A combination of salt and acid. Chemical properties of salts

Reasons

Bases are compounds containing only hydroxide ions OH - as an anion. The number of hydroxide ions that can be replaced by an acidic residue determines the acidity of the base. In this regard, bases are one-, two- and polyacid; however, true bases most often include one- and two-acid. Among them, water-soluble and water-insoluble bases should be distinguished. Please note that bases that are soluble in water and dissociate almost completely are called alkalis (strong electrolytes). These include hydroxides of alkali and alkaline earth elements and in no case a solution of ammonia in water.

The name of the base begins with the word hydroxide, after which the Russian name of the cation is given in the genitive case, and its charge is indicated in parentheses. It is allowed to list the number of hydroxide ions using the prefixes di-, tri-, tetra. For example: Mn(OH) 3 - manganese (III) hydroxide or manganese trihydroxide.

Note that there is a genetic relationship between bases and basic oxides: basic oxides correspond to bases. Therefore, base cations most often have a charge of one or two, which corresponds to the lowest oxidation states of metals.

Remember the basic ways to obtain bases

1. Interaction of active metals with water:

2Na + 2H 2 O = 2NaOH + H 2

La + 6H 2 O = 2La(OH) 3 + 3H 2

Interaction of basic oxides with water:

CaO + H 2 O = Ca (OH) 2

MgO + H 2 O = Mg(OH) 2.

3. Interaction of salts with alkalis:

MnSO 4 + 2KOH = Mn(OH) 2 ↓ + K 2 SO 4

NH 4 С1 + NaOH = NaCl + NH 3 ∙ H 2 O

Na 2 CO 3 + Ca(OH) 2 = 2NaOH + CaCO 3

MgOHCl + NaOH = Mg(OH) 2 + NaCl.

Electrolysis of aqueous salt solutions with a diaphragm:

2NaCl + 2H 2 O → 2NaOH + Cl 2 + H 2

Please note that in step 3, the starting reagents must be selected in such a way that among the reaction products there is either a sparingly soluble compound or a weak electrolyte.

Note that when considering the chemical properties of bases, reaction conditions depend on the solubility of the base.

1. Interaction with acids:

NaOH + H 2 SO 4 = NaHSO 4 + H 2 O

2NaOH + H 2 SO 4 = Na 2 SO 4 + 2H 2 O

2Mg(OH) 2 + H 2 SO 4 = (MgOH) 2 SO 4 + 2H 2 O

Mg(OH) 2 + H 2 SO 4 = MgSO 4 + 2H 2 O

Mg(OH) 2 + 2H 2 SO 4 = Mg(HSO 4) 2 + 2H 2 O

2. Interaction with acid oxides:

NaOH + CO 2 = NaHCO 3

2NaOH + CO 2 = Na 2 CO 3 + H 2 O

Fe(OH) 2 + P 2 O 5 = Fe(PO 3) 2 + H 2 O

3Fe(OH) 2 + P 2 O 5 = Fe 3 (PO 4) 2 + 2H 2 O

3. Interaction with amphoteric oxides:

A1 2 O 3 + 2NaOH p + 3H 2 O = 2Na

Al 2 O 3 + 2NaOH T = 2NaAlO 2 + H 2 O

Cr 2 O 3 + Mg(OH) 2 = Mg(CrO 2) 2 + H 2 O

4. Interaction with ampheteric hydroxides:

Ca(OH) 2 + 2Al(OH) 3 = Ca(AlO 2) 2 + 4H 2 O

3NaOH + Cr(OH) 3 = Na 3

Interaction with salts.

To the reactions described in point 3 of the production methods, the following should be added:

2ZnSO 4 + 2KOH = (ZnOH) 2 S0 4 + K 2 SO 4

NaHCO 3 + NaOH = Na 2 CO 3 + H 2 O

BeSO 4 + 4NaOH = Na 2 + Na 2 SO 4

Cu(OH) 2 + 4NH 3 ∙H 2 O = (OH) 2 + 4H 2 O

6. Oxidation to amphoteric hydroxides or salts:

4Fe(OH) 2 + O 2 + 2H 2 O = 4Fe(OH) 3

2Cr(OH) 2 + 2H 2 O + Na 2 O 2 + 4NaOH = 2Na 3.

7. Heat decomposition:

Ca(OH) 2 = CaO + H 2 O.

Please note that alkali metal hydroxides, except lithium, do not participate in such reactions.

!!!Are there alkaline precipitations?!!! Yes, there are, but they are not as widespread as acid precipitation, are little known, and their impact on environmental objects has been practically unstudied. Nevertheless, their consideration deserves attention.

The origin of alkaline precipitation can be explained as follows.

CaCO 3 →CaO + CO 2

In the atmosphere, calcium oxide combines with water vapor during condensation, with rain or sleet, forming calcium hydroxide:

CaO + H 2 O →Ca(OH) 2,

which creates an alkaline reaction of atmospheric precipitation. In the future, it is possible to react calcium hydroxide with carbon dioxide and water to form calcium carbonate and calcium bicarbonate:

Ca(OH) 2 + CO 2 → CaCO 3 + H 2 O;

CaCO 3 + CO 2 + H 2 O → Ca(HC0 3) 2.

Chemical analysis of rainwater showed that it contains sulfate and nitrate ions in small quantities (about 0.2 mg/l). As is known, the cause of the acidic nature of precipitation is sulfuric and nitric acids. At the same time, there is a high content of calcium cations (5-8 mg/l) and bicarbonate ions, the content of which in the area of ​​​​the construction complex enterprises is 1.5-2 times higher than in other areas of the city, and amounts to 18-24 mg /l. This shows that the calcium carbonate system and the processes occurring in it play a major role in the formation of local alkaline sediments, as mentioned above.

Alkaline precipitation affects plants; changes in the phenotypic structure of plants are noted. There are traces of “burns” on the leaf blades, a white coating on the leaves and a depressed state of herbaceous plants.

SALT, a class of chemical compounds. There is currently no generally accepted definition of the concept of “Salts,” as well as the terms “acids and bases,” the reaction products of which salts are. Salts can be considered as products of the replacement of acid hydrogen protons with metal ions, NH 4 +, CH 3 NH 3 + and other cations or OH groups of the base with acid anions (for example, Cl -, SO 4 2-).

Classification

The products of complete substitution are medium salts, for example. Na 2 SO 4, MgCl 2, partially acidic or basic salts, for example KHSO 4, СuСlОН. There are also simple salts, including one type of cations and one type of anions (for example, NaCl), double salts containing two types of cations (for example, KAl(SO 4) 2 12H 2 O), mixed salts, which contain two types of acid residues ( for example AgClBr). Complex salts contain complex ions, such as K4.

Physical properties

Typical salts are crystalline substances with an ionic structure, for example CsF. There are also covalent salts, for example AlCl 3. In fact, the nature of the chemical bond of many salts is mixed.

Based on their solubility in water, they distinguish between soluble, slightly soluble and practically insoluble salts. Soluble salts include almost all sodium, potassium and ammonium salts, many nitrates, acetates and chlorides, with the exception of polyvalent metal salts that hydrolyze in water, and many acidic salts.

Solubility of salts in water at room temperature

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 in either water or acids; G - completely hydrolyzes when dissolved 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 solutions of salts 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 crystal hydrates, from non-aqueous solutions - in the form of crystal 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 acidic oxide: MgO + CO 2 = MgCO 3

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

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

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

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

BaCl 2 + H 2 SO 4 = BaSO 4 + 2HCl

9) base solution with salt solution: Ba(OH) 2 + Na 2 SO 4 = 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 base. KOH + H2SO4 = KHSO4 + H2O

2. Interaction of the base with excess acid oxide

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

3. Interaction of the average salt with the acid Ca 3 (PO 4) 2 + 4H 3 PO 4 = 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 = Cl + HCl

2. Adding (drop by drop) 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 = 2 CO 3 + CO 2 + 4NaCl

4. Preparation of complex salts:

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

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

5. Preparation of double salts:

1. Joint crystallization of two salts:

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

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

2. Chemical properties of acid salts:

Thermal decomposition to form medium salt

Ca(HCO 3) 2 = CaCO 3 + CO 2 + H 2 O

Interaction with alkali. Getting medium salt.

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

3. Chemical properties of basic salts:

Thermal decomposition.

2 CO 3 = 2CuO + CO 2 + H 2 O

Interaction with acid: formation of medium 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 + K2S = CuS + 2KCl + 4NH3

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

K 2 + 6H 2 O = Cl 2 + 2KCl

5. Chemical properties of double salts:

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

2. Reduction: KCr(SO 4) 2 + 2H°(Zn, dil. H 2 SO 4) = 2CrSO 4 + H 2 SO 4 + K 2 SO 4

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

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

1. Main types of salts Borats

(oxoborates), salts of boric acids: metaboric HBO 2, orthoboric H3 BO 3 and polyboronic acids not isolated in a free state. Based on the number of boron atoms in the molecule, they are divided into mono-, di, tetra-, hexaborates, etc. Borates are also called by the acids that form them and by the number of moles of B 2 O 3 per 1 mole of the main oxide. Thus, various metaborates can be called monoborates if they contain the B(OH) 4 anion 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 ring anion (B 3 O 6) 3-.

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 frame polymerized ones. The latter are formed as a result of the elimination of water in hydrated borate molecules and the formation of bridging bonds through oxygen atoms; the process is sometimes accompanied by the cleavage of the B-O bond inside the polyanions. Polyanions can attach side groups - boron-oxygen tetrahedra or triangles, their dimers or extraneous anions.

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

Borates are colorless amorphous substances or crystals (mainly with a low-symmetric structure - monoclinic or orthorhombic). For anhydrous borates, melting temperatures range from 500 to 2000 °C; The highest melting points are alkali metaborates and ortho- and metaborates of alkaline earth metals. Most borates readily 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; The elimination of water due to OH groups coordinated around boron atoms occurs up to ~750°C. 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), and 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 SO 2 ;. Borates of alkaline earth and heavy metals interact with solutions of alkalis, carbonates and hydrocarbonates of alkali metals. Anhydrous borates are chemically more stable than hydrated borates. With some alcohols, in particular 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 VO 3 with metal oxides, hydroxides or carbonates; exchange reactions of alkali metal borates, most often Na, with salts of other metals; reaction of mutual transformation of poorly 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 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 charge components in the production of glass, glazes, enamels, ceramics; for fire-resistant coatings and impregnations; as components of fluxes for refining, welding and soldering metal”; as pigments and fillers for paints and varnishes; as dyeing mordants, 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 greater electronegativity than the other element. Halides are not formed by He, Ne and Ar. Simple, or binary, halides EXn (n is 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 (eg halofluorides). There are also mixed halides, polyhalides, hydrohalides, oxohalides, oxyhalides, hydroxohalides, thiohalides and complex halides. The oxidation number of halogens in halides is usually -1.

Based on the nature of the element-halogen bond, simple halides are divided into ionic and covalent. In reality, the connections are of a mixed nature with a predominance of the contribution of one or another component. 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 predominates. Most of them are relatively refractory, low-volatile, and highly soluble in water; in aqueous solutions almost completely dissociate into ions. Trihalides of rare earth elements also have the properties of salts. The solubility of ionic halides in water generally decreases from iodides to fluorides. Chlorides, bromides and iodides Ag + , Cu + , 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 charge of a metal to the radius of its ion leads to an increase in the covalent component of the bond, a decrease in solubility in water and the thermal stability of halides, an increase in volatility, an increase in oxidation, ability and tendency to hydrolysis. These dependencies are observed for metal halides of the same period and in a series of halides of the same metal. They can be easily observed using 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 CaCl2, 967 and 975°C for ScCl3, -24.1 and 136°C for TiCl4. 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 EXn compounds, with constant n, the bond covalency usually increases when going from fluorides to chlorides and decreases when going from the latter to bromides and iodides. So, for AlF 3 the sublimation temperature is 1280°C, AlC1 3 180°C, boiling point AlBr 3 254.8°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 series MFn and MC1n where M is a metal of one subgroup, the covalency of the bond decreases with increasing atomic mass of the metal. There are few metal fluorides and chlorides with approximately equal contributions from 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 (oxo- and oxyhalides, respectively), for example, vanadium oxotrifluoride VOF 3, niobium dioxyfluoride NbO 2 F, tungsten dioxo-iodide WO 2 I 2.

Complex halides (halometallates) contain complex anions in which the halogen atoms are ligands, for example, potassium hexachloroplatinate(IV) K2, sodium heptafluorotantalate(V), Na, lithium hexafluoroarsenate(V). Fluoro-, oxofluoro- and chlorometalates have the greatest thermal stability. By the nature of the bonds, ionic compounds with cations NF 4 +, N 2 F 3 +, C1F 2 +, XeF +, etc. are similar to complex halides.

Many halides are characterized by association and polymerization in the liquid and gas phases with the formation of bridging bonds. The most prone to this are metal halides of groups I and II, AlCl 3, pentafluorides of Sb and transition metals, oxofluorides of the composition MOF 4. Halides with a metal-to-metal bond are known, e.g. 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 halides.

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 = Ti + 2MgCl 2

UF 6 + H 2 = UF 4 + 2HF

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

WF 6 + ZN 2 = W + 6HF

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

KS1 + TaCl 5 = K

Lighter halogens can displace heavier halides. Oxygen can oxidize halides, releasing C1 2, Br 2, and I 2. One of the characteristic reactions of covalent halides is interaction with water (hydrolysis) or its vapor when heated (pyrohydrolysis), leading to the formation of oxides, oxy- or oxohalides, hydroxides and hydrogen halides.

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

Halides are widely used in technology as starting materials for the production of halogens, alkali and alkaline earth metals, 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 (for example, the minerals fluorite, cryolite) and chlorides (sylvite, carnallite). Bromine and iodine are present in some minerals as isomorphic impurities. Significant quantities of halides are contained in the water of seas and oceans, in salt and underground brines. Some halides, for example NaCl, KC1, CaCl 2, are part of living organisms.

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

Of the medium carbonates, only salts of alkali metals, ammonium and Tl(I) are soluble in water. As a result of significant hydrolysis, their solutions have an alkaline reaction. Metal carbonates are most difficult to dissolve 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 in cases where their solubility is significantly less than that of the corresponding hydroxides. This is the case for Ca, Sr and their analogs, the 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 intermediate, but basic crabonates or even hydroxides. Medium crabonates 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 with strong acids and the quantitative absorption of the resulting CO 2 by an alkali solution, or on the precipitation of the CO 3 2- ion from solution in the form of BaCO 3. When excess CO 2 acts on a medium carbonate precipitate, hydrogen carbonate is formed in solution, for example: CaCO 3 + H 2 O + CO 2 = Ca(HCO 3) 2. The presence of hydrocarbonates in natural water causes its temporary hardness. Hydrocarbonates, when slightly heated, even at low temperatures, again transform into medium carbonates, which, when heated, decompose to oxide and CO 2. The more active the metal, the higher the decomposition temperature of its carbonate. Thus, Na 2 CO 3 melts without decomposition at 857 °C, and for carbonates Ca, Mg and A1, 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 equilibria 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 calcite CaCO 3, magnesite MgCO 3, siderite FeCO 3, smithsonite ZnCO 3 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 (for example, MgCO 3 ZH 2 O, Na 2 CO 3 10H 2 O), double carbonates [for example, dolomite CaMg(CO 3) 2, trona Na 2 CO 3 NaHCO 3 2H 2 are also known O] and basic [malachite CuCO 3 Cu(OH) 2, hydrocerussite 2PbCO 3 Pb(OH) 2].

The most important are potassium carbonate, calcium carbonate and sodium carbonate. Many natural carbonates are very valuable metal ores (eg carbonates 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 is the oxidation state of the metal M) and in the form of crystalline hydrates M(NO 3)n xH 2 O (x = 1-9). Of aqueous solutions at temperatures close to room temperature, only alkali metal nitrates crystallize as 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 higher solubility in organic solvents, lower thermal stability, and their IR spectra are more complex; Some covalent nitrates are volatile at room temperature, and when dissolved in water, they partially decompose, releasing nitrogen oxides.

All anhydrous nitrates exhibit 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 ones - at 400-600°C (NaNO 3, KNO 3 and some others melt when heated). Decomposition products in solid and liquid phases. are successively nitrites, oxynitrates 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 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 quickly it can decompose with an explosion, in which 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 the 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 mainly retains its shape and size, which determines the space and structure of nitrates. The NO 3 - ion can act as a mono-, bi-, tridentate or bridging ligand, therefore nitrates are characterized by a wide variety of types of crystal structures.

Transition metals in high oxidation states due to steric. Anhydrous nitrates cannot form any difficulties, 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 internal 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 nitrates in that in their crystal structures the metal ion is in most cases associated with water molecules rather than with the NO 3 ion. Therefore, they are better soluble in water than anhydrous nitrates, but less soluble in organic solvents; they are weaker oxidizing agents and melt incongruently in water of crystallization in the range of 25-100°C. When hydrated nitrates are heated, anhydrous nitrates, as a rule, are not formed, but thermolysis occurs with the formation of hydroxonitrates and then oxonitrate 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. When nitrates are reduced, 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 environment and other factors.

Industrial methods for producing nitrates are based on the absorption of NH 3 by solutions of HNO 3 (for NH 4 NO 3) or on the absorption of nitrous gases (NO + NO 2) by solutions of alkalis or carbonates (for alkali metal nitrates, Ca, Mg, Ba), as well as 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) is the main nitrogen-containing fertilizer; Alkali metal nitrates and Ca are also used as fertilizers. Nitrates are components of rocket fuels, pyrotechnic compositions, etching 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 failure, etc. The lethal dose of nitrates for humans is 8-15 g, permissible daily intake is 5 mg/kg. For the sum of nitrates Na, K, Ca, NH3 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. Failure to comply with agrotechnical recommendations, excessive application of fertilizers sharply increases the nitrate content in agricultural products, surface runoff from fields ( 40-5500 mg/l), groundwater.

5. Nitrites, salts of nitrous acid HNO 2. Nitrites of alkali metals and ammonium are used primarily, and less - of alkaline earth and Zd metals, Pb and Ag. There is only fragmentary information about 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, e.g. CsNO 2 AgNO 2 or Ba(NO 2) 2 Ni(NO 2) 2 2KNO 2, as well as complex compounds, for example Na 3.

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

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

Nitrites are thermally unstable; Only nitrites of alkali metals melt without decomposition; 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 - metal oxide or elemental metal. The release of large amounts 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 environment 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 promote the decomposition of nitrogen-containing organic substances, in particular amines, amides, etc. With organic halides RXH. react to form both nitrites RONO and nitro compounds RNO 2 .

The industrial production of nitrites is based on the absorption of nitrous gas (mixture of NO + NO 2) with solutions of Na 2 CO 3 or NaOH with sequential crystallization of NaNO 2; Nitrites of other metals are obtained in industry and laboratories 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 agents and reducing agents in the rubber, textile and metalworking industries, as food preservatives. Nitrites, such as NaNO 2 and KNO 2, are toxic, causing headaches, vomiting, depressing breathing, etc. When NaNO 2 is poisoned, methemoglobin is formed in the blood and red blood cell membranes are damaged. It is possible to form nitrosamines from NaNO 2 and amines directly in the gastrointestinal tract.

6. Sulfates, salts of sulfuric acid. Medium sulfates with the SO 4 2- anion are known, or hydrosulfates, with the HSO 4 - anion, basic, containing, along with the SO 4 2- anion, OH groups, for example Zn 2 (OH) 2 SO 4. There are also double sulfates containing two different cations. These include two large groups of sulfates - alum, as well as schenites 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 (polyhalite mineral), 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 berkeite), MgSO 4 KCl 3H 2 O (kainite) .

Sulfates are crystalline substances, medium and acidic in most cases, highly soluble in water. Sulfates of calcium, strontium, lead and some others are slightly soluble; BaSO 4 and RaSO 4 are practically insoluble. Basic sulfates are usually poorly soluble or practically insoluble, or are hydrolyzed by water. From aqueous solutions, sulfates can crystallize in the form of crystalline hydrates. Crystal hydrates of some heavy metals are called vitriols; copper sulfate CuSO 4 5H 2 O, iron sulfate FeSO 4 7H 2 O.

Medium alkali metal sulfates are thermally stable, while acid sulfates decompose when heated, turning into pyrosulfates: 2KHSO 4 = H 2 O + K 2 S 2 O 7. Medium 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, for example, 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 volatile acid salts with sulfuric acid.

Inorganic sulfates are widely used. For example, ammonium sulfate is a nitrogen fertilizer, sodium sulfate is used in the glass, paper industries, 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 SO 3 2- anion and acidic (hydrosulfites) with the HSO 3 - anion. 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). Hydrosulfites are formed in aqueous solutions. 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 crystal 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, are disproportionately divided into sulfides and sulfates; when heated in a current of N 2, they lose SO 2, and when heated in air, they are easily oxidized to sulfates. With SO 2 in an aqueous 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 decompose with strong acids (for example, HC1) with the release of SO 2.

Crystalline hydrosulfites are known for K, Rb, Cs, NH 4 +, they are unstable. The remaining hydrosulfites exist only in aqueous solutions. Density of NH 4 HSO 3 2.03 g/cm 3 ; solubility in water (g in 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 teeming pulp solution is saturated with SO 2 M 2 SO 3, pyrosulfites (obsolete - metabisulfites) M 2 S 2 O 5 are formed - salts of the unknown free pyrosulfuric acid H 2 S 2 O 5; crystals, unstable; density (g/cm3): Na 2 S 2 O 5 1.48, K 2 S 2 O 5 2.34; above ~ 160 °C they decompose with the release of SO 2; dissolve in water (with decomposition to HSO 3 -), solubility (g in 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 prepared 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; They mainly use SO 2 from the exhaust gases of contact sulfuric acid production. Sulfites are used in bleaching, dyeing and printing of fabrics, fibers, leather for grain conservation, green feed, feed industrial waste (NaHSO 3,

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

Methods for separating mixtures

1. Filtration, separation of heterogeneous systems liquid - solid particles (suspensions) and gas - solid particles using porous filter partitions (FP), which 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 phase transition.

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 per m 3 passing through 1 m 2 of the FP surface per unit time). At a constant pressure difference, the suspension is supplied to the filter under vacuum or excess pressure, 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 layer of sediment does not form on the FP and solid particles enter the filtrate, filter using finely dispersed auxiliary materials (diatomaceous earth, 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.

There are continuous and periodic filters. For the latter, the main stages of work are filtering, washing the sediment, its dewatering and unloading. In this case, optimization according to the criteria of greatest productivity and lowest costs is applicable. If washing and dewatering are not carried out, and the hydraulic resistance of the partition can be neglected, then the greatest productivity is achieved when the filtering time is equal to the duration of the auxiliary operations.

Flexible FPs made from cotton, wool, synthetic and glass fabrics are applicable, as well as non-woven FPs made from natural and synthetic fibers and inflexible ones - ceramic, cermet and foam. 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 (see figure) of continuous action, in which the directions of movement of the filtrate and the action of gravity are opposite. The distribution device section connects zones I and II with a vacuum source and zones III and IV with a compressed air source. The filtrate and washing liquid from zones I and II enter separate receivers. An automated periodic filter press with horizontal chambers, filter fabric in the form of an endless belt and elastic membranes for dewatering sludge by pressing has also become widespread. It performs alternating operations of filling chambers with suspension, filtering, washing and dewatering sediment, disconnecting adjacent chambers and removing sediment.

2.Fractional crystallization

The following types of fractional crystallization are distinguished: mass, on cooled surfaces, directional, zone melting.

Mass crystallization. The method consists of simultaneously obtaining a large number of crystals throughout the entire volume of the apparatus. The industry has implemented several options for mass crystallization, which is carried out in periodically or continuously operating apparatuses: capacitive ones, equipped with external cooling jackets or internal coils and often mixing devices; tubular, scraper, disk, screw, etc. Due to the lack of calculation methods, the parameter a e during mass crystallization is found experimentally.

Crystallization with heat transfer through the wall. In the case of melts, the process is carried out by cooling them. During the crystallization of solutions, the choice of process mode is determined mainly by the nature of the dependence of the solubility of substances on temperature. If the solubility of a substance changes little with temperature (for example, NaCI in water), crystallization is carried out by partial or almost complete evaporation of a saturated solution at a constant temperature (isothermal crystallization). Substances whose solubility strongly depends on temperature (for example, KNO 3 in water) are crystallized by cooling hot solutions, while the initial amount of solvent contained in the mother liquid does not change in the system (isohydric crystallization). The resulting crystals, depending on their properties, shape and process conditions, capture different amounts of mother liquor. Its content in the solid phase in the form of inclusions in pores, cracks and cavities significantly depends on the method of separation of crystals and mother liquor. So, when separating crystals on a drum vacuum filter, the concentration of the mother solution in them is 10-30%, on a filter centrifuge - 3-10%.

The main advantages of the process: high productivity, lack of contact between the mixture being separated and the refrigerant, simplicity of hardware design; disadvantages: relatively low heat transfer coefficient, encrustation of cooling surfaces, large capture of mother liquor by crystals, the need to install additional equipment for separating solid and liquid phases, insufficiently high yield of crystalline product. Application examples: preparation of K and Na chlorides from sylvinite, separation of xylene isomers.

3. Evaporation is carried out to concentrate the solution, isolate the dissolved substance or obtain a pure solvent. Mainly aqueous solutions are subjected to evaporation. The coolant most often is water vapor (pressure 1.0-1.2 MPa), which is called heating, or primary; The steam formed when the solution boils is called secondary. The driving force for evaporation, the temperature difference between the heating steam and the boiling solution, is called useful. It is always less than the temperature difference between the primary and secondary steam. This is due to the fact that the solution boils at a higher temperature than the pure solvent (physicochemical, or concentration, depression). In addition, the boiling point of the solution increases due to the higher pressure in the solution than in the vapor space. Reasons for the increase in pressure: hydrostatic pressure of the solution; hydraulic resistance when moving a boiling (vapor-liquid) mixture; an increase in the speed of movement of this mixture due to the fact that it occupies a significantly larger volume than the original solution (hydrostatic, hydraulic and inertial depression, respectively).

For evaporation, devices operating under pressure or vacuum are used. Their main elements: heating chamber; a separator for separating the vapor-liquid mixture into collecting a concentrated solution; a circulation pipe through which the solution returns from the separator to the chamber (with repeated evaporation). The design of the apparatus is determined mainly by the composition, physicochemical properties, the required degree of concentration of solutions, their tendency to form scale and foam (scale sharply reduces the heat transfer coefficient, disrupts the circulation of the solution and can cause corrosion in welded joints, and heavy pricing increases the carryover of the solution by secondary ferry).

The most common are vertical devices with tubular heating chambers, the heating surface of which reaches 1250 m 2. In such devices, the solution is in the pipe, and the heating steam is in the inter-tube space of the chamber. The circulation of the solution in them can be natural or forced, created by a special pump.

Evaporation of low-viscosity (l up to 6-8 mPa -s) unsaturated solutions of highly soluble salts that do not precipitate during concentration (for example, NaNO 2, NaNO 3, NH 4 NO 3, KC1) and do not form scale is usually carried out in evaporating apparatus with natural circulation, in the heating tubes of which the solution not only heats up, but also boils. To evaporate solutions of poorly soluble substances that precipitate upon concentration [for example, CaCO 3, CaSO 4, Mg(OH) 2, Na aluminosilicate], as well as for the desalination of sea water, apparatus is used, above the heating chamber of which an additional lifting circulation system is installed a pipe that provides high speed natural circulation. To evaporate highly foaming and heat-sensitive products, for example in the production of yeast, enzymes, antibiotics, fruit juices, instant coffee, vertical film evaporators are used, in which concentration occurs as a result of a single movement of a thin layer (film) of solution together with secondary steam along tubes of length 6-8 m (heating surface up to 2200 m2). The advantages of these devices: absence of hydrostatic effect, low hydraulic resistance, high heat transfer coefficient, high productivity with relatively small volumes

4. Centrifugation, separation of suspensions, emulsions and three-component systems (emulsions containing a solid phase) under the action of centrifugal forces. It is used for isolating fractions from suspensions and emulsions, as well as for determining the molecular weights of polymers and dispersion analysis.

Centrifugation is carried out using special machines - centrifuges, the main part of which is a rotor (drum) rotating at high speed around its axis, thereby creating a field of centrifugal forces of up to 20,000 g in industrial centrifuges and up to 350,000 g in laboratory centrifuges (g - acceleration free fall). Centrifugation can be carried out according to the principles of sedimentation or filtration, respectively, in centrifuges with a solid or perforated rotor covered with filter material. There are two types of centrifuges: 1) periodic, in which the suspension is introduced into the center, part of the hollow rotor during its rotation; solid particles settle on the inner surface of the rotor and are discharged from it through a special nozzles or through periodically opening slits, clarified liquid (centrate) is removed from the top of its part; 2) continuous action, in which the suspension is fed along the axis of a hollow rotor, and the resulting sediment is unloaded using a screw rotating inside the rotor at a slightly different speed than the rotor (Fig. 1).

Centrifugation based on the filtration principle is most often used to separate suspensions and sludges with a relatively low content of the liquid phase and is carried out in cyclically operating machines. The suspension is fed into a continuously rotating rotor in portions; After part of the rotor is filled with sediment, the supply of suspension is stopped, the liquid phase is squeezed out, and the sediment is cut off with a knife and removed. Centrifuges with pulsating sediment unloading using a pusher (vibrating piston, with a pulsating piston) are also used, as well as with hydraulic unloading, when the condensed solid phase is removed from a rotor equipped with a package of conical plates through nozzles.

Bibliography

Ch. editor I.L. Knunyants. Large encyclopedic dictionary of Chemistry. Moscow 1998

Ch. editor I.L. Knunyants. Chemical encyclopedia. Moscow1998

N. Ya. Loginov, A. G. Voskresensky, I. S. Solodin. Analytical chemistry. Moscow 1979

R. A. Lidin. Handbook of general and inorganic chemistry. Moscow 1997

R. A. Lidin, V. A. Molochko, L. L. Andreeva. Chemical properties of inorganic substances. Moscow 1997

A. V. Suvorov, A. A. Kartsafa and others. The fascinating world of chemical transformations. St. Petersburg 1998

E. V. Barkovsky. Introduction to the chemistry of biogenic elements and chemical analysis. Minsk 1997

Salts are chemical compounds in which a metal atom is bonded to an acidic moiety. The difference between salts and other compounds is that they have a clearly expressed ionic bond. That is why the bond is called ionic. Ionic bonding is characterized by unsaturation and non-directionality. Examples of salts: sodium chloride or kitchen salt - NaCl, calcium sulfate or gypsum - CaSO4. Depending on how completely the hydrogen atoms in the acid or the hydroxo groups in the hydroxide are replaced, medium, acidic and basic salts are distinguished. A salt may contain several metal cations – these are double salts.

Medium salts

Medium salts are salts in which hydrogen atoms are completely replaced by metal ions. Kitchen salt and gypsum are such salts. Medium salts cover a large number of compounds often found in nature, for example, blende - ZnS, pyrrite - FeS2, etc. This type of salt is the most common.

Medium salts are obtained by a neutralization reaction when the base is taken in equimolar ratios, for example:
H2SO3 + 2 NaOH = Na2SO3 + 2 H2O
The result is medium salt. If you take 1 mole of sodium hydroxide, the reaction will proceed as follows:
H2SO3 + NaOH = NaHSO3 + H2O
The result is the acidic salt sodium hydrosulfite.

Acid salts

Acid salts are salts in which not all hydrogen atoms are replaced by a metal. Such salts are capable of forming only polybasic acids - sulfuric, phosphoric, sulfurous and others. Monobasic acids, such as hydrochloric, nitric and others, do not give.
Examples of salts: sodium bicarbonate or baking soda - NaHCO3, sodium dihydrogen phosphate - NaH2PO4.

Acid salts can also be obtained from medium salts with acid:
Na2SO3+ H2SO3 = 2NaHSO3

Basic salts

Basic salts are salts in which not all hydroxo groups are replaced by acidic residues. For example, – Al(OH)SO4, hydroxochloride – Zn(OH)Cl, copper dihydroxocarbonate or malachite – Cu2(CO3)(OH)2.

Double salts

Double salts are salts in which two metals replace hydrogen atoms in the acid moiety. Such salts are possible for polybasic acids. Examples of salts: potassium sodium carbonate - NaKCO3, potassium sulfate - KAl(SO4)2.. The most common double salts in everyday life are alum, for example, potassium alum - KAl(SO4)2 12H2O. They are used for water purification, leather tanning, and for loosening dough.

Mixed salts

Mixed salts are salts in which a metal atom is bonded to two different acidic residues, for example, bleach - Ca(OCl)Cl.

Reasons – complex substances that consist of a metal cation Me + (or a metal-like cation, for example, ammonium ion NH 4 +) and a hydroxide anion OH -.

Based on their solubility in water, bases are divided into soluble (alkalis) And insoluble bases . There is also unstable foundations, which spontaneously decompose.

Getting grounds

1. Interaction of basic oxides with water. In this case, only those oxides that correspond to a soluble base (alkali). Those. in this way you can only get alkalis:

basic oxide + water = base

For example , sodium oxide forms in water sodium hydroxide(sodium hydroxide):

Na 2 O + H 2 O → 2NaOH

At the same time about copper(II) oxide With water does not react:

CuO + H 2 O ≠

2. Interaction of metals with water. Wherein react with waterunder normal conditionsonly alkali metals(lithium, sodium, potassium, rubidium, cesium), calcium, strontium and barium.In this case, a redox reaction occurs, hydrogen is the oxidizing agent, and the metal is the reducing agent.

metal + water = alkali + hydrogen

For example, potassium reacts with water very stormy:

2K 0 + 2H 2 + O → 2K + OH + H 2 0

3. Electrolysis of solutions of some alkali metal salts. As a rule, to obtain alkalis, electrolysis is carried out solutions of salts formed by alkali or alkaline earth metals and oxygen-free acids (except for hydrofluoric acid) - chlorides, bromides, sulfides, etc. This issue is discussed in more detail in the article .

For example , electrolysis of sodium chloride:

2NaCl + 2H 2 O → 2NaOH + H 2 + Cl 2

4. Bases are formed by the interaction of other alkalis with salts. In this case, only soluble substances interact, and an insoluble salt or an insoluble base should be formed in the products:

or

alkali + salt 1 = salt 2 ↓ + alkali

For example: Potassium carbonate reacts in solution with calcium hydroxide:

K 2 CO 3 + Ca(OH) 2 → CaCO 3 ↓ + 2KOH

For example: Copper(II) chloride reacts in solution with sodium hydroxide. In this case it falls out blue copper(II) hydroxide precipitate:

CuCl 2 + 2NaOH → Cu(OH) 2 ↓ + 2NaCl

Chemical properties of insoluble bases

1. Insoluble bases react with strong acids and their oxides (and some medium acids). In this case, salt and water.

insoluble base + acid = salt + water

insoluble base + acid oxide = salt + water

For example ,Copper(II) hydroxide reacts with strong hydrochloric acid:

Cu(OH) 2 + 2HCl = CuCl 2 + 2H 2 O

In this case, copper (II) hydroxide does not interact with the acid oxide weak carbonic acid - carbon dioxide:

Cu(OH) 2 + CO 2 ≠

2. Insoluble bases decompose when heated into oxide and water.

For example, Iron(III) hydroxide decomposes into iron(III) oxide and water when heated:

2Fe(OH) 3 = Fe 2 O 3 + 3H 2 O

3. Insoluble bases do not reactwith amphoteric oxides and hydroxides.

insoluble base + amphoteric oxide ≠

insoluble base + amphoteric hydroxide ≠

4. Some insoluble bases can act asreducing agents. Reducing agents are bases formed by metals with minimum or intermediate oxidation state, which can increase their oxidation state (iron (II) hydroxide, chromium (II) hydroxide, etc.).

For example , Iron (II) hydroxide can be oxidized with atmospheric oxygen in the presence of water to iron (III) hydroxide:

4Fe +2 (OH) 2 + O 2 0 + 2H 2 O → 4Fe +3 (O -2 H) 3

Chemical properties of alkalis

1. Alkalis react with any acids - both strong and weak . In this case, medium salt and water are formed. These reactions are called neutralization reactions. Education is also possible sour salt, if the acid is polybasic, at a certain ratio of reagents, or in excess acid. IN excess alkali medium salt and water are formed:

alkali (excess) + acid = medium salt + water

alkali + polybasic acid (excess) = acid salt + water

For example , Sodium hydroxide, when interacting with tribasic phosphoric acid, can form 3 types of salts: dihydrogen phosphates, phosphates or hydrophosphates.

In this case, dihydrogen phosphates are formed in an excess of acid, or when the molar ratio (ratio of the amounts of substances) of the reagents is 1:1.

NaOH + H 3 PO 4 → NaH 2 PO 4 + H 2 O

When the molar ratio of alkali and acid is 2:1, hydrophosphates are formed:

2NaOH + H3PO4 → Na2HPO4 + 2H2O

In an excess of alkali, or with a molar ratio of alkali to acid of 3:1, alkali metal phosphate is formed.

3NaOH + H3PO4 → Na3PO4 + 3H2O

2. Alkalis react withamphoteric oxides and hydroxides. Wherein ordinary salts are formed in the melt , A in solution - complex salts .

alkali (melt) + amphoteric oxide = medium salt + water

alkali (melt) + amphoteric hydroxide = medium salt + water

alkali (solution) + amphoteric oxide = complex salt

alkali (solution) + amphoteric hydroxide = complex salt

For example , when aluminum hydroxide reacts with sodium hydroxide in the melt sodium aluminate is formed. A more acidic hydroxide forms an acidic residue:

NaOH + Al(OH) 3 = NaAlO 2 + 2H 2 O

A in solution a complex salt is formed:

NaOH + Al(OH) 3 = Na

Please note how the complex salt formula is composed:first we select the central atom (toAs a rule, it is an amphoteric hydroxide metal).Then we add to it ligands- in our case these are hydroxide ions. The number of ligands is usually 2 times greater than the oxidation state of the central atom. But the aluminum complex is an exception; its number of ligands is most often 4. We enclose the resulting fragment in square brackets - this is a complex ion. We determine its charge and add the required number of cations or anions on the outside.

3. Alkalis interact with acidic oxides. At the same time, education is possible sour or medium salt, depending on the molar ratio of alkali and acid oxide. In an excess of alkali, a medium salt is formed, and in an excess of an acidic oxide, an acid salt is formed:

alkali (excess) + acid oxide = medium salt + water

or:

alkali + acid oxide (excess) = acid salt

For example , when interacting excess sodium hydroxide With carbon dioxide, sodium carbonate and water are formed:

2NaOH + CO 2 = Na 2 CO 3 + H 2 O

And when interacting excess carbon dioxide with sodium hydroxide only sodium bicarbonate is formed:

2NaOH + CO 2 = NaHCO 3

4. Alkalis interact with salts. Alkalis react only with soluble salts in solution, provided that Gas or sediment forms in the food . Such reactions proceed according to the mechanism ion exchange.

alkali + soluble salt = salt + corresponding hydroxide

Alkalies interact with solutions of metal salts, which correspond to insoluble or unstable hydroxides.

For example, sodium hydroxide reacts with copper sulfate in solution:

Cu 2+ SO 4 2- + 2Na + OH - = Cu 2+ (OH) 2 - ↓ + Na 2 + SO 4 2-

Also alkalis react with solutions of ammonium salts.

For example , Potassium hydroxide reacts with ammonium nitrate solution:

NH 4 + NO 3 - + K + OH - = K + NO 3 - + NH 3 + H 2 O

! When salts of amphoteric metals interact with excess alkali, a complex salt is formed!

Let's look at this issue in more detail. If the salt formed by the metal to which it corresponds amphoteric hydroxide , interacts with a small amount of alkali, then the usual exchange reaction occurs, and a precipitate occurshydroxide of this metal .

For example , excess zinc sulfate reacts in solution with potassium hydroxide:

ZnSO 4 + 2KOH = Zn(OH) 2 ↓ + K 2 SO 4

However, in this reaction it is not a base that is formed, but mphoteric hydroxide. And, as we already indicated above, amphoteric hydroxides dissolve in excess alkalis to form complex salts . T Thus, when zinc sulfate reacts with excess alkali solution a complex salt is formed, no precipitate forms:

ZnSO 4 + 4KOH = K 2 + K 2 SO 4

Thus, we obtain 2 schemes for the interaction of metal salts, which correspond to amphoteric hydroxides, with alkalis:

amphoteric metal salt (excess) + alkali = amphoteric hydroxide↓ + salt

amph.metal salt + alkali (excess) = complex salt + salt

5. Alkalis interact with acidic salts.In this case, medium salts or less acidic salts are formed.

sour salt + alkali = medium salt + water

For example , Potassium hydrosulfite reacts with potassium hydroxide to form potassium sulfite and water:

KHSO 3 + KOH = K 2 SO 3 + H 2 O

It is very convenient to determine the properties of acidic salts by mentally breaking the acidic salt into 2 substances - acid and salt. For example, we break sodium bicarbonate NaHCO 3 into uolic acid H 2 CO 3 and sodium carbonate Na 2 CO 3. The properties of bicarbonate are largely determined by the properties of carbonic acid and the properties of sodium carbonate.

6. Alkalis interact with metals in solution and melt. In this case, an oxidation-reduction reaction occurs, forming in the solution complex salt And hydrogen, in the melt - medium salt And hydrogen.

Note! Only those metals whose oxide with the minimum positive oxidation state of the metal is amphoteric react with alkalis in solution!

For example , iron does not react with alkali solution, iron (II) oxide is basic. A aluminum dissolves in aqueous alkali solution, aluminum oxide is amphoteric:

2Al + 2NaOH + 6H 2 + O = 2Na + 3H 2 0

7. Alkalies interact with non-metals. In this case, redox reactions occur. Usually, nonmetals are disproportionate in alkalis. They don't react with alkalis oxygen, hydrogen, nitrogen, carbon and inert gases (helium, neon, argon, etc.):

NaOH +O 2 ≠

NaOH +N 2 ≠

NaOH +C ≠

Sulfur, chlorine, bromine, iodine, phosphorus and other non-metals disproportionate in alkalis (i.e. they self-oxidize and self-recover).

For example, chlorinewhen interacting with cold lye goes into oxidation states -1 and +1:

2NaOH +Cl 2 0 = NaCl - + NaOCl + + H 2 O

Chlorine when interacting with hot lye goes into oxidation states -1 and +5:

6NaOH +Cl 2 0 = 5NaCl - + NaCl +5 O 3 + 3H 2 O

Silicon oxidized by alkalis to oxidation state +4.

For example, in solution:

2NaOH + Si 0 + H 2 + O= NaCl - + Na 2 Si +4 O 3 + 2H 2 0

Fluorine oxidizes alkalis:

2F 2 0 + 4NaO -2 H = O 2 0 + 4NaF - + 2H 2 O

You can read more about these reactions in the article.

8. Alkalis do not decompose when heated.

The exception is lithium hydroxide:

2LiOH = Li 2 O + H 2 O

1. Bases react with acids to form salt and water:

Cu(OH) 2 + 2HCl = CuCl 2 + 2H 2 O

2. With acid oxides, forming salt and water:

Ca(OH) 2 + CO 2 = CaCO 3 + H 2 O

3. Alkalis react with amphoteric oxides and hydroxides, forming salt and water:

2NaOH + Cr 2 O 3 = 2NaCrO 2 + H 2 O

KOH + Cr(OH) 3 = KCrO 2 + 2H 2 O

4. Alkalis react with soluble salts, forming either a weak base, a precipitate, or a gas:

2NaOH + NiCl 2 = Ni(OH) 2 ¯ + 2NaCl

base

2KOH + (NH 4) 2 SO 4 = 2NH 3 + 2H 2 O + K 2 SO 4

Ba(OH) 2 + Na 2 CO 3 = BaCO 3 ¯ + 2NaOH

5. Alkalis react with some metals, which correspond to amphoteric oxides:

2NaOH + 2Al + 6H 2 O = 2Na + 3H 2

6. Effect of alkali on the indicator:

OH - + phenolphthalein ® crimson color

OH - + litmus ® blue color

7. Decomposition of some bases when heated:

Сu(OH) 2 ® CuO + H 2 O

Amphoteric hydroxides– chemical compounds exhibiting the properties of both bases and acids. Amphoteric hydroxides correspond to amphoteric oxides (see paragraph 3.1).

Amphoteric hydroxides are usually written in the form of a base, but they can also be represented in the form of an acid:

Zn(OH) 2 Û H 2 ZnO 2

foundation

Chemical properties of amphoteric hydroxides

1. Amphoteric hydroxides interact with acids and acid oxides:

Be(OH) 2 + 2HCl = BeCl 2 + 2H 2 O

Be(OH) 2 + SO 3 = BeSO 4 + H 2 O

2. Interact with alkalis and basic oxides of alkali and alkaline earth metals:

Al(OH) 3 + NaOH = NaAlO 2 + 2H 2 O;

H 3 AlO 3 acid sodium metaaluminate

(H 3 AlO 3 ® HAlO 2 + H 2 O)

2Al(OH) 3 + Na 2 O = 2NaAlO 2 + 3H 2 O

All amphoteric hydroxides are weak electrolytes

Salts

Salts- These are complex substances consisting of metal ions and an acid residue. Salts are products of complete or partial replacement of hydrogen ions with metal (or ammonium) ions in acids. Types of salts: medium (normal), acidic and basic.

Medium salts- these are the products of complete replacement of hydrogen cations in acids with metal (or ammonium) ions: Na 2 CO 3, NiSO 4, NH 4 Cl, etc.

Chemical properties of medium salts

1. Salts interact with acids, alkalis and other salts, forming either a weak electrolyte or a precipitate; or gas:

Ba(NO 3) 2 + H 2 SO 4 = BaSO 4 ¯ + 2HNO 3

Na 2 SO 4 + Ba(OH) 2 = BaSO 4 ¯ + 2NaOH

CaCl 2 + 2AgNO 3 = 2AgCl¯ + Ca(NO 3) 2

2CH 3 COONa + H 2 SO 4 = Na 2 SO 4 + 2CH 3 COOH

NiSO 4 + 2KOH = Ni(OH) 2 ¯ + K 2 SO 4

base

NH 4 NO 3 + NaOH = NH 3 + H 2 O + NaNO 3

2. Salts interact with more active metals. A more active metal displaces a less active metal from the salt solution (Appendix 3).

Zn + CuSO 4 = ZnSO 4 + Cu

Acid salts- these are products of incomplete replacement of hydrogen cations in acids with metal (or ammonium) ions: NaHCO 3, NaH 2 PO 4, Na 2 HPO 4, etc. Acid salts can only be formed by polybasic acids. Almost all acid salts are highly soluble in water.

Obtaining acidic salts and converting them to medium salts

1. Acid salts are obtained by reacting an excess of acid or acid oxide with a base:

H 2 CO 3 + NaOH = NaHCO 3 + H 2 O

CO 2 + NaOH = NaHCO 3

2. When excess acid interacts with the basic oxide:

2H 2 CO 3 + CaO = Ca(HCO 3) 2 + H 2 O

3. Acid salts are obtained from medium salts by adding acid:

· eponymous

Na 2 SO 3 + H 2 SO 3 = 2NaHSO 3;

Na 2 SO 3 + HCl = NaHSO 3 + NaCl

4. Acid salts are converted to medium salts using alkali:

NaHCO 3 + NaOH = Na 2 CO 3 + H 2 O

Basic salts– these are products of incomplete substitution of hydroxo groups (OH - ) bases with an acidic residue: MgOHCl, AlOHSO 4, etc. Basic salts can only be formed by weak bases of polyvalent metals. These salts are generally sparingly soluble.

Obtaining basic salts and converting them to medium salts

1. Basic salts are obtained by reacting an excess of base with an acid or acid oxide:

Mg(OH) 2 + HCl = MgOHCl¯ + H 2 O

hydroxo-

magnesium chloride

Fe(OH) 3 + SO 3 = FeOHSO 4 ¯ + H 2 O

hydroxo-

iron(III) sulfate

2. Basic salts are formed from medium salt by adding a lack of alkali:

Fe 2 (SO 4) 3 + 2NaOH = 2FeOHSO 4 + Na 2 SO 4

3. Basic salts are converted to medium salts by adding an acid (preferably the one that corresponds to the salt):

MgOHCl + HCl = MgCl 2 + H 2 O

2MgOHCl + H 2 SO 4 = MgCl 2 + MgSO 4 + 2H 2 O

ELECTROLYTES

Electrolytes- these are substances that disintegrate into ions in solution under the influence of polar solvent molecules (H 2 O). Based on their ability to dissociate (break down into ions), electrolytes are conventionally divided into strong and weak. Strong electrolytes dissociate almost completely (in dilute solutions), while weak electrolytes dissociate into ions only partially.

Strong electrolytes include:

· strong acids (see p. 20);

· strong bases – alkalis (see p. 22);

· almost all soluble salts.

Weak electrolytes include:

weak acids (see p. 20);

· bases are not alkali;

One of the main characteristics of a weak electrolyte is dissociation constant – TO . For example, for a monobasic acid,

HA Û H + +A - ,

where, is the equilibrium concentration of H + ions;

– equilibrium concentration of acid anions A - ;

– equilibrium concentration of acid molecules,

Or for a weak foundation,

MOH Û M + +OH - ,

,

where, is the equilibrium concentration of M + cations;

– equilibrium concentration of hydroxide ions OH - ;

– equilibrium concentration of weak base molecules.

Dissociation constants of some weak electrolytes (at t = 25°C)