > ? 2 – hydrolysis (and the salt solution will be alkaline). Thus, solutions of salts with the anions HSO 3 -, H 2 PO 4 -, H 2 AsO 4 - and HSeO 3 - will be acidic, solutions of salts with other anions (the majority of them) will be alkaline. In other words, the name “acidic” for salts with a majority of hydroanions does not imply that these anions will behave like acids in solution (hydrolysis of hydroanions and the calculation of the ratio between α1 and α2 are studied only in high school).

Basic salts MgCl(OH), Cu 2 CO 3 (OH) 2 and others are mostly practically insoluble in water, and it is impossible to discuss their behavior in an aqueous solution.

13.3. Dissociation of water. Solution medium

The water itself is very weak electrolyte:

Concentrations of H + cation and OH - anion in clean water very small and amount to 1 10 -7 mol/l at 25 °C.

The hydrogen cation H + is the simplest nucleus - a proton p +(the electron shell of the H + cation is empty, 1s 0). A free proton has high mobility and penetrating ability; surrounded by polar H 2 O molecules, it cannot remain free. The proton immediately attaches to the water molecule:

In what follows, for simplicity, the notation H + is retained (but H 3 O + is implied).

Types aqueous solution environments:

For water at room temperature we have:

therefore, in clean water:

This equality is also true for aqueous solutions:

The practical pH scale corresponds to the range 1-13 (dilute solutions of acids and bases):

In a practically neutral environment with pH = 6–7 and pH = 7–8, the concentration of H + and OH - is very small (1 10 -6 – 1 10 -7 mol/l) and is almost equal to the concentration of these ions in pure water. Such solutions of acids and bases are considered extremely diluted (contain very little substance).

To practically establish the type of medium of aqueous solutions, use indicators– substances that give a characteristic color to neutral, acidic and/or alkaline solutions.

Common indicators in the laboratory are litmus, methyl orange and phenolphthalein.

Methyl orange (an indicator of an acidic environment) becomes pink in a strongly acidic solution (Table 16), phenolphthalein (an indicator for an alkaline environment) - crimson in a strongly alkaline solution, and litmus is used in all environments.

13.4. Ion exchange reactions

In dilute solutions of electrolytes (acids, bases, salts), chemical reactions usually occur with the participation ions. In this case, all elements of the reagents can retain their oxidation states ( exchange reactions) or change them ( redox reactions). The examples given below relate to exchange reactions (for the occurrence of redox reactions, see Section 14).

In accordance with Berthollet's ruleionic reactions proceed practically irreversibly if solid, slightly soluble substances are formed(they precipitate) highly volatile substances(they are released as gases) or soluble substances – weak electrolytes(including water). Ionic reactions are represented by a system of equations - molecular, complete And short ionic. The complete ionic equations are omitted below (the reader is encouraged to compose them himself).

When writing equations ionic reactions It is imperative to follow the solubility table (see Table 8).

Examples reactions with precipitation:

Attention! The slightly soluble (“m”) and practically insoluble (“n”) salts indicated in the solubility table (see Table 15) precipitate exactly as they are presented in the table (CaF 2 v, PbI 2 v, Ag 2 SO 4 v, AlPO 4 v, etc.).

In table 15 not specified carbonates– medium salts with the CO 3 2- anion. Please keep in mind that:

1) K 2 CO 3, (NH 4) 2 CO 3 and Na 2 CO 3 are soluble in water;

2) Ag 2 CO 3, BaCO 3 and CaCO 3 are practically insoluble in water and precipitate as such, for example:

3) salts of other cations, such as MgCO 3, CuCO 3, FeCO 3, ZnCO 3 and others, although insoluble in water, do not precipitate from an aqueous solution during ionic reactions (i.e., they cannot be obtained in this way).

For example, iron (II) carbonate FeCO 3, obtained “dry” or taken in the form of a mineral siderite, when added to water, it precipitates without visible interaction. However, when you try to obtain it by an exchange reaction in a solution between FeSO 4 and K 2 CO 3, a precipitate of the main salt precipitates (the conditional composition is given, in practice the composition is more complex) and carbon dioxide is released:

Similar to FeCO 3, sulfide chromium (III) Cr 2 S 3 (insoluble in water) does not precipitate from solution:

In table 15 also does not indicate salts that decompose water - sulfide aluminum Al 2 S 3 (as well as BeS) and acetate chromium (III) Cr(CH 3 COO) 3:

Consequently, these salts also cannot be obtained by an exchange reaction in solution:

(in the latter reaction the composition of the precipitate is more complex; such reactions are studied in more detail in higher education).

Examples reactions with gas release:

Examples reactions with the formation of weak electrolytes:

If the reagents and products of the exchange reaction are not strong electrolytes, ionic species there is no equation, for example:

13.5. Hydrolysis of salts

Hydrolysis of a salt is the interaction of its ions with water, leading to the appearance of acidic or alkaline environment, but not accompanied by the formation of sediment or gas (below we're talking about about medium salts).

The hydrolysis process occurs only with the participation soluble salts and consists of two stages:

1) dissociation salts in solution - irreversible reaction (degree of dissociation? = 1, or 100%);

2) actually hydrolysis, i.e. the interaction of salt ions with water, – reversible reaction (degree of hydrolysis?< 1, или 100 %).

Equations of the 1st and 2nd stages - the first of them is irreversible, the second is reversible - you cannot add them!

Note that salts formed by cations alkalis and anions strong acids do not undergo hydrolysis; they only dissociate when dissolved in water. In solutions of salts KCl, NaNO 3, Na 2 SO 4 and BaI 2 the environment neutral.

In case of interaction anion hydrolysis of the salt at the anion.

Dissociation of the KNO 2 salt occurs completely, hydrolysis of the NO 2 anion occurs to a very small extent (for a 0.1 M solution - by 0.0014%), but this is enough for the solution to become alkaline(among the hydrolysis products there is an OH - ion), it has a pH = 8.14.

Anions undergo hydrolysis only weak acids (in in this example– nitrite ion NO 2 - corresponding to weak nitrous acid HNO 2). The anion of a weak acid attracts the hydrogen cation present in water and forms a molecule of this acid, while the hydroxide ion remains free:

List of hydrolyzable anions:

Please note that in examples (c – e) you cannot increase the number of water molecules and instead of hydroanions (HCO 3 -, HPO 4 2-, HS -) write the formulas of the corresponding acids (H 2 CO 3, H 3 PO 4, H 2 S ). Hydrolysis – reversible reaction, and it cannot proceed “to the end” (until the formation of acid H n A).

If such an unstable acid as H 2 CO 3 were formed in a solution of its salt Na 2 CO 3, then CO 2 gas would be released from the solution (H 2 CO 3 = CO 2 v + H 2 O). However, when soda is dissolved in water, a transparent solution is formed without gas evolution, which is evidence of incomplete hydrolysis of the CO| anion. with the appearance of only hydroanion in solution carbonic acid HCOg.

The degree of hydrolysis of a salt by anion depends on the degree of dissociation of the hydrolysis product - acid (HNO 2, HClO, HCN) or its hydroanion (HCO 3 -, HPO 4 2-, HS -); the weaker the acid, the higher the degree of hydrolysis. For example, CO 3 2-, PO 4 3- and S 2- ions undergo hydrolysis into to a greater extent(in 0.1 M solutions ~ 5%, 37% and 58%, respectively) than the NO 2 ion, since the dissociation of H 2 CO 3 and H 2 S is in the 2nd step, and H 3 PO 4 is in the 3rd step steps (i.e., the dissociation of HCO 3 -, HS - and HPO 4 2- ions) occurs significantly less than the dissociation of the acid HNO 2. Therefore, solutions, for example, Na 2 CO 3, K 3 PO 4 and BaS will be highly alkaline(which is easy to verify by the soapiness of the soda solution to the touch). Excess OH ions in solution can be easily detected with an indicator or measured special devices(pH meters).

If in concentrated solution a salt that is strongly hydrolyzed by the anion, for example Na 2 CO 3, add aluminum, then the latter (due to amphotericity) will react with OH -

and hydrogen evolution will be observed. This - additional proof hydrolysis of the CO 3 2- ion occurs (after all, we did not add alkali NaOH to the Na 2 CO 3 solution!).

In case of interaction cation dissolved salt with water the process is called hydrolysis of salt by cation:

The dissociation of the Ni(NO 3) 2 salt occurs completely, the hydrolysis of the Ni 2+ cation occurs to a very small extent (for a 0.1 M solution - by 0.001%), but this is enough for the solution to become sour(H+ ion is present among the hydrolysis products), pH = 5.96.

Only poorly soluble basic and amphoteric hydroxides and ammonium cation NH 4 +. The hydrolyzed cation attracts the OH - anion present in water and forms the corresponding hydroxocation, while the H + cation remains free:

The ammonium cation in this case forms weak foundation– ammonia hydrate:

List of hydrolyzable cations:

Examples:

Please note that in examples (a – c) you cannot increase the number of water molecules and instead of the hydroxocations FeOH 2+, CrOH 2+, ZnOH + write the formulas of the hydroxides FeO(OH), Cr(OH) 3, Zn(OH) 2. If hydroxides were formed, then precipitation would form from solutions of FeCl 3 , Cr 2 (SO 4) 3 and ZnBr 2 salts, which is not observed (these salts form transparent solutions).

Excess H+ cations can be easily detected with an indicator or measured with special devices. You can also

do such an experiment. In a concentrated solution of a salt that is strongly hydrolyzed by the cation, for example AlCl 3:

magnesium or zinc is added. The latter will react with H +:

and hydrogen evolution will be observed. This experiment is additional evidence of the hydrolysis of the Al 3+ cation (after all, we did not add acid to the AlCl 3 solution!).

Examples of tasks for parts A, B

1. Strong electrolyte- This

1) C 6 H 5 OH

2) CH 3 COOH

3) C 2 H 4 (OH) 2

2. Weak electrolyte- This

1) hydrogen iodide

2) hydrogen fluoride

3) ammonium sulfate

4) barium hydroxide

3. In an aqueous solution, every 100 molecules form 100 hydrogen cations for acid

1) coal

2) nitrogenous

3) nitrogen

4-7. In the equation of dissociation of a weak acid in all possible steps

the sum of the coefficients is equal

8-11. For dissociation equations in a solution of two alkalis set

8. NaOH, Ba(OH) 2

9. Sr(OH) 2, Ca(OH) 2

10. KOH, LiOH

11. CsOH, Ca(OH) 2

the total sum of the coefficients is

12.V lime water contains a set of particles

1) CaOH+, Ca 2+, OH -

2) Ca 2+, OH -, H 2 O

3) Ca 2+, H 2 O, O 2-

4) CaOH +, O 2-, H+

13-16. When dissociating one formula unit of salt

14. K 2 Cr 2 O 7

16. Cr 2 (SO 4) 3

the number of ions formed is equal to

17. Greatest the amount of PO 4 -3 ion can be detected in a solution containing 0.1 mol

18. A reaction with precipitation is

1) MgSO 4 + H 2 SO 4 >...

2) AgF + HNO 3 >...

3) Na 2 HPO 4 + NaOH >...

4) Na 2 SiO 3 + HCl >...

19. A reaction with the release of gas is

1) NaOH + CH 3 COOH >...

2) FeSO 4 + KOH >...

3) NaHCO 3 + HBr >…

4) Pl(NO 3) 2 + Na 2 S >...

20. The short ionic equation OH - + H + = H 2 O corresponds to the interaction

1) Fe(OH) 2 + HCl >…

2) NaOH + HNO 2 >...

3) NaOH + HNO 3 >...

4) Ba(OH) 2 + KHSO 4 >...

21. B ionic equation reactions

SO 2 + 2ON = SO 3 2- + H 2 O

OH ion - may correspond to the reagent

4) C 6 H 5 OH

22-23. Ionic equation

22. ZCa 2+ + 2PO 4 3- = Ca 3 (PO 4) 2 v

23. Ca 2+ + HPO 4 2- = CaHPO 4 v

corresponds to the reaction between

1) Ca(OH) 2 and K 3 PO 4

2) CaCl 2 and NaH 2 PO 4

3) Ca(OH) 2 and H 3 PO 4

4) CaCl and K 2 HPO 4

24-27. In the molecular reaction equation

24. Na 3 PO 4 + AgNO 3 >...

25. Na 2 S + Cu(NO 3) 2 >…

26. Ca(HSO 3) 2 >…

27. K 2 SO 3 + 2HBr >... the sum of the coefficients is

28-29. For a complete neutralization reaction

28. Fe(OH) 2 + HI >…

29. Ba(OH) 2 + H 2 S >…

the sum of the coefficients in the complete ionic equation is

30-33. In the short ionic reaction equation

30. NaF + AlCl 3 >…

31. K 2 CO 3 + Sr(NO 3) 2 >...

32. Mgl 2 + K 3 PO 4 >...

33. Na 2 S + H 2 SO 4 >...

the sum of the coefficients is equal

34-36. In an aqueous salt solution

34. Ca(ClO 4) 2

36. Fe 2 (SO 4) 3

environment is formed

1) acidic

2) neutral

3) alkaline

37. The concentration of hydroxide ion increases after salt is dissolved in water

38. Neutral environment will be in the final solution after mixing the solutions of the original salts in sets

1) BaCl 2, Fe(NO 3) 3

2) Na 2 CO 3, SrS

4) MgCl 2, RbNO 3

39. Match salt with its ability to hydrolyze.

40. Match the salt with the solution medium.

41. Establish a correspondence between salt and the concentration of hydrogen cation after dissolving the salt in water.

Dissolution is a spontaneous, reversible physicochemical process that includes three main stages.

The atomization stage is the destruction of the crystal lattice of the substance being dissolved; the process is endothermic (D at H>O).

2) Stage of solvation (hydration) - the formation of solvation (hydrate) shells around the particles of the dissolved substance; exothermic process, (D sol H<О).

3) Diffusion stage - uniform distribution of the dissolved substance throughout the entire volume of the solution (D dif H ≈ O).

Thus, the heat of solution (D р Н) is an integral value:

D p H = D at H + D sol H + D diff H

Heat of solution is the thermal effect of dissolving 1 mole of a substance in an infinitely large volume of solvent.

The dissolution of most solids in water is an endothermic process (D p H > 0), because the heat absorbed at the atomization stage is not compensated by the heat released at the solvation stage. When gases dissolve, heat is released (D p H< 0), т.к. их растворение не включает стадию атомизация (газообразные вещества не образуют кристаллических решеток). Растворение жидкостей друг в друге протекает без заметного теплового эффекта (D p H ≈ 0), т.к. главной стадией их растворения является диффузия.

Like any reversible process, dissolution reaches equilibrium. A solution that is in equilibrium with an excess of solute is called saturated. At equilibrium, the rate of dissolution is equal to the rate of crystallization.

According to the degree of saturation, solutions are:

unsaturated: contain less solute than saturated ones,

rich,

oversaturated: contain more solute than saturated ones (they are unstable).

4.3. Solubility of gases, liquids and solids in water

Solubility (S) is the ability of a substance to dissolve in a given solvent. It is equal to the content of the solute in its saturated solution at a given temperature.

Solubility depends on the nature of the substances and the thermodynamic parameters of the system. The influence of the nature of substances on solubility is described by the rule: “ Like dissolves into like" In other words, polar substances dissolve well in polar solvents, and non-polar substances dissolve well in non-polar ones. For example: salt NaCl is highly soluble in water and poorly soluble in benzene; I 2 is highly soluble in benzene and poorly soluble in water.

The dissolution of gases in water can be represented by the diagram:

A (gas) + H 2 OA (solution), D р Н<О

In accordance with Le Chatelier's principle, as the temperature increases, the equilibrium shifts to the left, i.e. solubility decreases, and with decreasing temperature - to the right, solubility increases (Table 3).

Table 3 - Solubility of gases (l/1l H 2 O) at p = 1 atm.

In accordance with Le Chatelier's principle, as pressure increases, equilibrium shifts to the right, i.e. the solubility of gases increases. The quantitative dependence of gas solubility on pressure is described by Henry’s equation (1803):

where k is Henry's constant,

p - gas pressure above the solution.

Henry's law allows us to reveal the causes of decompression sickness. It occurs in divers, pilots and representatives of other professions who, due to their occupation, quickly move from a high-pressure environment to a low-pressure environment.

During a person’s stay in an environment with high pressure, his blood and tissues are saturated with nitrogen (N 2) and partially carbon dioxide (CO 2). There is no accumulation of oxygen, as it is spent on physiological processes in the body. When a person quickly moves into a low-pressure environment, excess amounts of dissolved gases are released, which do not have time to diffuse through the lungs and form gas plugs in the tissues and blood vessels. This leads to blockage and rupture of blood capillaries, accumulation of gas bubbles in the subcutaneous fatty tissue, in the joints, and in the bone marrow. Symptoms of decompression sickness include dizziness, itching, muscle and chest pain, respiratory failure, paralysis and death.

The solubility of gases is affected by the presence of electrolytes in solution. This dependence is described by the Sechenov equation (1859):

where S and S o are the solubility of the gas in an electrolyte solution and pure water,

c - electrolyte concentration,

k is the Sechenov constant.

From Sechenov’s equation it follows that the higher the electrolyte concentration in the solution, the lower the solubility of gases. This is why the solubility of gases in water is greater than in plasma (Table 4).

Table 4 - Solubility of gases in pure water and blood plasma at 38ºС

The dissolution of a liquid in water can be represented by the diagram:

A (g) + H 2 OA (solution)

The main stage of dissolution of a liquid in a liquid is diffusion, the rate of which increases with increasing temperature. Accordingly, the mutual solubility of liquids increases with increasing temperature.

There are three types of liquids:

a) unlimitedly soluble in each other: H 2 SO 4 / H 2 O, C 2 H 5 OH / H 2 O;

b) sparingly soluble: C 6 H 6 / H 2 O

c) absolutely insoluble: Hg / H 2 O.

If a third component is added to a system of two immiscible liquids, then the ratio of its concentrations in each liquid is a constant value at a given temperature (Nernst-Shilov distribution law) (Figure 6).

Drawing6 - Nernst-Shilov distribution law

The Nernst-Shilov law is the theoretical basis of extraction, one of the methods for separating mixtures.

The dissolution of solids in water is described by the following scheme:

A (k) + H 2 OA (solution), Dр Н > O

If a sparingly soluble electrolyte (salt, base or acid) dissolves, then the heterogeneous equilibrium between the solid and its ions in a saturated solution can be represented by the diagram:

A n B m (k) nA m+ (aq) + mB n- (aq).

This equilibrium is characterized using the solubility constant Ks, which is a heterogeneous equilibrium constant:

K s = n m

For binary electrolytes n = m= 1, therefore

K s = · .

Accordingly, S 2 =K s, and S =

For example, when the sparingly soluble salt BaSO 4 is dissolved in water, a heterogeneous equilibrium is established between the crystals of the substance and its ions in a saturated solution:

BaSO 4 (k) Ba 2+ (aq) + SO 4 2- (aq)

According to the law of mass action, K S = = 1.1·10 -10.

Hence S =
.

The lower Ks, the lower the solubility of the substance and the easier it is for a precipitate of sparingly soluble electrolyte to form.

The condition for the formation of a precipitate of sparingly soluble electrolyte can be formulated as follows: Precipitates form from saturated and supersaturated solutions. In a saturated solution · = K s , and in a supersaturated solution · > K s

One of the most important heterogeneous processes in vivo is the formation of bone tissue. The main mineral component of bone tissue is calcium hydroxyphosphate (hydroxyapatite) Ca 5 (RO 4 ) 3 HE.

The process of bone tissue formation can be represented as follows. In the blood at pH = 7.4, the anions HPO 4 2– and H 2 PO 4 –, as well as Ca 2+ cations, are present in approximately equal quantities. After comparing the solubility constants of CaHPO 4 (K S = 2.7∙10 –7) and Ca(H 2 PO 4) 2 (K S = 1∙10 –3), it becomes obvious that the CaHPO 4 salt is less soluble. As a result, it is CaHPO 4 that is formed at the first stage of bone tissue formation:

Ca 2+ + NPO 4 2– CaHPO 4 .

Further formation of hydroxoapatite proceeds in accordance with the equations:

3 CaHPO 4 + Ca 2+ + 2 OH – Ca 4 H(PO 4) 3 + 2 H 2 O,

Ca 4 H (PO 4) 3 + Ca 2+ + 2 OH – Ca 5 (PO 4) 3 OH + H 2 O.

The solubility constant of hydroxoapatite is very small (K S = 10 -58), which indicates the high stability of bone tissue.

With an excess of Ca 2+ ions in the blood, the balance shifts to the right, which leads to calcification of the bones. With a lack of Ca 2+, the equilibrium shifts to the left; bone tissue is destroyed. In children this leads to rickets, develops in adults osteoporosis.

If there is a lack of calcium in bone tissue, its place can be taken by the closest electronic analogues: beryllium and strontium. Their accumulation causes accordingly beryllium and strontium rickets(increased fragility and fragility of bones). When the radioisotope Sr-90 is incorporated into bone tissue, the bone marrow is irradiated, which can lead to leukemia and other cancers. Calcium blocks the body's accumulation of radioactive strontium.

A solution is a homogeneous system consisting of two or more substances, the content of which can be changed within certain limits without disturbing the homogeneity.

Water solutions consist of water(solvent) and dissolved substance. The state of substances in an aqueous solution is, if necessary, indicated by a subscript (p), for example, KNO 3 in solution - KNO 3 (p).

Solutions that contain a small amount of solute are often called diluted and solutions with a high solute content - concentrated. A solution in which further dissolution of a substance is possible is called unsaturated and a solution in which a substance ceases to dissolve under given conditions is saturated. The latter solution is always in contact (in heterogeneous equilibrium) with an undissolved substance (one crystal or more).

Under special conditions, for example when carefully (without stirring) cooling a hot unsaturated solution solid substances that can form oversaturated solution. When a crystal of a substance is introduced, such a solution is divided into a saturated solution and a precipitate of the substance.

In accordance with chemical theory of solutions D.I. Mendeleev, the dissolution of a substance in water is accompanied, firstly, by destruction chemical bonds between molecules (intermolecular bonds in covalent substances) or between ions (in ionic substances), and thus the particles of the substance mix with water (in which some of the hydrogen bonds between molecules are also destroyed). The breaking of chemical bonds occurs due to the thermal energy of movement of water molecules, and this occurs cost energy in the form of heat.

Secondly, once in water, particles (molecules or ions) of the substance are subjected to hydration. As a result, hydrates– compounds of uncertain composition between particles of a substance and water molecules (the internal composition of the particles of the substance themselves does not change upon dissolution). This process is accompanied highlighting energy in the form of heat due to the formation of new chemical bonds in hydrates.

In general, the solution is either cools down(if the heat consumption exceeds its release), or heats up (otherwise); sometimes - if the heat input and its release are equal - the temperature of the solution remains unchanged.

Many hydrates turn out to be so stable that they do not collapse even when the solution is completely evaporated. Thus, solid crystalline hydrates of the salts CuSO 4 5H 2 O, Na 2 CO 3 10H 2 O, KAl(SO 4) 2 12H 2 O, etc. are known.

The content of a substance in a saturated solution at T= const quantitatively characterizes solubility of this substance. Solubility is usually expressed as the mass of solute per 100 g of water, for example 65.2 g KBr/100 g H 2 O at 20 °C. Therefore, if 70 g of solid potassium bromide is added to 100 g of water at 20 °C, then 65.2 g of salt will go into solution (which will be saturated), and 4.8 g of solid KBr (excess) will remain at the bottom of the glass.

It should be remembered that the solute content in rich solution equals, V unsaturated solution less and in oversaturated solution more its solubility at a given temperature. Thus, a solution prepared at 20 °C from 100 g of water and sodium sulfate Na 2 SO 4 (solubility 19.2 g/100 g H 2 O), containing

15.7 g salt – unsaturated;

19.2 g salt – saturated;

2O.3 g of salt – supersaturated.

The solubility of solid substances (Table 14) usually increases with increasing temperature (KBr, NaCl), and only for some substances (CaSO 4, Li 2 CO 3) the opposite is observed.

The solubility of gases decreases with increasing temperature, and increases with increasing pressure; for example, at a pressure of 1 atm, the solubility of ammonia is 52.6 (20 °C) and 15.4 g/100 g H 2 O (80 °C), and at 20 °C and 9 atm it is 93.5 g/100 g H 2 O.

In accordance with solubility values, substances are distinguished:

– highly soluble, the mass of which in a saturated solution is comparable to the mass of water (for example, KBr - at 20 °C solubility 65.2 g/100 g H 2 O; 4.6 M solution), they form saturated solutions with a molarity of more than 0.1 M;

– slightly soluble, whose mass in a saturated solution is significantly less than the mass of water (for example, CaSO 4 - at 20 °C solubility 0.206 g/100 g H 2 O; 0.015 M solution), they form saturated solutions with a molarity of 0.1–0.001 M;

– practically insoluble, whose mass in a saturated solution is negligible compared to the mass of the solvent (for example, AgCl - at 20 °C solubility 0.00019 g per 100 g H 2 O; 0.0000134 M solution), they form saturated solutions with a molarity of less than 0.001 M.

Compiled based on reference data solubility table common acids, bases and salts (Table 15), which indicates the type of solubility, substances unknown to science (not obtained) or completely decomposed by water are noted.

Conventions used in the table:

“r” – highly soluble substance

“m” – slightly soluble substance

“n” – practically insoluble substance

“–” – substance not received (does not exist)

» – the substance mixes with water unlimitedly

Note. This table corresponds to the preparation of a saturated solution at room temperature by adding the substance (in the appropriate state of aggregation) to water. It should be taken into account that obtaining precipitation of poorly soluble substances using ion exchange reactions is not always possible (for more details, see 13.4).

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Each orbital can accommodate no more than two electrons.
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Molecules. Chemical bond. Structure of substances
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Manganese
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Elements with metallic properties are located in the IA – VIA groups of the Periodic Table (Table 7).

Hydrogen
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Chlorine. Hydrogen chloride
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Chlorides
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Oxygen
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Nitrites. Nitrates
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Free carbon
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