Useful salt insoluble in water x. One of the possible solutions. Chemical properties of acid salts

A salt can be defined as a compound that is formed by the reaction between an acid and a base, but is not water. In this section, those properties of salts that are associated with ionic equilibria will be considered.

REACTIONS OF SALT IN WATER

Somewhat later it will be shown that solubility is a relative concept. However, for the purposes of the following discussion, we can roughly classify all salts into water-soluble and water-insoluble salts.

Some salts, when dissolved in water, form neutral solutions. Other salts form acidic or alkaline solutions. This is due to the occurrence of a reversible reaction between salt ions and water, as a result of which conjugate acids or bases are formed. Whether a salt solution is neutral, acidic, or alkaline depends on the type of salt. In this sense, there are four types of salts.

Salts formed by strong acids and weak bases. Salts of this type, when dissolved in water, form an acidic solution. As an example, consider ammonium chloride. When this salt is dissolved in water, the ammonium ion acts as an acid, donating a proton to the water.

The excess amount of ions formed in this process determines the acidic properties of the solution.

Salts formed by a weak acid and a strong base. Salts of this type, when dissolved in water, form an alkaline solution. Let's take sodium acetate as an example. The acetate ion acts as a base by accepting a proton from water, which in this case acts as an acid:

An excess of OH ions formed in this process determines the alkaline properties of the solution.

Salts formed by strong acids and strong bases. When salts of this type are dissolved in water, a neutral solution is formed. Let us take sodium chloride as an example. When dissolved in water, this salt is completely ionized, and, therefore, the concentration of ions is equal to the concentration of ions. Since neither ion enters into acid-base reactions with water, in

the solution does not form an excess of ions or OH Therefore, the solution is neutral.

Salts formed by weak acids and weak bases. An example of salts of this type is ammonium acetate. When dissolved in water, the ammonium ion reacts with water as an acid, and the acetate ion reacts with water as a base. Both of these reactions are described above. An aqueous solution of a salt formed by a weak acid and a weak base can be weakly acidic, slightly alkaline, or neutral, depending on the relative concentrations of ions formed as a result of the reactions of cations and anions of the salt with water. It depends on the ratio between the values ​​of the dissociation constants of the cation and anion.

The solubility table of salts, acids and bases is the foundation, without which it is impossible to fully master chemical knowledge. The solubility of bases and salts helps in teaching not only schoolchildren, but also professional people. The creation of many life products cannot do without this knowledge.

Table of solubility of acids, salts and bases in water

The table of solubility of salts and bases in water is a manual that helps in mastering the basics of chemistry. The following notes will help you understand the table below.

• P - indicates a soluble substance;
• H is an insoluble substance;
• M - the substance is slightly soluble in the aquatic environment;
• RK - a substance can dissolve only when exposed to strong organic acids;
• The dash will say that such a creature does not exist in nature;
• NK - does not dissolve in either acids or water;
• ? - a question mark indicates that today there is no exact information about the dissolution of the substance.

Often, the table is used by chemists and schoolchildren, students for laboratory research, during which it is necessary to establish the conditions for the occurrence of certain reactions. According to the table, it turns out to find out how the substance behaves in a hydrochloric or acidic environment, whether a precipitate is possible. Precipitate during research and experiments indicates the irreversibility of the reaction. This is a significant point that can affect the course of the entire laboratory work.

Task 1. "Useful salt"

Salt X, which is insoluble in water, is a component of many useful substances - white paints, refractory materials, drilling fluids, contrast agents for radiography. It consists of three elements, one of which is sulfur. When calcined with an excess of carbon, X turns into a soluble salt Y, which consists of only two elements in equal amounts. The masses of the elements in Y differ by a factor of 4.28.

1. Determine the formulas for salts X and Y.
2. Write the equations for the reactions X → Y and Y → X.
3. Suggest three ways to obtain X from substances belonging to different classes of compounds.

Decision

1. When calcined with coal, salt X loses oxygen, and sulfur and the metal element remain in equal proportions, i.e. Y is divalent metal sulfide, MeS.

From the mass ratio we find:

M(Me) \u003d 32 4.28 \u003d 137 g / mol - this is barium. Y—BaS, X—BaSO4.

4 points(on 2 points for each salt).

The answer X - BaSO 3 is also considered correct.

2. X → Y. BaSO 4 + 4C = BaS + 4CO

1.5 points

(the equation BaSO 4 + 2C = BaS + 2CO 2 and similar equations with BaSO 3 are also accepted),

Y → X. BaS + H 2 SO4 = BaSO 4 + H2S

1.5 points

3. BaO + H 2 SO 4 \u003d BaSO 4 + H 2 O

Ba(OH) 2 + H 2 SO 4 = BaSO 4 + 2H 2 O

BaCl 2 + H 2 SO 4 \u003d BaSO 4 ↓ + 2HCl

(any reasonable reactions of BaSO 3 formation are also accepted)

Each equation is 1 point, maximum - 3 points.

Total - 10 points

Problem 2. "Incomplete reaction equations"

Below are the equations of chemical reactions in which some substances and coefficients are omitted. Fill in all gaps.

1. … + Br2 = S + 2…
2. 2NaCl + 2… = …NaOH + … + Cl 2
3. … + 5O 2 = 3CO 2 + …H 2 O
4. Pb 3 O 4 + 4 ... = ... + 2Pb (NO 3) 2 + ... H 2 O
5. ...NaHCO 3 \u003d Na 2 CO 3 + ... + H 2 O

Decision

1. H 2 S+ Br2 = S + 2 HBr or Na 2 S+ Br2 = S + 2 NaBr
2. 2NaCl + 2 H2O = 2 NaOH + H2+Cl2
3. C 3 H 8+ 5O 2 = 3CO 2 + 4 H2O
4. Pb3O4+ 4HNO 3 = PbO2+ 2Pb(NO 3) 2 + 2 H2O
5. 2 NaHCO 3 \u003d Na 2 CO 3 + CO2+ H2O

For each correct equation, 2 points.

Total - 10 points

Task 3. "Experiments with chips"

Calcium chips weighing 4.0 g were calcined in air and then thrown into water. When the chips were dissolved in water, 560 ml of gas (n.a.) was released, which is practically insoluble in water.

1. Write down the reaction equations.
2. Determine by how many grams the weight of the chips increased during calcination.
3. Calculate the composition of the calcined chips in mass percent.

Decision

1. When calcium chips are calcined, the reaction occurs: 2Ca + O 2 \u003d 2CaO

(The condition that the gas is practically insoluble in water excludes the reaction of calcium with nitrogen, which can lead to calcium nitride hydrolyzing to form NH 3 .)

Since calcium melts at a high temperature, and the reaction product is also refractory, the oxidation of the metal initially occurs only from the surface.

The calcined chips are metal coated on the outside with a layer of oxide. When placed in water, both the metal and the oxide react with it:

• CaO + H 2 O \u003d Ca (OH) 2;
• Ca + 2H 2 O \u003d Ca (OH) 2 + H 2.

2. The amount of metal substance that did not react with oxygen is equal to the amount of substance of the released gas (hydrogen): n (Ca) \u003d n (H 2) \u003d 0.56 / 22.4 \u003d 0.025 mol.

In total, in the original chip n(Ca) = 4/40 = 0.1 mol.

Thus, 0.1 - 0.025 = 0.075 mol of calcium reacted with oxygen, which is m(Ca) = 0.075∙40 = 3 g.

The increase in the mass of chips is associated with the addition of oxygen. The mass of oxygen that reacted with calcium is m (O 2) \u003d 32 ∙ 0.0375 \u003d 1.2 g.

So, the mass of chips after calcination increased by 1.2 g.

3. The calcined chips consist of calcium (0.025 mol) weighing 1 g and calcium oxide (0.075 mol) weighing 4.2 g. Composition in mass percent: Ca - 19.2%; CaO - 80.8%.

Grading system:

Task 4. "Unknown salt"

The unknown salt is formed by two ions with the electronic configuration of argon. It is known that when it is introduced into an aqueous solution of silver nitrate, a precipitate forms, when it is treated with hydrochloric acid, gas is released, and an aqueous solution of sodium carbonate does not cause any changes.

1. Name the salt. Write down the electronic configuration of the ions that make up the salt.
2. Write the equations of the described reactions in molecular and abbreviated ionic form.
3. Suggest two ways to obtain this salt. Write down the reaction equations.

Decision

1. Ions with the configuration of the inert gas argon 1s 2 2s 2 2p 6 3s 2 3p 6 are the cations of the beginning of the fourth period (for example, K + , Ca 2+) and the anions of the end of the third period (for example, S 2– , Cl –). Only potassium sulfide K 2 S satisfies the conditions described in the problem.

2. Reaction equations:

• K 2 S + 2AgNO 3 \u003d Ag2S ↓ + 2KNO3
• 2Ag + + S 2– = Ag 2 S↓
• K 2 S + 2HCl \u003d 2KCl + H 2 S
• 2H + + S 2– = H 2 S

3. Salt can be obtained in different ways, for example, by the interaction of simple substances, the interaction of potassium hydroxide with hydrogen sulfide:

• 2K + S = K 2 S;
• 2KOH + H 2 S = K 2 S + 2H 2 O.

Grading system:

Task 5. "Unknown metal"

A piece of silver-white unknown metal was brought to the chemistry room.

The teacher instructed one of the students to do the analysis of the metal. Student compiled

research plan. When the atmospheric pressure became equal to 760 mm Hg. Art., the student cooled the installation to 0 ° C and proceeded to the analysis of the metal.

Taking an exact weight of the metal - 1.00 g, he dissolved it in hydrochloric acid. In this case, hydrogen was released with a volume of 2.49 liters. This was enough to identify the metal.

1. Based on the experimental data, determine the metal. Write the reaction equation.
2. Why is it important to consider atmospheric pressure and temperature in this study?
3. What additional reactions can confirm the metal identification?

Decision

1. The metal beryllium was determined and the reaction equation was written

5 points

One possible solution:

The amount of released hydrogen was determined

The metal reacts with hydrochloric acid according to the equation:

Me+ x HCl = MeCl x + 1/2 x H2

where: m- weight of metal sample, x- metal valence, n- amount of hydrogen substance. Of all the possible selection options for valence, beryllium is suitable. M = 9.09 g/mol

Be + 2HCl \u003d BeCl 2 + H 2

2. The dependence of gas volume on pressure and temperature is explained

2 points

3. Beryllium hydroxide has amphoteric properties. The reaction equation for the production of beryllium hydroxide and the reaction of Be(OH) 2 with acid and alkali is given

3 points

• BeCl 2 + 2NaOH = Be(OH)2↓ + 2NaCl
• Be(OH) 2 + 2HCl = BeCl 2 + 2H 2 O
• Be(OH) 2 + 2NaOH = Na 2

• Be(OH) 2 + 2NaOH = Na 2 BeO 2 + 2H 2 O

The student can offer different ways to identify beryllium, it is not required to prove the difference between beryllium and aluminum in this problem.

Total - 10 points

Task 6. "Gas that does not support combustion"

Granules of substance X were placed in the device shown in Figure 1 and liquid Y was poured. After the tap was opened, liquid Y descended from the funnel into the lower part of the device and came into contact with substance X, a reaction began, accompanied by the evolution of colorless gas Z. Gas Z was collected in a flask by the method of displacement of air (see. rice. 6.1).

A burning candle is introduced into a flask filled with gas Z (see Fig. rice. 6.2), and the candle went out. However, when the candle was taken out of the flask, it caught fire again.

1. What gas was obtained in the device shown in Figure 1? What is the name of this device?
2. What can substances X and Y be? Write an equation for a possible reaction between X and Y to form Z.
3. Explain the candle experiment. Why did the candle go out when it was brought into the flask, and flare up again when it was taken out of the flask? How long can this experience continue?
4. According to the safety regulations, before conducting an experiment with a candle, it is necessary to check the gas Z "for purity". What does it mean? How to do it? What can happen if this safety rule is neglected? Explain the answer.

Decision

1. Hydrogen (gas Z) was produced in a Kipp apparatus.

2 points

2. Substance X is an active metal, such as zinc; Y is an acid, such as hydrochloric or dilute sulfuric. Possible interaction option:

Zn + 2HCl \u003d ZnCl 2 + H 2

2 points

3. The candle goes out in a flask filled with hydrogen, since this gas does not support combustion. However, when a lit candle is introduced into the flask, hydrogen ignites at the opening of the flask. Hydrogen burns with a colorless flame, so it is almost invisible. When the candle is removed from the flask, the burning hydrogen ignites the wick and the candle flares up again.

This experiment can be continued (put the candle into the flask and take it out) until the hydrogen burns quietly in the flask. Gradually, as the hydrogen burns out, the combustion front will rise higher in the flask. Combustion will be more and more unstable due to the "mixing" of oxygen in the air.

3 points

4. Testing hydrogen "for purity" is an experimental test for the absence of gas impurities that form "explosive mixtures" with it, such as oxygen, air, chlorine. To test hydrogen for purity, it is collected in a test tube turned upside down and brought to the flame of an alcohol lamp. Pure hydrogen ignites with a slight “p” sound. "Dirty" hydrogen explodes with a loud pop or whistle.

If “dirty” hydrogen is collected in the flask for this experiment, then when a burning candle is introduced, the explosive mixture will explode.

3 points

Total - 10 points

Out of 6 tasks, 5 solutions are included in the final assessment, for which the participant
scores the highest, that is one of the tasks with the lowest score is not
taken into account.

Cations	Anions
F-	Cl-	br-	I-	S2-	NO 3 -	CO 3 2-	SiO 3 2-	SO 4 2-	PO 4 3-
Na+	R	R	R	R	R	R	R	R	R	R
K+	R	R	R	R	R	R	R	R	R	R
NH4+	R	R	R	R	R	R	R	R	R	R
Mg2+	RK	R	R	R	M	R	H	RK	R	RK
Ca2+	NK	R	R	R	M	R	H	RK	M	RK
Sr2+	NK	R	R	R	R	R	H	RK	RK	RK
Ba 2+	RK	R	R	R	R	R	H	RK	NK	RK
sn 2+	R	R	R	M	RK	R	H	H	R	H
Pb 2+	H	M	M	M	RK	R	H	H	H	H
Al 3+	M	R	R	R	G	R	G	NK	R	RK
Cr3+	R	R	R	R	G	R	G	H	R	RK
Mn2+	R	R	R	R	H	R	H	H	R	H
Fe2+	M	R	R	R	H	R	H	H	R	H
Fe3+	R	R	R	-	-	R	G	H	R	RK
Co2+	M	R	R	R	H	R	H	H	R	H
Ni2+	M	R	R	R	RK	R	H	H	R	H
Cu2+	M	R	R	-	H	R	G	H	R	H
Zn2+	M	R	R	R	RK	R	H	H	R	H
CD 2+	R	R	R	R	RK	R	H	H	R	H
Hg2+	R	R	M	NK	NK	R	H	H	R	H
Hg 2 2+	R	NK	NK	NK	RK	R	H	H	M	H
Ag+	R	NK	NK	NK	NK	R	H	H	M	H

Legend:

P - the substance is highly soluble in water; M - slightly soluble; H - practically insoluble in water, but easily soluble in weak or dilute acids; RK - insoluble in water and soluble only in strong inorganic acids; NK - insoluble neither in water nor in acids; G - completely hydrolyzes upon dissolution and does not exist in contact with water. A dash means that such a substance does not exist at all.

In aqueous solutions, salts completely or partially dissociate into ions. Salts of weak acids and/or weak bases undergo hydrolysis. Aqueous salt solutions contain hydrated ions, ion pairs, and more complex chemical forms, including hydrolysis products, etc. A number of salts are also soluble in alcohols, acetone, acid amides, and other organic solvents.

From aqueous solutions, salts can crystallize in the form of crystalline hydrates, from non-aqueous solutions - in the form of crystalline solvates, for example CaBr 2 3C 2 H 5 OH.

Data on various processes occurring in water-salt systems, on the solubility of salts in their joint presence depending on temperature, pressure and concentration, on the composition of solid and liquid phases can be obtained by studying the solubility diagrams of water-salt systems.

General methods for the synthesis of salts.

1. Obtaining medium salts:

1) metal with non-metal: 2Na + Cl 2 = 2NaCl

2) metal with acid: Zn + 2HCl = ZnCl 2 + H 2

3) metal with a salt solution of a less active metal Fe + CuSO 4 = FeSO 4 + Cu

4) basic oxide with acid oxide: MgO + CO 2 = MgCO 3

5) basic oxide with acid CuO + H 2 SO 4 \u003d CuSO 4 + H 2 O

6) bases with acidic oxide Ba (OH) 2 + CO 2 = BaCO 3 + H 2 O

7) bases with acid: Ca (OH) 2 + 2HCl \u003d CaCl 2 + 2H 2 O

8) acid salts: MgCO 3 + 2HCl = MgCl 2 + H 2 O + CO 2

BaCl 2 + H 2 SO 4 \u003d BaSO 4 + 2HCl

9) a base solution with a salt solution: Ba (OH) 2 + Na 2 SO 4 \u003d 2NaOH + BaSO 4

10) solutions of two salts 3CaCl 2 + 2Na 3 PO 4 = Ca 3 (PO 4) 2 + 6NaCl

2. Obtaining acid salts:

1. Interaction of an acid with a lack of a base. KOH + H 2 SO 4 \u003d KHSO 4 + H 2 O

2. Interaction of a base with an excess of acid oxide

Ca(OH) 2 + 2CO 2 = Ca(HCO 3) 2

3. Interaction of an average salt with acid Ca 3 (PO 4) 2 + 4H 3 PO 4 \u003d 3Ca (H 2 PO 4) 2

3. Obtaining basic salts:

1. Hydrolysis of salts formed by a weak base and a strong acid

ZnCl 2 + H 2 O \u003d Cl + HCl

2. Addition (drop by drop) of small amounts of alkalis to solutions of medium metal salts AlCl 3 + 2NaOH = Cl + 2NaCl

3. Interaction of salts of weak acids with medium salts

2MgCl 2 + 2Na 2 CO 3 + H 2 O \u003d 2 CO 3 + CO 2 + 4NaCl

4. Obtaining complex salts:

1. Reactions of salts with ligands: AgCl + 2NH 3 = Cl

FeCl 3 + 6KCN] = K 3 + 3KCl

5. Obtaining double salts:

1. Joint crystallization of two salts:

Cr 2 (SO 4) 3 + K 2 SO 4 + 24H 2 O \u003d 2 + NaCl

4. Redox reactions due to the properties of the cation or anion. 2KMnO 4 + 16HCl = 2MnCl 2 + 2KCl + 5Cl 2 + 8H 2 O

2. Chemical properties of acid salts:

1. Thermal decomposition with the formation of medium salt

Ca (HCO 3) 2 \u003d CaCO 3 + CO 2 + H 2 O

2. Interaction with alkali. Obtaining medium salt.

Ba(HCO 3) 2 + Ba(OH) 2 = 2BaCO 3 + 2H 2 O

3. Chemical properties of basic salts:

1. Thermal decomposition. 2 CO 3 \u003d 2CuO + CO 2 + H 2 O

2. Interaction with acid: the formation of an average salt.

Sn(OH)Cl + HCl = SnCl 2 + H 2 O

4. Chemical properties of complex salts:

1. Destruction of complexes due to the formation of poorly soluble compounds:

2Cl + K 2 S \u003d CuS + 2KCl + 4NH 3

2. Exchange of ligands between the outer and inner spheres.

K 2 + 6H 2 O \u003d Cl 2 + 2KCl

5. Chemical properties of double salts:

1. Interaction with alkali solutions: KCr(SO 4) 2 + 3KOH = Cr(OH) 3 + 2K 2 SO 4

2. Recovery: KCr (SO 4) 2 + 2H ° (Zn, diluted H 2 SO 4) \u003d 2CrSO 4 + H 2 SO 4 + K 2 SO 4

The raw materials for the industrial production of a number of chloride salts, sulfates, carbonates, Na, K, Ca, Mg borates are sea and ocean water, natural brines formed during its evaporation, and solid deposits of salts. For a group of minerals that form sedimentary salt deposits (sulfates and chlorides of Na, K and Mg), the code name “natural salts” is used. The largest deposits of potassium salts are located in Russia (Solikamsk), Canada and Germany, powerful deposits of phosphate ores - in North Africa, Russia and Kazakhstan, NaNO3 - in Chile.

Salts are used in food, chemical, metallurgical, glass, leather, textile industries, agriculture, medicine, etc.

The main types of salts

1. Borates (oxoborates), salts of boric acids: metaboric HBO 2, orthoboric H 3 BO 3 and polyboric acids not isolated in the free state. According to the number of boron atoms in the molecule, they are divided into mono-, di, tetra-, hexaborates, etc. Borates are also called according to the acids that form them and according to the number of moles of B 2 O 3 per 1 mole of the basic oxide. So various metaborates can be called monoborates if they contain an anion B (OH) 4 or a chain anion (BO 2) n n - diborates - if they contain a chain double anion (B 2 O 3 (OH) 2) n 2n- triborates - if they contain a ring anion (B 3 O 6) 3-.

The structures of borates include boron-oxygen groups - “blocks” containing from 1 to 6, and sometimes 9 boron atoms, for example:

The coordination number of boron atoms is 3 (boron-oxygen triangular groups) or 4 (tetrahedral groups). Boron-oxygen groups are the basis of not only island, but also more complex structures - chain, layered and framework polymerized. The latter are formed as a result of the elimination of water in the molecules of hydrated borates and the appearance of bridging bonds through oxygen atoms; the process is sometimes accompanied by the breaking of the B-O bond within the polyanions. Polyanions can attach side groups - boron-oxygen tetrahedra or triangles, their dimers or extraneous anions.

Ammonium, alkali, and also other metals in the +1 oxidation state most often form hydrated and anhydrous metaborates of the MBO 2 type, M 2 B 4 O 7 tetraborates, MB 5 O 8 pentaborates, and also M 4 B 10 O 17 decaborates n H 2 O. Alkaline earth and other metals in the + 2 oxidation state usually give hydrated metaborates, M 2 B 6 O 11 triborates and MB 6 O 10 hexaborates. as well as anhydrous meta-, ortho- and tetraborates. Metals in the + 3 oxidation state are characterized by hydrated and anhydrous MBO 3 orthoborates.

Borates are colorless amorphous substances or crystals (mainly with a low-symmetrical structure - monoclinic or rhombic). For anhydrous borates, the melting points are in the range from 500 to 2000 °C; the most high-melting metaborates are alkali and ortho- and metaborates of alkaline earth metals. Most borates easily form glasses when their melts are cooled. The hardness of hydrated borates on the Mohs scale is 2-5, anhydrous - up to 9.

Hydrated monoborates lose water of crystallization up to ~180°C, polyborates - at 300-500°C; elimination of water due to OH groups , coordinated around boron atoms occurs up to ~750°С. With complete dehydration, amorphous substances are formed, which at 500-800 ° C in most cases undergo “borate rearrangement” - crystallization, accompanied (for polyborates) by partial decomposition with the release of B 2 O 3.

Borates of alkali metals, ammonium and T1 (I) are soluble in water (especially meta- and pentaborates), hydrolyze in aqueous solutions (solutions have an alkaline reaction). Most borates are easily decomposed by acids, in some cases by the action of CO 2; and SO2;. Borates of alkaline earth and heavy metals interact with solutions of alkalis, carbonates and bicarbonates of alkali metals. Anhydrous borates are chemically more stable than hydrated ones. With some alcohols, in particular with glycerol, borates form water-soluble complexes. Under the action of strong oxidizing agents, in particular H 2 O 2, or during electrochemical oxidation, borates are converted into peroxoborates .

About 100 natural borates are known, which are mainly salts of Na, Mg, Ca, Fe.

Hydrated borates are obtained by: neutralization of H 3 BO 3 with metal oxides, hydroxides or carbonates; exchange reactions of alkali metal borates, most often Na, with salts of other metals; the reaction of mutual transformation of sparingly soluble borates with aqueous solutions of alkali metal borates; hydrothermal processes using alkali metal halides as mineralizing additives. Anhydrous borates are obtained by fusion or sintering of B 2 O 3 with metal oxides or carbonates or by dehydration of hydrates; single crystals are grown in solutions of borates in molten oxides, for example Bi 2 O 3 .

Borates are used: to obtain other boron compounds; as components of the charge in the production of glasses, glazes, enamels, ceramics; for fire-resistant coatings and impregnations; as components of fluxes for refining, welding and soldering of metal”; as pigments and fillers of paints and varnishes; as mordants in dyeing, corrosion inhibitors, components of electrolytes, phosphors, etc. Borax and calcium borates are most widely used.

2. Halides, chemical compounds of halogens with other elements. Halides usually include compounds in which the halogen atoms have a higher electronegativity than another element. Halides do not form He, Ne and Ar. To simple, or binary, halides EX n (n- most often an integer from 1 for monohalides to 7 for IF 7, and ReF 7, but can also be fractional, for example 7/6 for Bi 6 Cl 7) include, in particular, salts of hydrohalic acids and interhalogen compounds (for example, halofluorides). There are also mixed halides, polyhalides, hydrohalides, oxohalides, oxyhalides, hydroxohalides, thiohalides, and complex halides. The oxidation state of halogens in halides is usually -1.

According to the nature of the element-halogen bond, simple halides are divided into ionic and covalent. In reality, the relationships are of a mixed nature with the predominance of the contribution of one or another component. The halides of alkali and alkaline earth metals, as well as many mono- and dihalides of other metals, are typical salts in which the ionic nature of the bond prevails. Most of them are relatively refractory, low volatile, highly soluble in water; in aqueous solutions, they almost completely dissociate into ions. The properties of salts are also possessed by trihalides of rare earth elements. The water solubility of ionic halides generally decreases from iodides to fluorides. Chlorides, bromides and iodides Ag + , Сu + , Hg + and Pb 2+ are poorly soluble in water.

An increase in the number of halogen atoms in metal halides or the ratio of the metal charge to the radius of its ion leads to an increase in the covalent component of the bond, a decrease in water solubility and thermal stability of halides, an increase in volatility, an increase in oxidization, ability and tendency to hydrolysis. These dependences are observed for metal halides of the same period and in the series of halides of the same metal. They are easy to trace on the example of thermal properties. For example, for metal halides of the 4th period, the melting and boiling points are respectively 771 and 1430°C for KC1, 772 and 1960°C for CaCl 2, 967 and 975°C for ScCl 3 , -24.1 and 136°C for TiCl 4 . For UF 3, the melting point is ~ 1500 ° C, UF 4 1036 ° C, UF 5 348 ° C, UF 6 64.0 ° C. In the series of EC compounds n with the same n the covalence of the bond usually increases on going from fluorides to chlorides and decreases on going from the latter to bromides and iodides. So, for AlF 3, the sublimation temperature is 1280 ° C, A1C1 3 180 ° C, the boiling point of A1Br 3 is 254.8 ° C, AlI 3 407 ° C. In the series ZrF 4 , ZrCl 4 ZrBr 4 , ZrI 4 the sublimation temperature is 906, 334, 355 and 418°C, respectively. In the MF ranks n and MS1 n where M is a metal of one subgroup, the covalence of the bond decreases with increasing atomic mass of the metal. There are few metal fluorides and chlorides with approximately the same contribution of the ionic and covalent bond components.

The average element-halogen bond energy decreases when moving from fluorides to iodides and with increasing n(see table).

Many metal halides containing isolated or bridging O atoms (respectively, oxo- and oxyhalides), for example, vanadium oxotrifluoride VOF 3, niobium dioxyfluoride NbO 2 F, tungsten dioxodiiodide WO 2 I 2.

Complex halides (halogenometallates) contain complex anions in which the halogen atoms are ligands, for example, potassium hexachloroplatinate (IV) K 2 , sodium heptafluorotantalate (V) Na, lithium hexafluoroarsenate (V) Li. Fluoro-, oxofluoro- and chlorometallates have the highest thermal stability. By the nature of the bonds, ionic compounds with cations NF 4 + , N 2 F 3 + , C1F 2 + , XeF + and others are close to complex halides.

Many halides are characterized by association and polymerization in the liquid and gas phases with the formation of bridge bonds. The most prone to this are the halides of metals of groups I and II, AlCl 3 , pentafluorides of Sb and transition metals, oxofluorides of the composition MOF 4 . Known halides with a metal-metal bond, for example. Cl-Hg-Hg-Cl.

Fluorides differ significantly in properties from other halides. However, in simple halides, these differences are less pronounced than in the halogens themselves, and in complex halides, they are less pronounced than in simple ones.

Many covalent halides (especially fluorides) are strong Lewis acids, e.g. AsF 5 , SbF 5 , BF 3 , A1C1 3 . Fluorides are part of superacids. Higher halides are reduced by metals and hydrogen, for example:

5WF 6 + W = 6WF 5

TiCl 4 + 2Mg \u003d Ti + 2MgCl 2

UF 6 + H 2 \u003d UF 4 + 2HF

Metal halides of groups V-VIII, except for Cr and Mn, are reduced by H 2 to metals, for example:

WF 6 + ZN 2 = W + 6HF

Many covalent and ionic metal halides interact with each other to form complex halides, for example:

KC1 + TaCl 5 = K

The lighter halogens can displace the heavier ones from the halides. Oxygen can oxidize halides with the release of C1 2 , Br 2 , and I 2 . One of the characteristic reactions of covalent halides is the interaction with water (hydrolysis) or its vapors when heated (pyrohydrolysis), leading to the formation of oxides, oxy- or oxo halides, hydroxides and hydrogen halides.

Halides are obtained directly from the elements, by the interaction of hydrogen halides or hydrohalic acids with elements, oxides, hydroxides or salts, as well as by exchange reactions.

Halides are widely used in engineering as starting materials for the production of halogens, alkali and alkaline earth metals, and as components of glasses and other inorganic materials; they are intermediate products in the production of rare and some non-ferrous metals, U, Si, Ge, etc.

In nature, halides form separate classes of minerals, which include fluorides (eg, the minerals fluorite, cryolite) and chlorides (sylvite, carnallite). Bromine and iodine are present in some minerals as isomorphic impurities. Significant amounts of halides are found in the water of the seas and oceans, in salt and underground brines. Some halides, such as NaCl, KC1, CaCl 2, are part of living organisms.

3. Carbonates (from lat. carbo, genus case carbonis coal), salts of carbonic acid. There are medium carbonates with the CO 3 2- anion and acidic, or bicarbonates (obsolete bicarbonates), with the HCO 3 - anion. Carbonates are crystalline substances. Most of the medium metal salts in the oxidation state + 2 crystallize into a hexagon. lattice type of calcite or rhombic type of aragonite.

Of the medium carbonates, only salts of alkali metals, ammonium and Tl (I) dissolve in water. As a result of significant hydrolysis, their solutions have an alkaline reaction. The most difficult soluble metal carbonates in the oxidation state + 2. On the contrary, all bicarbonates are highly soluble in water. During exchange reactions in aqueous solutions between metal salts and Na 2 CO 3, precipitates of medium carbonates are formed when their solubility is much lower than that of the corresponding hydroxides. This is the case for Ca, Sr and their analogues, lanthanides, Ag(I), Mn(II), Pb(II), and Cd(II). The remaining cations, when interacting with dissolved carbonates as a result of hydrolysis, can give not average, but basic carbonates or even hydroxides. Medium carbonates containing multiply charged cations can sometimes be precipitated from aqueous solutions in the presence of a large excess of CO 2 .

The chemical properties of carbonates are due to their belonging to the class of inorganic salts of weak acids. The characteristic features of carbonates are associated with their poor solubility, as well as the thermal instability of both the crabonates themselves and H 2 CO 3 . These properties are used in the analysis of crabonates, based either on their decomposition by strong acids and the quantitative absorption of the CO 2 released in this case by an alkali solution, or on the precipitation of the CO 3 2- ion from the solution in the form of ВаСО 3 . Under the action of an excess of CO 2 on a precipitate of an average carbonate in a solution, a bicarbonate is formed, for example: CaCO 3 + H 2 O + CO 2 \u003d Ca (HCO 3) 2. The presence of bicarbonates in natural water determines its temporary hardness. Hydrocarbonates upon slight heating already at low temperatures are again converted into medium carbonates, which, upon heating, decompose to oxide and CO 2. The more active the metal, the higher the decomposition temperature of its carbonate. So, Na 2 CO 3 melts without decomposition at 857 °C, and for Ca, Mg and Al carbonates, the equilibrium decomposition pressures reach 0.1 MPa at temperatures of 820, 350 and 100 °C, respectively.

Carbonates are very widespread in nature, which is due to the participation of CO 2 and H 2 O in the processes of mineral formation. carbonates play a large role in global equilibriums between gaseous CO 2 in the atmosphere and dissolved CO 2 ;

and HCO 3 - and CO 3 2- ions in the hydrosphere and solid salts in the lithosphere. The most important minerals are CaCO 3 calcite, MgCO 3 magnesite, FeCO 3 siderite, ZnCO 3 smithsonite and some others. Limestone consists mainly of calcite or calcite skeletal remains of organisms, rarely of aragonite. Natural hydrated carbonates of alkali metals and Mg are also known (for example, MgCO 3 ZH 2 O, Na 2 CO 3 10H 2 O), double carbonates [for example, dolomite CaMg (CO 3) 2, throne Na 2 CO 3 NaHCO 3 2H 2 O] and basic [malachite CuCO 3 Cu(OH) 2, hydrocerussite 2РbСО 3 Pb(OH) 2].

The most important are potassium carbonate, calcium carbonate and sodium carbonate. Many natural carbonates are very valuable metal ores (for example, carbonates of Zn, Fe, Mn, Pb, Cu). Bicarbonates play an important physiological role, being buffer substances that regulate the constancy of blood pH.

4. Nitrates, salts of nitric acid HNO 3. Known for almost all metals; exist both in the form of anhydrous salts M (NO 3) n (n- the degree of oxidation of the metal M), and in the form of crystalline hydrates M (NO 3) n x H 2 O ( X= 1-9). From aqueous solutions at a temperature close to room temperature, only alkali metal nitrates crystallize anhydrous, the rest - in the form of crystalline hydrates. The physicochemical properties of anhydrous and hydrated nitrate of the same metal can be very different.

Anhydrous crystalline compounds of d-element nitrates are colored. Conventionally, nitrates can be divided into compounds with a predominantly covalent type of bond (salts of Be, Cr, Zn, Fe, and other transition metals) and with a predominantly ionic type of bond (salts of alkali and alkaline earth metals). Ionic nitrates are characterized by higher thermal stability, the predominance of crystal structures of higher symmetry (cubic) and the absence of splitting of the nitrate ion bands in the IR spectra. Covalent nitrates have a higher solubility in organic solvents, lower thermal stability, their IR spectra are more complex; some covalent nitrates are volatile at room temperature, and when dissolved in water, they partially decompose with the release of nitrogen oxides.

All anhydrous nitrates show strong oxidizing properties due to the presence of the NO 3 - ion, while their oxidizing ability increases when moving from ionic to covalent nitrates. The latter decompose in the range of 100-300°C, ionic - at 400-600°C (NaNO 3 , KNO 3 and some others melt when heated). Decomposition products in solid and liquid phases. are sequentially nitrites, oxonitrates and oxides, sometimes - free metals (when the oxide is unstable, for example Ag 2 O), and in the gas phase - NO, NO 2, O 2 and N 2. The composition of the decomposition products depends on the nature of the metal and its degree of oxidation, heating rate, temperature, composition of the gaseous medium, and other conditions. NH 4 NO 3 detonates, and when heated rapidly it can decompose with an explosion, in this case N 2 , O 2 and H 2 O are formed; when heated slowly, it decomposes into N 2 O and H 2 O.

The free NO 3 - ion in the gas phase has the geometric structure of an equilateral triangle with an N atom in the center, ONO angles ~ 120°, and N-O bond lengths of 0.121 nm. In crystalline and gaseous nitrates, the NO 3 ion - basically retains its shape and size, which determines the space and structure of nitrates. Ion NO 3 - can act as a mono-, bi-, tridentate or bridging ligand, so nitrates are characterized by a wide variety of types of crystal structures.

Transition metals in high oxidation states due to steric. difficulties cannot form anhydrous nitrates, and they are characterized by oxonitrates, for example UO 2 (NO 3) 2, NbO (NO 3) 3. Nitrates form a large number of double and complex salts with the NO 3 ion - in the inner sphere. In aqueous media, as a result of hydrolysis, transition metal cations form hydroxonitrates (basic nitrates) of variable composition, which can also be isolated in the solid state.

Hydrated nitrates differ from anhydrous ones in that in their crystal structures, the metal ion is in most cases associated with water molecules, and not with the NO 3 ion. Therefore, they dissolve better than anhydrous nitrates in water, but worse - in organic solvents, weaker oxidizing agents melt incongruently in crystallization water in the range of 25-100°C. When hydrated nitrates are heated, as a rule, anhydrous nitrates are not formed, but thermolysis occurs with the formation of hydroxonitrates and then oxonitrates and metal oxides.

In many of their chemical properties, nitrates are similar to other inorganic salts. The characteristic features of nitrates are due to their very high solubility in water, low thermal stability and the ability to oxidize organic and inorganic compounds. During the reduction of nitrates, a mixture of nitrogen-containing products NO 2 , NO, N 2 O, N 2 or NH 3 is formed with the predominance of one of them depending on the type of reducing agent, temperature, reaction of the medium, and other factors.

Industrial methods for producing nitrates are based on the absorption of NH 3 by HNO 3 solutions (for NH 4 NO 3) or the absorption of nitrous gases (NO + NO 2) by alkali or carbonate solutions (for alkali metal nitrates, Ca, Mg, Ba), as well as on various exchange reactions of metal salts with HNO 3 or alkali metal nitrates. In the laboratory, to obtain anhydrous nitrates, reactions of transition metals or their compounds with liquid N 2 O 4 and its mixtures with organic solvents or reactions with N 2 O 5 are used.

Nitrates Na, K (sodium and potassium nitrate) are found in the form of natural deposits.

Nitrates are used in many industries. Ammonium nitrite (ammonium nitrate) - the main nitrogen-containing fertilizer; nitrates of alkali metals and Ca are also used as fertilizers. Nitrates - components of rocket fuels, pyrotechnic compositions, pickling solutions for dyeing fabrics; they are used for hardening metals, food preservation, as medicines, and for the production of metal oxides.

Nitrates are toxic. They cause pulmonary edema, cough, vomiting, acute cardiovascular insufficiency, etc. The lethal dose of nitrates for humans is 8-15 g, the allowable daily intake is 5 mg / kg. For the sum of Na, K, Ca, NH3 nitrates MPC: in water 45 mg/l", in soil 130 mg/kg (hazard class 3); in vegetables and fruits (mg/kg) - potatoes 250, late white cabbage 500, late carrots 250, beets 1400, onions 80, zucchini 400, melons 90, watermelons, grapes, apples, pears 60. Non-compliance with agrotechnical recommendations, excessive fertilization dramatically increases the content of nitrates in agricultural products, surface runoff from fields ( 40-5500 mg/l), ground water.

5. Nitrites, salts of nitrous acid HNO 2. First of all, nitrites of alkali metals and ammonium are used, less - alkaline earth and Z d-metals, Pb and Ag. There is only fragmentary information about the nitrites of other metals.

Metal nitrites in the +2 oxidation state form crystal hydrates with one, two or four water molecules. Nitrites form double and triple salts, for example. CsNO 2 AgNO 2 or Ba (NO 2) 2 Ni (NO 2) 2 2KNO 2, as well as complex compounds, such as Na 3.

Crystal structures are known only for a few anhydrous nitrites. The NO 2 anion has a non-linear configuration; ONO angle 115°, H-O bond length 0.115 nm; the type of connection M-NO 2 is ionic-covalent.

K, Na, Ba nitrites are well soluble in water, Ag, Hg, Cu nitrites are poorly soluble. With increasing temperature, the solubility of nitrites increases. Almost all nitrites are poorly soluble in alcohols, ethers, and low-polarity solvents.

Nitrites are thermally unstable; melt without decomposition only nitrites of alkali metals, nitrites of other metals decompose at 25-300 °C. The mechanism of nitrite decomposition is complex and includes a number of parallel-sequential reactions. The main gaseous decomposition products are NO, NO 2, N 2 and O 2, solid ones are metal oxide or elemental metal. The release of a large amount of gases causes the explosive decomposition of some nitrites, for example NH 4 NO 2, which decomposes into N 2 and H 2 O.

The characteristic features of nitrites are associated with their thermal instability and the ability of the nitrite ion to be both an oxidizing agent and a reducing agent, depending on the medium and the nature of the reagents. In a neutral environment, nitrites are usually reduced to NO, in an acidic environment they are oxidized to nitrates. Oxygen and CO 2 do not interact with solid nitrites and their aqueous solutions. Nitrites contribute to the decomposition of nitrogen-containing organic substances, in particular amines, amides, etc. With organic halides RXH. react to form both RONO nitrites and RNO 2 nitro compounds.

The industrial production of nitrites is based on the absorption of nitrous gas (a mixture of NO + NO 2) with solutions of Na 2 CO 3 or NaOH with successive crystallization of NaNO 2; nitrites of other metals in industry and laboratories are obtained by the exchange reaction of metal salts with NaNO 2 or by the reduction of nitrates of these metals.

Nitrites are used for the synthesis of azo dyes, in the production of caprolactam, as oxidizing and reducing agents in the rubber, textile and metalworking industries, as food preservatives. Nitrites such as NaNO 2 and KNO 2 are toxic, causing headache, vomiting, respiratory depression, etc. When NaNO 2 is poisoned, methemoglobin is formed in the blood, erythrocyte membranes are damaged. Perhaps the formation of nitrosamines from NaNO 2 and amines directly in the gastrointestinal tract.

6. Sulfates, salts of sulfuric acid. Medium sulfates with the anion SO 4 2- are known, acidic, or hydrosulfates, with the anion HSO 4 - , basic, containing along with the anion SO 4 2- - OH groups, for example Zn 2 (OH) 2 SO 4 . There are also double sulfates, which include two different cations. These include two large groups of sulfates - alum , as well as chenites M 2 E (SO 4) 2 6H 2 O , where M is a singly charged cation, E is Mg, Zn and other doubly charged cations. Known triple sulfate K 2 SO 4 MgSO 4 2CaSO 4 2H 2 O (mineral polygalite), double basic sulfates, for example, minerals of the alunite and jarosite groups M 2 SO 4 Al 2 (SO 4) 3 4Al (OH 3 and M 2 SO 4 Fe 2 (SO 4) 3 4Fe (OH) 3, where M is a singly charged cation.Sulfates can be part of mixed salts, for example 2Na 2 SO 4 Na 2 CO 3 (mineral berkite), MgSO 4 KCl 3H 2 O (kainite) .

Sulfates are crystalline substances, medium and acidic, in most cases they are highly soluble in water. Slightly soluble sulfates of calcium, strontium, lead and some others, practically insoluble BaSO 4 , RaSO 4 . Basic sulfates are usually sparingly soluble or practically insoluble, or hydrolyzed by water. Sulfates can crystallize from aqueous solutions in the form of crystalline hydrates. The crystalline hydrates of some heavy metals are called vitriol; copper sulfate СuSO 4 5H 2 O, ferrous sulfate FeSO 4 7H 2 O.

Medium alkali metal sulfates are thermally stable, while acid sulfates decompose when heated, turning into pyrosulfates: 2KHSO 4 \u003d H 2 O + K 2 S 2 O 7. Average sulfates of other metals, as well as basic sulfates, when heated to sufficiently high temperatures, as a rule, decompose with the formation of metal oxides and the release of SO 3 .

Sulfates are widely distributed in nature. They are found in the form of minerals, such as gypsum CaSO 4 H 2 O, mirabilite Na 2 SO 4 10H 2 O, and are also part of sea and river water.

Many sulfates can be obtained by the interaction of H 2 SO 4 with metals, their oxides and hydroxides, as well as the decomposition of salts of volatile acids with sulfuric acid.

Inorganic sulfates are widely used. For example, ammonium sulfate is a nitrogen fertilizer, sodium sulfate is used in the glass, paper industry, viscose production, etc. Natural sulfate minerals are raw materials for the industrial production of compounds of various metals, building materials, etc.

7.sulfites, salts of sulfurous acid H 2 SO 3 . There are medium sulfites with the anion SO 3 2- and acidic (hydrosulfites) with the anion HSO 3 - . Medium sulfites are crystalline substances. Ammonium and alkali metal sulfites are highly soluble in water; solubility (g in 100 g): (NH 4) 2 SO 3 40.0 (13 ° C), K 2 SO 3 106.7 (20 ° C). In aqueous solutions they form hydrosulfites. Sulfites of alkaline earth and some other metals are practically insoluble in water; solubility of MgSO 3 1 g in 100 g (40°C). Known crystalline hydrates (NH 4) 2 SO 3 H 2 O, Na 2 SO 3 7H 2 O, K 2 SO 3 2H 2 O, MgSO 3 6H 2 O, etc.

Anhydrous sulfites, when heated without access to air in sealed vessels, disproportionate into sulfides and sulfates, when heated in a stream of N 2 they lose SO 2, and when heated in air, they are easily oxidized to sulfates. With SO 2 in the aquatic environment, medium sulfites form hydrosulfites. Sulfites are relatively strong reducing agents; they are oxidized in solutions with chlorine, bromine, H 2 O 2, etc. to sulfates. They are decomposed by strong acids (for example, HC1) with the release of SO 2.

Crystalline hydrosulfites are known for K, Rb, Cs, NH 4 + , they are unstable. Other hydrosulfites exist only in aqueous solutions. Density NH 4 HSO 3 2.03 g/cm 3 ; solubility in water (g per 100 g): NH 4 HSO 3 71.8 (0 ° C), KHSO 3 49 (20 ° C).

When crystalline hydrosulfites Na or K are heated, or when the slurry solution of the pulp M 2 SO 3 is saturated with SO 2, pyrosulfites (obsolete - metabisulfites) M 2 S 2 O 5 are formed - salts of pyrosulfurous acid unknown in the free state H 2 S 2 O 5; crystals, unstable; density (g / cm 3): Na 2 S 2 O 5 1.48, K 2 S 2 O 5 2.34; above ~ 160 °С they decompose with the release of SO 2; dissolve in water (with decomposition to HSO 3 -), solubility (g per 100 g): Na 2 S 2 O 5 64.4, K 2 S 2 O 5 44.7; form hydrates Na 2 S 2 O 5 7H 2 O and ZK 2 S 2 O 5 2H 2 O; reducing agents.

Medium alkali metal sulfites are obtained by reacting an aqueous solution of M 2 CO 3 (or MOH) with SO 2 , and MSO 3 by passing SO 2 through an aqueous suspension of MCO 3 ; mainly SO 2 is used from the off-gases of contact sulfuric acid production. Sulfites are used in bleaching, dyeing and printing of fabrics, fibers, leather for grain conservation, green fodder, industrial feed waste (NaHSO 3 ,

Na 2 S 2 O 5). CaSO 3 and Ca(HSO 3) 2 - disinfectants in winemaking and sugar industry. NaНSO 3 , MgSO 3 , NH 4 НSO 3 - components of sulfite liquor during pulping; (NH 4) 2 SO 3 - SO 2 absorber; NaHSO 3 is an H 2 S absorber from production waste gases, a reducing agent in the production of sulfur dyes. K 2 S 2 O 5 - component of acid fixers in photography, antioxidant, antiseptic.

Mixture separation methods

Filtration, separation of inhomogeneous systems liquid - solid particles (suspensions) and gas - solid particles using porous filter partitions (FP) that allow liquid or gas to pass through, but retain solid particles. The driving force of the process is the pressure difference on both sides of the FP.

When separating suspensions, solid particles usually form a layer of wet sediment on the FP, which, if necessary, is washed with water or other liquid, and also dehydrated by blowing air or other gas through it. Filtration is carried out at a constant pressure difference or at a constant process speed w(the amount of filtrate in m 3 passing through 1 m 2 of the FP surface per unit time). At a constant pressure difference, the suspension is fed to the filter under vacuum or overpressure, as well as by a piston pump; when using a centrifugal pump, the pressure difference increases and the process speed decreases.

Depending on the concentration of suspensions, several types of filtration are distinguished. At a concentration of more than 1%, filtration occurs with the formation of a precipitate, and at a concentration of less than 0.1%, with clogging of the pores of the FP (clarification of liquids). If a sufficiently dense sediment layer is not formed on the FP and solid particles get into the filtrate, it is filtered using finely dispersed auxiliary materials (diatomite, perlite), which are previously applied to the FP or added to the suspension. At an initial concentration of less than 10%, partial separation and thickening of suspensions is possible.

A distinction is made between continuous and intermittent filters. For the latter, the main stages of work are filtration, washing of the sediment, its dehydration and unloading. At the same time, optimization is applicable according to the criteria of the highest productivity and the lowest costs. If washing and dehydration are not performed, and the hydraulic resistance of the partition can be neglected, then the highest productivity is achieved when the filtration time is equal to the duration of the auxiliary operations.

Applicable flexible FP made of cotton, wool, synthetic and glass fabrics, as well as non-woven FP made of natural and synthetic fibers and inflexible - ceramic, cermet and foam plastic. The directions of movement of the filtrate and the action of gravity can be opposite, coincide or be mutually perpendicular.

Filter designs are varied. One of the most common is a rotating drum vacuum filter. (cm. Fig.) of continuous action, in which the directions of movement of the filtrate and the action of gravity are opposite. The switchgear section connects zones I and II to a vacuum source and zones III and IV to a compressed air source. The filtrate and wash liquid from zones I and II enter separate receivers. The automated intermittent filter press with horizontal chambers, filter cloth in the form of an endless belt and elastic membranes for sludge dewatering by pressing has also become widespread. It performs alternating operations of filling the chambers with a suspension, filtering, washing and dehydrating the sediment, separating adjacent chambers and removing the sediment.

Determination of dynamic shear stress, effective and plastic viscosity at normal temperature

Determination of dynamic shear stress, effective and plastic viscosity at elevated temperature

Experience 2. Obtaining and studying the properties of phosphoric acid salts.

In everyday life, people rarely encounter Most objects are mixtures of substances.

A solution is one in which the components are uniformly mixed. There are several types according to particle size: coarse systems, molecular solutions and colloidal systems, which are often called sols. In this article we are talking about molecular (or Solubility of substances in water - one of the main conditions affecting the formation of compounds.

Solubility of substances: what is it and why is it needed

To understand this topic, you need to know the solubility of substances. In simple terms, this is the ability of a substance to combine with another and form a homogeneous mixture. From a scientific point of view, a more complex definition can be considered. The solubility of substances is their ability to form homogeneous (or heterogeneous) compositions with one or more substances with a dispersed distribution of components. There are several classes of substances and compounds:

• soluble;
• sparingly soluble;
• insoluble.

What is the measure of the solubility of a substance

The content of a substance in a saturated mixture is a measure of its solubility. As mentioned above, for all substances it is different. Soluble are those that can dilute more than 10g of themselves in 100g of water. The second category is less than 1 g under the same conditions. Practically insoluble are those in the mixture of which less than 0.01 g of the component passes. In this case, the substance cannot transfer its molecules to water.

What is the solubility coefficient

The solubility coefficient (k) is an indicator of the maximum mass of a substance (g) that can be dissolved in 100 g of water or another substance.

Solvents

This process involves a solvent and a solute. The first differs in that initially it is in the same state of aggregation as the final mixture. As a rule, it is taken in larger quantities.

However, many people know that water occupies a special place in chemistry. There are separate rules for it. A solution in which H 2 O is present is called an aqueous solution. When talking about them, the liquid is an extractant even when it is in a smaller amount. An example is an 80% solution of nitric acid in water. The proportions here are not equal. Although the proportion of water is less than that of acids, it is incorrect to call the substance a 20% solution of water in nitric acid.

There are mixtures that do not contain H 2 O. They will be called non-aqueous. Such electrolyte solutions are ionic conductors. They contain single or mixtures of extractants. They are composed of ions and molecules. They are used in industries such as medicine, the production of household chemicals, cosmetics and other areas. They can combine several desired substances with different solubility. The components of many products that are applied externally are hydrophobic. In other words, they do not interact well with water. In these, they can be volatile, non-volatile and combined. Organic substances in the first case dissolve fats well. The volatiles include alcohols, hydrocarbons, aldehydes, and others. They are often included in household chemicals. Non-volatile are most often used for the manufacture of ointments. These are fatty oils, liquid paraffin, glycerin and others. Combined is a mixture of volatile and non-volatile, for example, ethanol with glycerin, glycerin with dimexide. They may also contain water.

Types of solutions by degree of saturation

A saturated solution is a mixture of chemicals containing the maximum concentration of one substance in a solvent at a certain temperature. It will not breed further. In the preparation of a solid substance, precipitation is noticeable, which is in dynamic equilibrium with it. This concept means a state that persists in time due to its flow simultaneously in two opposite directions (forward and reverse reactions) at the same speed.

If a substance can still decompose at a constant temperature, then this solution is unsaturated. They are stable. But if you continue to add a substance to them, then it will be diluted in water (or other liquid) until it reaches its maximum concentration.

Another type is oversaturated. It contains more solute than can be at a constant temperature. Due to the fact that they are in an unstable equilibrium, crystallization occurs when they are physically affected.

How can you tell a saturated solution from an unsaturated one?

This is easy enough to do. If the substance is a solid, then a precipitate can be seen in a saturated solution. In this case, the extractant can thicken, as, for example, in a saturated composition, water to which sugar has been added.
But if you change the conditions, increase the temperature, then it will no longer be considered saturated, since at a higher temperature the maximum concentration of this substance will be different.

Theories of interaction of components of solutions

There are three theories regarding the interaction of elements in a mixture: physical, chemical and modern. The authors of the first one are Svante August Arrhenius and Wilhelm Friedrich Ostwald. They assumed that, due to diffusion, the particles of the solvent and the solute were evenly distributed throughout the volume of the mixture, but there was no interaction between them. The chemical theory put forward by Dmitri Ivanovich Mendeleev is the opposite of it. According to it, as a result of chemical interaction between them, unstable compounds of constant or variable composition are formed, which are called solvates.

At present, the unified theory of Vladimir Aleksandrovich Kistyakovsky and Ivan Alekseevich Kablukov is used. It combines physical and chemical. The modern theory says that in the solution there are both non-interacting particles of substances and the products of their interaction - solvates, the existence of which Mendeleev proved. In the case when the extractant is water, they are called hydrates. The phenomenon in which solvates (hydrates) are formed is called solvation (hydration). It affects all physical and chemical processes and changes the properties of the molecules in the mixture. Solvation occurs due to the fact that the solvation shell, consisting of molecules of the extractant closely associated with it, surrounds the solute molecule.

Factors affecting the solubility of substances

Chemical composition of substances. The rule "like attracts like" applies to reagents as well. Substances that are similar in physical and chemical properties can mutually dissolve faster. For example, non-polar compounds interact well with non-polar ones. Substances with polar molecules or an ionic structure are diluted in polar ones, for example, in water. Salts, alkalis and other components decompose in it, while non-polar ones do the opposite. A simple example can be given. To prepare a saturated solution of sugar in water, a larger amount of substance is required than in the case of salt. What does it mean? Simply put, you can dilute much more sugar in water than salt.

Temperature. To increase the solubility of solids in liquids, you need to increase the temperature of the extractant (works in most cases). An example can be shown. If you put a pinch of sodium chloride (salt) in cold water, this process will take a long time. If you do the same with a hot medium, then the dissolution will be much faster. This is explained by the fact that as a result of an increase in temperature, kinetic energy increases, a significant amount of which is often spent on the destruction of bonds between molecules and ions of a solid. However, when the temperature rises in the case of lithium, magnesium, aluminum and alkali salts, their solubility decreases.

Pressure. This factor only affects gases. Their solubility increases with increasing pressure. After all, the volume of gases is reduced.

Changing the dissolution rate

Do not confuse this indicator with solubility. After all, different factors influence the change in these two indicators.

The degree of fragmentation of the solute. This factor affects the solubility of solids in liquids. In the whole (lumpy) state, the composition is diluted longer than the one that is broken into small pieces. Let's take an example. A solid block of salt will take much longer to dissolve in water than salt in the form of sand.

Stirring speed. As is known, this process can be catalyzed by stirring. Its speed is also important, because the faster it is, the faster the substance will dissolve in the liquid.

Why is it important to know the solubility of solids in water?

First of all, such schemes are needed to correctly solve chemical equations. In the solubility table there are charges of all substances. They need to be known in order to correctly record the reagents and draw up the equation of a chemical reaction. Solubility in water indicates whether the salt or base can dissociate. Aqueous compounds that conduct current have strong electrolytes in their composition. There is another type. Those that conduct current poorly are considered weak electrolytes. In the first case, the components are substances that are completely ionized in water. Whereas weak electrolytes show this indicator only to a small extent.

Chemical reaction equations

There are several types of equations: molecular, complete ionic and short ionic. In fact, the last option is a shortened form of molecular. This is the final answer. The complete equation contains the reactants and products of the reaction. Now comes the turn of the solubility table of substances. First you need to check whether the reaction is feasible, that is, whether one of the conditions for the reaction is met. There are only 3 of them: the formation of water, the release of gas, precipitation. If the first two conditions are not met, you need to check the last one. To do this, you need to look at the solubility table and find out if there is an insoluble salt or base in the reaction products. If it is, then this will be the sediment. Further, the table will be required to write the ionic equation. Since all soluble salts and bases are strong electrolytes, they will decompose into cations and anions. Further, unbound ions are reduced, and the equation is written in a short form. Example:

1. K 2 SO 4 + BaCl 2 \u003d BaSO 4 ↓ + 2HCl,
2. 2K + 2SO 4 + Ba + 2Cl \u003d BaSO 4 ↓ + 2K + 2Cl,
3. Ba+SO4=BaSO4 ↓.

Thus, the table of solubility of substances is one of the key conditions for solving ionic equations.

A detailed table helps you find out how much component you need to take to prepare a rich mixture.

Solubility table

This is what the usual incomplete table looks like. It is important that the temperature of the water is indicated here, as it is one of the factors that we have already mentioned above.

How to use the table of solubility of substances?

The table of solubility of substances in water is one of the main assistants of a chemist. It shows how various substances and compounds interact with water. The solubility of solids in a liquid is an indicator without which many chemical manipulations are impossible.

The table is very easy to use. Cations (positively charged particles) are written on the first line, anions (negatively charged particles) are written on the second line. Most of the table is occupied by a grid with certain symbols in each cell. These are the letters "P", "M", "H" and the signs "-" and "?".

• "P" - the compound is dissolved;
• "M" - dissolves a little;
• "H" - does not dissolve;
• "-" - connection does not exist;
• "?" - there is no information about the existence of the connection.

There is one empty cell in this table - it is water.

Simple example

Now about how to work with such material. Let's say you need to find out if salt is soluble in water - MgSo 4 (magnesium sulfate). To do this, you need to find the Mg 2+ column and go down it to the SO 4 2- line. At their intersection is the letter P, which means the compound is soluble.

Conclusion

So, we have studied the issue of the solubility of substances in water and not only. Without a doubt, this knowledge will be useful in the further study of chemistry. After all, the solubility of substances plays an important role there. It is useful in solving chemical equations and various problems.