salts complex substances are called, the molecules of which consist of metal atoms and acid residues (sometimes they may contain hydrogen). For example, NaCl is sodium chloride, CaSO 4 is calcium sulfate, etc.
Practically All salts are ionic compounds therefore, in salts, ions of acid residues and metal ions are interconnected:
Na + Cl - - sodium chloride
Ca 2+ SO 4 2– - calcium sulfate, etc.
Salt is a product of partial or complete replacement of acid hydrogen atoms by a metal. Hence, the following types of salts are distinguished:
1. Medium salts- all hydrogen atoms in the acid are replaced by a metal: Na 2 CO 3, KNO 3, etc.
2. Acid salts- not all hydrogen atoms in the acid are replaced by a metal. Of course, acid salts can only form dibasic or polybasic acids. Monobasic acids cannot give acid salts: NaHCO 3, NaH 2 PO 4, etc. d.
3. Double salts- hydrogen atoms of a dibasic or polybasic acid are replaced not by one metal, but by two different ones: NaKCO 3, KAl(SO 4) 2, etc.
4. Basic salts can be considered as products of incomplete or partial substitution of hydroxyl groups of bases by acidic residues: Al(OH)SO 4 , Zn(OH)Cl, etc.
According to international nomenclature, the name of the salt of each acid comes from the Latin name of the element. For example, salts of sulfuric acid are called sulfates: CaSO 4 - calcium sulfate, Mg SO 4 - magnesium sulfate, etc.; salts of hydrochloric acid are called chlorides: NaCl - sodium chloride, ZnCI 2 - zinc chloride, etc.
The particle "bi" or "hydro" is added to the name of salts of dibasic acids: Mg (HCl 3) 2 - magnesium bicarbonate or bicarbonate.
Provided that in a tribasic acid only one hydrogen atom is replaced by a metal, then the prefix "dihydro" is added: NaH 2 PO 4 - sodium dihydrogen phosphate.
Salts are solid substances that have a wide range of solubility in water.
Chemical properties of salts
The chemical properties of salts are determined by the properties of the cations and anions that are part of their composition.
1. Some salts decompose when calcined:
CaCO 3 \u003d CaO + CO 2
2. React with acids to form a new salt and a new acid. For this reaction to occur, it is necessary that the acid be stronger than the salt that the acid acts on:
2NaCl + H 2 SO 4 → Na 2 SO 4 + 2HCl.
3. Interact with bases, forming a new salt and a new base:
Ba(OH) 2 + MgSO 4 → BaSO 4 ↓ + Mg(OH) 2 .
4. Interact with each other with the formation of new salts:
NaCl + AgNO 3 → AgCl + NaNO 3 .
5. Interact with metals, which are in the range of activity to the metal that is part of the salt:
Fe + CuSO 4 → FeSO 4 + Cu↓.
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DEFINITION
Salts are electrolytes, the dissociation of which forms metal cations (ammonium ion or complex ions) and anions of acid residues:
\(\ \mathrm(NaNOZ) \mapsto \mathrm(Na)++\mathrm(NOZ)_(-) \);
\(\ \mathrm(NH) 4 \mathrm(NO) 3 \leftrightarrow \mathrm(NH) 4++\mathrm(NO) 3_(-) \);
\(\ \mathrm(KAl)(\mathrm(SO) 4) 2 \leftrightarrow \mathrm(K)++\mathrm(Al) 3++2 \mathrm(SO) 42- \);
\(\ [\mathrm(Zn)(\mathrm(NH) 3) 4] \mathrm(Cl) 2[\mathrm(Zn)(\mathrm(NH) 3) 4] 2++2 \mathrm(Cl) \).
Salts are usually divided into three groups: medium (\(\ \mathrm(NaCl) \)), acidic (\(\ \mathrm(NaHCO) 3 \)) and basic (\(\ \mathrm(Fe)(\mathrm( OH)) \mathrm(Cl) \)). In addition, there are double (mixed) and complex salts. Double salts are formed by two cations and one anion. They exist only in solid form.
Chemical properties of salts
a) acid salts
Acid salts during dissociation give metal cations (ammonium ion), hydrogen ions and anions of the acid residue:
\(\ \mathrm(NaHCO) 3+\mathrm(Na)++\mathrm(H)++\mathrm(CO) 32 \).
Acid salts are products of incomplete replacement of hydrogen atoms by the corresponding acid with metal atoms.
Acid salts are thermally unstable and decompose when heated to form medium salts:
\(\ \mathrm(Ca)(\mathrm(HCO) 3) 2=\mathrm(CaCOZ) \downarrow+\mathrm(CO) 2 \uparrow+\mathrm(H) 2 \mathrm(O) \).
Neutralization reactions with alkalis are characteristic of acid salts:
\(\ \mathrm(Ca)(\mathrm(HCO) 3) 2+\mathrm(Ca)(\mathrm(OH)) 2=2 \mathrm(Ca) \mathrm(CO) 3 \downarrow+2 \mathrm (H) 2 \mathrm(O)\).
b) basic salts
During dissociation, basic salts give metal cations, acid anions, and OH ions:
\(\ \mathrm(Fe)(\mathrm(OH)) \mathrm(Cl) \rightarrow \mathrm(Fe)(\mathrm(OH))++\mathrm(Cl)-+\mathrm(Fe) 2+ +\mathrm(OH)-+\mathrm(Cl) \).
Basic salts are products of incomplete replacement of the hydroxyl groups of the corresponding base with acidic residues.
Basic salts, as well as acid salts, are thermally unstable and decompose when heated:
\(\ [\mathrm(Cu)(\mathrm(OH))] 2 \mathrm(CO) 3=2 \mathrm(CuO)+\mathrm(CO) 2+\mathrm(H) 2 \mathrm(O) \).
Neutralization reactions with acids are typical for basic salts:
\(\ \mathrm(Fe)(\mathrm(OH)) \mathrm(Cl)+\mathrm(HCl) \& \text ( bull; ) \mathrm(FeCl) 2+\mathrm(H) 2 \mathrm( O)\).
c) medium salt
During dissociation, middle salts give only metal cations (ammonium ion) and acid residue anions (see above). Medium salts are products of the complete replacement of hydrogen atoms of the corresponding acid with metal atoms.
Most medium salts are thermally unstable and decompose when heated:
\(\ \mathrm(CaCO) 3=\mathrm(CaO)+\mathrm(CO) 2 \);
\(\ \mathrm(NH) 4 \mathrm(Cl)=\mathrm(NH) 3+\mathrm(HCl) \);
\(\ 2 \mathrm(Cu)(\mathrm(NO) 3) 2=2 \mathrm(CuO)+4 \mathrm(NO) 2+\mathrm(O) 2 \).
In an aqueous solution, salt salts undergo hydrolysis:
\(\ \mathrm(Al) 2 \mathrm(S) 3+6 \mathrm(H) 2 \mathrm(O) 2 \mathrm(Al)(\mathrm(OH)) 3+3 \mathrm(H) 2 \mathrm(S)\);
\(\ \mathrm(K) 2 \mathrm(S)+\mathrm(H) 2 \mathrm(O) \rightarrow \mathrm(KHS)+\mathrm(KOH) \);
\(\ \mathrm(Fe)(\mathrm(NO) 3) 3+\mathrm(H) 2 \mathrm(O) \rightarrow \mathrm(Fe)(\mathrm(OH))(\mathrm(NO) 3 ) 2+\mathrm(HNO) 3 \).
Medium salts enter into exchange reactions with acids, bases and other salts:
\(\ \mathrm(Pb)(\mathrm(NO) 3) 2+\mathrm(H) 2 \mathrm(S)=\mathrm(PbS) \downarrow+2 \mathrm(HNO) 3 \);
\(\ \mathrm(Fe) 2(\mathrm(SO) 4) 3+3 \mathrm(Ba)(\mathrm(OH)) 2=2 \mathrm(Fe)(\mathrm(OH)) 3 \downarrow +3 \mathrm(BaSO) 4\downarrow \);
\(\ \mathrm(CaBr) 2+\mathrm(K) 2 \mathrm(CO) 3=\mathrm(CaCO) 3 \downarrow+2 \mathrm(KBr) \).
Physical properties of salts
Most often, salts are crystalline substances with an ionic crystal lattice. Salts have high melting points. When n. salts are dielectrics. The solubility of salts in water is different.
Getting salts
a) acid salts
The main methods for obtaining acid salts are the incomplete neutralization of acids, the effect of excess acid oxides on bases, and the action of acids on salts:
\(\ \mathrm(NaOH)+\mathrm(H) 2 \mathrm(SO) 4=\mathrm(NaHSO) 4+\mathrm(H) 2 \mathrm(O) \);
\(\ \mathrm(Ca)(\mathrm(OH)) 2+2 \mathrm(CO) 2=\mathrm(Ca)(\mathrm(HCO) 3) 2 \);
\(\ \mathrm(CaCO) 3+\mathrm(CO) 2+\mathrm(H) 2 \mathrm(O)=\mathrm(Ca)(\mathrm(HCO) 3) 2 \).
b) basic salts
Basic salts are obtained by carefully adding a small amount of alkali to a saline solution or by the action of salts of weak acids on medium salts:
\(\ \mathrm(AICl) 3+2 \mathrm(NaOH)=\mathrm(Al)(\mathrm(OH)) 2 \mathrm(Cl)+2 \mathrm(NaCl) \);
\(\ 2 \mathrm(MgCl) 2+2 \mathrm(Na) 2 \mathrm(CO) 3+\mathrm(H) 2 \mathrm(O)=[\mathrm(Mg)(\mathrm(OH)) ] 2 \mathrm(CO) 3 \downarrow+\mathrm(CO) 2+2 \mathrm(NaCl) \).
c) medium salt
The main methods for obtaining salts of the medium are the reaction of the interaction of acids with metals, basic or amphoteric oxides and bases, as well as the reaction of the interaction of bases with acidic or amphoteric oxides and acids, the reaction of the interaction of acids and basic oxides and the exchange reaction:
\(\ \mathrm(Mg)+\mathrm(H) 2 \mathrm(SO) 4=\mathrm(MgSO) 4+\mathrm(H) 2 \);
\(\ \mathrm(Ag) 2 \mathrm(O)+2 \mathrm(HNO) \mathbf(3)=2 \mathrm(AgNO) \mathbf(3)+\mathrm(H) 2 \mathrm(O) \);
\(\ \mathrm(Cu)(\mathrm(OH)) 2+2 \mathrm(HCl)=\mathrm(CuCl) 2+2 \mathrm(H) 20 \);
\(\ 2 \mathrm(KOH)+\mathrm(SO) 2=\mathrm(K) 2 \mathrm(SO) 3+\mathrm(H) 20 \);
\(\ \mathrm(CaO)+\mathrm(SO) 3=\mathrm(CaSO) 4 \);
\(\ \mathrm(BaCl) 2+\mathrm(MgSO) 4=\mathrm(MgCl) 2+\mathrm(BaSO) 4\downarrow \).
Problem Solving Examples
Determine the mass of ammonium chloride, which is formed by the interaction of 5.9 g of ammonia with 5.6 l (N.O.) of hydrogen chloride.
Let us write the equation for the formation of ammonium chloride from ammonia and hydrogen chloride: \(\ \mathrm(NH) 3+\mathrm(HCl)=\mathrm(NH) 4 \mathrm(Cl) \).
Determine which of the substances is in excess and which is in short supply:
\(\ \mathrm(v)(\mathrm(NH) 3)=\mathrm(m)(\mathrm(NH) 3) / \mathrm(M)(\mathrm(NH) 3)=5,6 / 17 \u003d 0.33 \) mol;
\(\ \mathrm(v)(\mathrm(HCl))=\mathrm(V)(\mathrm(HCl)) / \mathrm(Vm)=5.6 / 22.4=0.25 \) mol.
The calculation is made on a substance that is in short supply - on hydrochloric acid. Calculate the mass of ammonium chloride:
\(\ \mathrm(v)(\mathrm(HCl))=\mathrm(v)(\mathrm(NH) 4 \mathrm(Cl))=0.25 \) mol;
\(\ (\mathrm(NH) 4 \mathrm(Cl))=0.25 \times 53.5=13.375 \mathrm(r)\).
The mass of ammonium chloride is 13.375 g.
Determine the amount of substance, volume (n.o.s.) and mass of ammonia required to obtain 250 g of ammonium sulfate used as a fertilizer.
Let's write the equation for the reaction of obtaining ammonium sulfate from ammonia and sulfuric acid:
\(\ 2 \mathrm(NH) 3+\mathrm(H) 2 \mathrm(SO) 4=(\mathrm(NH) 4) \quad 2 \mathrm(SO) 4 \).
Molar mass of ammonium sulfate calculated using the D.I. Mendeleev - 132 g / mol. Then the amount of ammonium sulfate:
\(\ \mathrm(v)((\mathrm(NH) 4) \quad 2 \mathrm(SO) 4)=\mathrm(m)((NH 4) 2 S 04) / M((NH 4) 2 S 04) \)
\(\ \mathrm(v)((\mathrm(NH) 4) \quad 2 \mathrm(S) 04)=250 / 132=1.89 \) mol
According to the reaction equation \(\ \mathrm(v)((\mathrm(NH) 4) \quad 2 \mathrm(S) 04) : \mathrm(v)(\mathrm(NH) 3)=1: 2 \) , so the amount of ammonia is:
\(\ \mathrm(v)(\mathrm(NH) 3)=2 \times \mathrm(v)((\mathrm(NH) 4) 2 \mathrm(SO) 4)=2 \times 1,89= 3.79 \) mol.
Determine the volume of ammonia:
\(\ \mathrm(V)(\mathrm(NH) 3)=\mathrm(v)(\mathrm(NH) 3) \times \mathrm(V)_(\mathrm(m)) \);
\(\ V(N H 3)=3.79 \times 22.4=84.8 l\).
The molar mass of ammonia, calculated using the table of chemical elements of D.I. Mendeleev - 17 g/mol. Then, find the mass of ammonia:
\(\ \mathrm(m)(\mathrm(NH) 3)=\mathrm(v)(\mathrm(NH) 3) \times \mathrm(M)(\mathrm(NH) 3) \);
\(\ \mathrm(m)(\mathrm(NH) 3)=3.79 \times 17=64.43 \mathrm(r)\).
The amount of ammonia substance is 3.79 mol, the volume of ammonia is 84.8 l, the mass of ammonia is 64.43 g.
SALT, a class of chemical compounds. A generally accepted definition of the concept of “Salts”, as well as the terms “acids and bases”, the products of the interaction of which salts are, currently does not exist. Salts can be considered as products of substitution of acid hydrogen protons for metal ions, NH 4 + , CH 3 NH 3 + and other cations or OH groups of the base for acid anions (eg, 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 , СuСlOH. 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 include two types of acid residues ( e.g. AgClBr). Complex salts contain complex ions such as K 4 .
Physical properties
Typical salts are crystalline substances with an ionic structure, such as CsF. There are also covalent salts, such as AlCl 3 . In fact, the nature of the chemical bond v of many salts is mixed.
By solubility in water, soluble, slightly soluble and practically insoluble salts are distinguished. Soluble include almost all salts of sodium, potassium and ammonium, many nitrates, acetates and chlorides, with the exception of salts of polyvalent metals that hydrolyze in water, many acidic salts.
Solubility of salts in water at room temperature
| 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. Getting 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:
Thermal decomposition to medium salt
Ca (HCO 3) 2 \u003d CaCO 3 + CO 2 + H 2 O
Interaction with alkali. Obtaining 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 \u003d 2CuO + CO 2 + H 2 O
Interaction with acid: 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:
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 double chain anion (B 2 O 3 (OH) 2) n 2n-triborates - if they contain 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 nH 2 decaborates O. Alkaline earth and other metals in the oxidation state + 2 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°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), 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. The 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 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 EHn compounds at a constant n, the bond covalence 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 series MFn and MC1n 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 with the transition from fluorides to iodides and with an increase in 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 in cases where 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 is the oxidation state of the metal M), and in the form of crystalline hydrates M (NO 3) n xH 2 O (x \u003d 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 oxidizers melt incongruently in water of crystallization 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 НNO 2 . First of all, nitrites of alkali metals and ammonium are used, less - alkaline earth and Zd-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 polyhalite), 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
1. Filtration, separation of inhomogeneous 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 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 speed of the process 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 (see 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.
2. Fractional crystallization
There are the following types of fractional crystallization: mass, on cooled surfaces, directional, zone melting.
Mass crystallization. The method consists in the simultaneous obtaining of a large number of crystals in 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, equipped with external cooling jackets or internal coils and often mixing devices; tubular, scraper, disk, screw, etc. Due to the lack of a calculation method, the parameter ae, 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 the 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 (eg 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) crystallize 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 the 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 liquid. So, when separating crystals on a drum vacuum filter, the concentration of the mother liquor in them is 10-30%, on a filtering centrifuge - 3-10%.
The main advantages of the process: high productivity, no contact between the mixture to be separated and the refrigerant, simplicity of instrumentation; disadvantages: relatively low heat transfer coefficient, inlay of cooling surfaces, large capture of the mother liquid by crystals, the need to install additional equipment for the separation of solid and liquid phases, insufficiently high yield of a crystalline product. Application examples: preparation of K and Na chlorides from sylvinite, separation of xylene isomers.
3. Evaporation, carried out to concentrate the solution, isolate the solute or obtain a pure solvent. Evaporation is mainly applied to aqueous solutions. The coolant is most often 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 of evaporation - the temperature difference between the heating steam and the boiling solution, is called useful. It is always less than the temperature difference between 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 rises due to the higher pressure in the solution than in the vapor space. Causes of pressure increase: hydrostatic pressure of the solution; hydraulic resistance during the movement of a boiling (vapor-liquid) mixture; an increase in the speed of movement of this mixture due to the fact that it occupies a much larger volume than the original solution (respectively, hydrostatic, hydraulic and inertial depression).
For evaporation, devices operating under pressure or vacuum are used. Their main elements are: heating chamber; a separator for separating the vapor-liquid mixture in the selection of a concentrated solution; circulation pipe, through which the solution is returned from the separator to the chamber (with multiple 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 abundant 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 annulus 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 evaporators with natural circulation, in the heating tubes of which the solution not only heats up, but also boils. For evaporation of solutions of poorly soluble substances that precipitate during concentration [for example, CaCO 3 , CaSO 4 , Mg (OH) 2 , Na aluminosilicate], as well as for the desalination of sea water, apparatuses are used, above the heating chamber of which an additional, lifting circulation a pipe that provides a high rate of natural circulation. For the evaporation of highly foaming and temperature-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 the solution together with secondary steam along tubes of length 6-8 m (heating surface up to 2200 m2). Advantages of these devices: no hydrostatic effect, low hydraulic resistance, high heat transfer coefficient, high performance 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 to isolate fractions from suspensions and emulsions, as well as to determine the molecular weights of polymers, 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, which creates a field of centrifugal forces up to 20,000 g in industrial centrifuges and up to 350,000 g in laboratory ones (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 sediment, centrifuges: 1) periodic action, 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 slots, clarified liquid (centrate) is discharged from the top of its part; 2) continuous action, in which the suspension is fed along the axis of the 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 according to the filtration principle is most often used to separate suspensions and sludges with a relatively low liquid phase content and is carried out in cyclic machines. The suspension is fed into the continuously rotating rotor in portions; after filling the part of the rotor with sediment, the supply of the suspension is stopped, the liquid phase is squeezed out, and the sediment is cut off with a knife and removed. Centrifuges are also used with pulsating sludge discharge using a pusher (vibrating piston, with a pulsating piston), as well as with hydraulic unloading, when the thickened solid phase is removed from the rotor equipped with a package of conical plates through nozzles.
Bibliography
Ch. editor I.L. Knunyants. Big Encyclopedic Dictionary 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 et al. Fascinating world of chemical transformations. Saint Petersburg 1998
E. V. Barkovsky. Introduction to the chemistry of biogenic elements and chemical analysis. Minsk 1997
Interaction of medium salts with metals
The reaction of a salt with a metal proceeds if the original free metal is more active than the one that is part of the original salt. You can find out which metal is more active by using the electrochemical series of metal voltages.
So, for example, iron interacts with copper sulfate in an aqueous solution, since it is more active than copper (to the left in the activity series):
At the same time, iron does not react with a solution of zinc chloride, since it is less active than zinc:
It should be noted that such active metals as alkali and alkaline earth, when added to aqueous solutions of salts, will primarily react not with salt, but with water included in the solutions.
Interaction of medium salts with metal hydroxides
Let us make a reservation that, in this case, metal hydroxides are understood to mean compounds of the form Me (OH) x.
In order for the middle salt to react with the metal hydroxide, simultaneously (!) two requirements are met:
- a precipitate or gas must be detected in the intended products;
- the original salt and the original metal hydroxide must be soluble.
Consider a couple of cases in order to learn this rule.
Let us determine which of the reactions below proceed, and write the equations of the proceeding reactions:
- 1) PbS + KOH
- 2) FeCl 3 + NaOH
Consider the first interaction of lead sulfide and potassium hydroxide. Let's write down the supposed ion exchange reaction and mark it on the left and right with “blinds”, denoting it in such a way that it is not yet known whether the reaction actually proceeds:
In the proposed products, we see lead (II) hydroxide, which, according to the solubility table, is insoluble and must precipitate. However, the conclusion that the reaction proceeds cannot yet be made, since we have not verified the satisfaction of one more mandatory requirement - the solubility of the initial salt and hydroxide. Lead sulfide is an insoluble salt, which means that the reaction does not proceed, since one of the mandatory requirements for the reaction between the salt and the metal hydroxide is not met. Those.:
Consider the second proposed interaction between iron(III) chloride and potassium hydroxide. Let's write down the expected ion exchange reaction and mark it on the left and right with "curtains", as in the first case:
In the proposed products, we see iron(III) hydroxide, which is insoluble and must precipitate. However, it is still impossible to draw a conclusion about the course of the reaction. To do this, we must also verify the solubility of the initial salt and hydroxide. Both starting materials are soluble, so we can conclude that the reaction is proceeding. Let's write down its equation:
Reactions of medium salts with acids
The middle salt reacts with an acid if a precipitate or a weak acid is formed.
It is almost always possible to recognize a precipitate among the intended products by looking at the solubility table. So, for example, sulfuric acid reacts with barium nitrate, since insoluble barium sulfate precipitates:
It is impossible to recognize a weak acid from a solubility table, since many weak acids are soluble in water. Therefore, the list of weak acids should be learned. Weak acids include H 2 S, H 2 CO 3 , H 2 SO 3 , HF, HNO 2 , H 2 SiO 3 and all organic acids.
So, for example, hydrochloric acid reacts with sodium acetate, since a weak organic acid (acetic acid) is formed:
It should be noted that hydrogen sulfide H 2 S is not only a weak acid, but also poorly soluble in water, and therefore it is released from it in the form of a gas (with the smell of rotten eggs):
In addition, it should be remembered that weak acids - carbonic and sulfurous - are unstable and decompose almost immediately after formation into the corresponding acid oxide and water:
It was said above that the reaction of a salt with an acid occurs if a precipitate or a weak acid is formed. Those. if there is no precipitate and a strong acid is present in the intended products, then the reaction will not proceed. However, there is a case that does not formally fall under this rule, when concentrated sulfuric acid displaces hydrogen chloride when it acts on solid chlorides:
However, if we take not concentrated sulfuric acid and solid sodium chloride, but solutions of these substances, then the reaction will really not go:
Reactions of medium salts with other medium salts
The reaction between medium salts proceeds if simultaneously (!) two requirements are met:
- initial salts are soluble;
- there is sediment or gas in the intended products.
For example, barium sulfate does not react with potassium carbonate, because despite the fact that there is a precipitate (barium carbonate) in the intended products, the requirement for the solubility of the original salts is not met.
At the same time, barium chloride reacts with potassium carbonate in solution, since both initial salts are soluble, and there is a precipitate in the products:
Gas during the interaction of salts is formed in the only case - if you mix a solution of any nitrite with a solution of any ammonium salt when heated:
The reason for the formation of gas (nitrogen) is that the solution simultaneously contains NH 4 + cations and NO 2 - anions, which form thermally unstable ammonium nitrite, which decomposes in accordance with the equation:
Thermal decomposition reactions of salts
Decomposition of carbonates
All insoluble carbonates, as well as lithium and ammonium carbonates, are thermally unstable and decompose when heated. Metal carbonates decompose to metal oxide and carbon dioxide:
and ammonium carbonate gives three products - ammonia, carbon dioxide and water:
Decomposition of nitrates
Absolutely all nitrates decompose when heated, and the type of decomposition depends on the position of the metal in the activity series. The decomposition scheme of metal nitrates is shown in the following illustration:
So, for example, in accordance with this scheme, the decomposition equations for sodium nitrate, aluminum nitrate and mercury nitrate are written as follows:
It should also be noted the specifics of the decomposition of ammonium nitrate:
Decomposition of ammonium salts
Thermal decomposition of ammonium salts is most often accompanied by the formation of ammonia:
If the acid residue has oxidizing properties, instead of ammonia, some product of its oxidation is formed, for example, molecular nitrogen N 2 or nitrogen oxide (I):
Chemical properties of acid salts
The ratio of acid salts to alkalis and acids
Acid salts react with alkalis. Moreover, if the alkali contains the same metal as the acid salt, then medium salts are formed:
Also, if two or more mobile hydrogen atoms remain in the acid residue of the acid salt, as, for example, in sodium dihydrogen phosphate, then the formation of both an average is possible:
and another acid salt with a smaller number of hydrogen atoms in the acid residue:
It is important to note that acid salts react with any alkalis, including those formed by another metal. For instance:
Acid salts formed by weak acids react with strong acids in the same way as the corresponding medium salts:
Thermal decomposition of acid salts
All acidic salts decompose when heated. As part of the USE program in chemistry, from the reactions of decomposition of acid salts, one should learn how hydrocarbonates decompose. Metal bicarbonates decompose already at a temperature of more than 60 ° C. In this case, metal carbonate, carbon dioxide and water are formed:
The last two reactions are the main cause of scale formation on the surface of water heating elements in electric kettles, washing machines, etc.
Ammonium bicarbonate decomposes without a solid residue with the formation of two gases and water vapor:
Chemical properties of basic salts
Basic salts always react with all strong acids. In this case, medium salts can be formed if an acid with the same acid residue is used as in the basic salt, or mixed salts if the acid residue in the basic salt differs from the acid residue of the acid reacting with it:
Also, basic salts are characterized by decomposition reactions when heated, for example:
Chemical properties of complex salts (on the example of aluminum and zinc compounds)
As part of the USE program in chemistry, one should learn the chemical properties of such complex compounds of aluminum and zinc as tetrahydroxoaluminates and tetrahydroxozincates.
Tetrahydroxoaluminates and tetrahydroxozincates are salts whose anions have the formulas - and 2-, respectively. Consider the chemical properties of such compounds using sodium salts as an example:
These compounds, like other soluble complex ones, dissociate well, while almost all complex ions (in square brackets) remain intact and do not dissociate further:
The action of an excess of a strong acid on these compounds leads to the formation of two salts:
Under the action of a lack of strong acids on them, only the active metal passes into a new salt. Aluminum and zinc in the composition of hydroxides precipitate:
Precipitation of aluminum and zinc hydroxides with strong acids is not a good choice, since it is difficult to add the strictly necessary amount of strong acid for this without dissolving part of the precipitate. For this reason, carbon dioxide is used for this, which has very weak acidic properties and is therefore incapable of dissolving the hydroxide precipitate:
In the case of tetrahydroxoaluminate, hydroxide precipitation can also be carried out using sulfur dioxide and hydrogen sulfide:
In the case of tetrahydroxozincate, precipitation with hydrogen sulfide is impossible, since its sulfide precipitates instead of zinc hydroxide:
Upon evaporation of solutions of tetrahydroxozincate and tetrahydroxoaluminate followed by calcination, these compounds are converted into zincate and aluminate, respectively.
Video lesson 1: Classification of inorganic salts and their nomenclature
Video lesson 2: Methods for obtaining inorganic salts. Chemical properties of salts
Lecture: Characteristic chemical properties of salts: medium, acidic, basic; complex (on the example of aluminum and zinc compounds)
Characteristics of salts
salt- these are chemical compounds consisting of metal cations (or ammonium) and acidic residues.
Salts should also be considered as a product of the interaction of an acid and a base. As a result of this interaction, the following can be formed:
basic salts.
normal (medium),
normal salts are formed when the amount of acid and base is sufficient for complete interaction. For example:
H 3 RO 4 + 3KOH → K 3 RO 4 + 3H 2 O.
The names of normal salts consist of two parts. First, the anion (acid residue) is called, then the cation. For example: sodium chloride - NaCl, iron (III) sulfate - Fe 2 (SO 4) 3, potassium carbonate - K 2 CO 3, potassium phosphate - K 3 PO 4, etc.
Acid salts are formed with an excess of acid and an insufficient amount of alkali, because in this case there are not enough metal cations to replace all the hydrogen cations present in the acid molecule. For example:
H 3 RO 4 + 2KOH \u003d K 2 HRO 4 + 2H 2 O;
H 3 RO 4 + KOH \u003d KN 2 RO 4 + H 2 O.
As part of the acid residues of this type of salt, you will always see hydrogen. Acid salts are always possible for polybasic acids, but not for monobasic acids.
The names of acid salts are prefixed hydro- to the anion. For example: iron (III) hydrogen sulfate - Fe (HSO 4) 3, potassium bicarbonate - KHCO 3, potassium hydrogen phosphate - K 2 HPO 4, etc.
Basic salts are formed with an excess of base and an insufficient amount of acid, because in this case the anions of acid residues are not enough to completely replace the hydroxo groups present in the base. For example:
Cr(OH) 3 + HNO 3 → Cr(OH) 2 NO 3 + H 2 O;
Cr(OH) 3 + 2HNO 3 → CrOH(NO 3) 2 + 2H 2 O.
Thus, the basic salts in the composition of cations contain hydroxo groups. Basic salts are possible for polyacid bases, but not for monoacid ones. Some basic salts are able to decompose on their own, while releasing water, forming oxosalts, which have the properties of basic salts. For example:
Sb(OH) 2 Cl → SbOCl + H 2 O;
Bi(OH) 2 NO 3 → BiONO 3 + H 2 O.
The name of the basic salts is built as follows: the prefix is added to the anion hydroxo-. For example: iron (III) hydroxosulfate - FeOHSO 4, aluminum hydroxosulfate - AlOHSO 4, iron (III) dihydroxochloride - Fe (OH) 2 Cl, etc.
Many salts, being in a solid state of aggregation, are crystalline hydrates: CuSO4.5H2O; Na2CO3.10H2O etc.
Chemical properties of salts
Salts are fairly solid crystalline substances that have an ionic bond between cations and anions. The properties of salts are due to their interaction with metals, acids, bases and salts.
Typical reactions of normal salts
They react well with metals. At the same time, more active metals displace less active ones from solutions of their salts. For example:
Zn + CuSO 4 → ZnSO 4 + Cu;
Cu + Ag 2 SO 4 → CuSO 4 + 2Ag.
With acids, alkalis and other salts, the reactions go to completion, provided that a precipitate, gas, or poorly dissociated compounds are formed. For example, in the reactions of salts with acids, substances such as hydrogen sulfide H 2 S are formed - gas; barium sulfate BaSO 4 - precipitate; acetic acid CH 3 COOH is a weak electrolyte, a poorly dissociated compound. Here are the equations for these reactions:
K 2 S + H 2 SO 4 → K 2 SO 4 + H 2 S;
BaCl 2 + H 2 SO 4 → BaSO 4 + 2HCl;
CH 3 COONa + HCl → NaCl + CH 3 COOH.
In the reactions of salts with alkalis, substances such as nickel (II) hydroxide Ni (OH) 2 are formed - a precipitate; ammonia NH 3 - gas; water H 2 O is a weak electrolyte, a low-dissociation compound:
NiCl 2 + 2KOH → Ni(OH) 2 + 2KCl;
NH 4 Cl + NaOH → NH 3 + H 2 O + NaCl.
Salts react with each other if a precipitate forms:
Ca(NO 3) 2 + Na 2 CO 3 → 2NaNO 3 + CaCO 3.
Or in the case of the formation of a more stable compound:
Ag 2 CrO 4 + Na 2 S → Ag 2 S + Na 2 CrO 4 .
In this reaction, brick-red silver chromate produces black silver sulfide, due to the fact that it is a more insoluble precipitate than chromate.
Many normal salts decompose when heated to form two oxides - acidic and basic:
CaCO 3 → CaO + CO 2.
Nitrates decompose in a different way than other normal salts. When heated, alkali and alkaline earth metal nitrates release oxygen and turn into nitrites:
2NaNO 3 → 2NaNO 2 + O 2.
Nitrates of almost all other metals decompose to oxides:
2Zn(NO 3) 2 → 2ZnO + 4NO 2 + O 2 .
Nitrates of some heavy metals (silver, mercury, etc.) decompose when heated to metals:
2AgNO 3 → 2Ag + 2NO 2 + O 2.
A special position is occupied by ammonium nitrate, which, up to the melting point (170 ° C), partially decomposes according to the equation:
NH 4 NO 3 → NH 3 + HNO 3.
At temperatures of 170 - 230 ° C, according to the equation:
NH 4 NO 3 → N 2 O + 2H 2 O.
At temperatures above 230 ° C - with an explosion, according to the equation:
2NH 4 NO 3 → 2N 2 + O 2 + 4H 2 O.
Ammonium chloride NH 4 Cl decomposes to form ammonia and hydrogen chloride:
NH 4 Cl → NH 3 + HCl.
Typical reactions of acid salts
They enter into all those reactions that acids enter into. They react with alkalis as follows, if the acid salt and alkali contain the same metal, then a normal salt is formed as a result. For example:
NaH CO3+ Na Oh→ Na 2 CO3+ H2O.
NaHCO 3 + Li Oh → Li NaCO 3+ H2O.
Typical reactions major salts
These salts undergo the same reactions as the bases. They react with acids as follows, if the basic salt and acid contain the same acid residue, then a normal salt is formed as a result. For example:
Cu( Oh)Cl+ H Cl → Cu Cl 2 + H2O.
Cu( Oh)Cl + HBr → Cu Br Cl+ H2O.
Complex salts
complex connection- a compound whose crystal lattice sites contain complex ions.
Consider the complex compounds of aluminum - tetrahydroxoaluminates and zinc - tetrahydroxozincates. Complex ions are indicated in square brackets of the formulas of these substances.
Chemical properties of sodium tetrahydroxoaluminate Na and sodium tetrahydroxozincate Na 2:
1. Like all complex compounds, the above substances dissociate:
- Na → Na + + - ;
- Na 2 → 2Na + + - .
Keep in mind that further dissociation of complex ions is not possible.
2. In reactions with an excess of strong acids, they form two salts. Consider the reaction of sodium tetrahydroxoaluminate with a dilute solution of hydrogen chloride:
- Na + 4HCl→ Al Cl3 + Na Cl + H2O.
We see the formation of two salts: aluminum chloride, sodium chloride and water. A similar reaction will occur in the case of sodium tetrahydroxozincate.
3. If a strong acid is not enough, let's say instead of 4 HCl We took 2 HCl then the salt forms the most active metal, in this case sodium is more active, which means that sodium chloride is formed, and the resulting aluminum and zinc hydroxides will precipitate. Let us consider this case in the reaction equation with sodium tetrahydroxozincate:
Na 2 + 2HCl→ 2Na Cl+ Zn (OH) 2 ↓ +2H2O.
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