The solubility table of chemical elements is a table with the water solubilities of the most famous inorganic acids, bases and salts.
Definition 1
The solubility table in chemistry shows solubility at 20 °C, solubility increases with increasing temperature.
A substance is soluble in water if its solubility is more than 1 g per 100 g of water and insoluble if less than 0.1 g/100 g. For example, by finding lithium in the solubility table in chemistry, you can be sure that almost all of its salts form solutions.
In Fig. 1 and fig. 2 shows a photo of the complete solubility table in chemistry with the names of acid residues.
Figure 1. Photo solubility table in chemistry 2018-2019
Figure 2. Chemistry table with acids and acid residues
To make up the name of a salt, you need to use the periodic table and solubility. The name of the acid residue is added to the name of the metal from the periodic table, for example:
$\mathrm(Zn_3(PO_4)_2)$ - zinc phosphate; $\mathrm(FeSO_4)$ - iron (II) sulfate.
In brackets with the text name, you must indicate the valence of the metal, if there are several of them. In the case of iron, there is also a salt $\mathrm(Fe_2(SO_4)_3)$ - iron (III) sulfate.
What can you learn using the solubility table in chemistry?
The solubility table for substances in chemistry with precipitates is used to determine the possibility of any reaction occurring, since the formation of a precipitate or gas is necessary for the irreversible reaction to occur.
The solubility table for salts, acids and bases is the foundation without which it is impossible to fully master chemical knowledge. The solubility of bases and salts helps in learning not only for schoolchildren, but also for professional people. The creation of many life products cannot do without this knowledge.
Table of solubility of acids, salts and bases in water
The table of solubility of salts and bases in water is a guide that helps in mastering the basics of chemistry. The following notes will help you understand the table below.
- P - indicates a soluble substance;
- H is an insoluble substance;
- M - the substance is slightly soluble in the aquatic environment;
- RK - a substance that can dissolve only when exposed to strong organic acids;
- The dash will say that such a creature does not exist in nature;
- NK – does not dissolve in either acids or water;
- ? – a question mark indicates that today there is no accurate information about the dissolution of the substance.
Often, the table is used by chemists and schoolchildren, students to conduct laboratory research, during which it is necessary to establish the conditions for the occurrence of certain reactions. Using the table, it is possible to determine how a substance will behave in a salt or acidic environment, and whether a precipitate may appear. A precipitate during research and experiments indicates the irreversibility of the reaction. This is a significant point that can affect the course of all laboratory work.
Water is one of the main chemical compounds on our planet. One of its most interesting properties is the ability to form aqueous solutions. And in many areas of science and technology, the solubility of salt in water plays an important role.
Solubility is understood as the ability of various substances to form homogeneous (homogeneous) mixtures with liquids - solvents. It is the volume of material that is used to dissolve and form a saturated solution that determines its solubility, comparable to the mass fraction of this substance or its amount in a concentrated solution.
According to their ability to dissolve, salts are classified as follows:
- Soluble substances include substances that can be dissolved in 100 g of water more than 10 g;
- Slightly soluble include those whose amount in the solvent does not exceed 1 g;
- the concentration of insolubles in 100 g of water is less than 0.01.
When the polarity of the substance used for dissolution is similar to the polarity of the solvent, it is soluble. With different polarities, it is most likely not possible to dilute the substance.
How does dissolution occur?
If we talk about whether salt dissolves in water, then for most salts this is a fair statement. There is a special table according to which you can accurately determine the solubility value. Since water is a universal solvent, it mixes well with other liquids, gases, acids and salts.
One of the most obvious examples of the dissolution of a solid in water can be observed almost every day in the kitchen, while preparing dishes using table salt. So why does salt dissolve in water?
Many people remember from their school chemistry course that the molecules of water and salt are polar. This means that their electrical poles are opposite, resulting in a high dielectric constant. Water molecules surround ions of another substance, for example, in the case we are considering, NaCl. This produces a liquid that is homogeneous in consistency.
Effect of temperature
There are some factors that influence the solubility of salts. First of all, this is the temperature of the solvent. The higher it is, the greater the diffusion coefficient of particles in the liquid, and mass transfer occurs faster.
Although, for example, the solubility of table salt (NaCl) in water practically does not depend on temperature, since its solubility coefficient is 35.8 at 20° C and 38.0 at 78° C. But copper sulfate (CaSO4) with increasing temperature water dissolves less well.
Other factors that affect solubility include:
- Size of dissolved particles – with a larger area of phase separation, dissolution occurs faster.
- A mixing process that, when performed intensively, promotes more efficient mass transfer.
- The presence of impurities: some accelerate the dissolution process, while others, by complicating diffusion, reduce the speed of the process.
Video about the mechanism of salt dissolution
| Cations | Anions | |||||||||
| F- | Cl- | Br- | I - | S 2- | 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 |
| Mg 2+ | RK | R | R | R | M | R | N | RK | R | RK |
| Ca2+ | NK | R | R | R | M | R | N | RK | M | RK |
| Sr 2+ | NK | R | R | R | R | R | N | RK | RK | RK |
| Ba 2+ | RK | R | R | R | R | R | N | RK | NK | RK |
| Sn 2+ | R | R | R | M | RK | R | N | N | R | N |
| Pb 2+ | N | M | M | M | RK | R | N | N | N | N |
| Al 3+ | M | R | R | R | G | R | G | NK | R | RK |
| Cr 3+ | R | R | R | R | G | R | G | N | R | RK |
| Mn 2+ | R | R | R | R | N | R | N | N | R | N |
| Fe 2+ | M | R | R | R | N | R | N | N | R | N |
| Fe 3+ | R | R | R | - | - | R | G | N | R | RK |
| Co2+ | M | R | R | R | N | R | N | N | R | N |
| Ni 2+ | M | R | R | R | RK | R | N | N | R | N |
| Cu 2+ | M | R | R | - | N | R | G | N | R | N |
| Zn 2+ | M | R | R | R | RK | R | N | N | R | N |
| Cd 2+ | R | R | R | R | RK | R | N | N | R | N |
| Hg 2+ | R | R | M | NK | NK | R | N | N | R | N |
| Hg 2 2+ | R | NK | NK | NK | RK | R | N | N | M | N |
| Ag+ | R | NK | NK | NK | NK | R | N | N | M | N |
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 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 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 \u003d 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 + H 2 SO 4 \u003d KHSO 4 + H 2 O
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 \u003d 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. Obtaining complex salts:
1. Reactions of salts with ligands: AgCl + 2NH 3 = Cl
FeCl 3 + 6KCN] = K 3 + 3KCl
5. Obtaining double salts:
1. Joint crystallization of two salts:
Cr 2 (SO 4) 3 + K 2 SO 4 + 24H 2 O = 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:
1. Thermal decomposition with the formation of medium salt
Ca(HCO 3) 2 = CaCO 3 + CO 2 + H 2 O
2. Interaction with alkali. Getting medium salt.
Ba(HCO 3) 2 + Ba(OH) 2 = 2BaCO 3 + 2H 2 O
3. Chemical properties of basic salts:
1. Thermal decomposition. 2 CO 3 = 2CuO + CO 2 + H 2 O
2. 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:
1. 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.
Main types of salts
1. Borates (oxoborates), salts of boric acids: metaboric HBO 2, orthoboric H3 BO 3 and polyboronic acids not isolated in the 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 (BO2) n n - diborates - if they contain a chain double anion (B 2 O 3 (OH) 2) n 2n- triborates - if they contain a ring anion (B 3 O 6) 3-.
The structures of borates include boron-oxygen groups - “blocks” containing from 1 to 6, and sometimes 9 boron atoms, for example:
The coordination number of boron atoms is 3 (boron-oxygen triangular groups) or 4 (tetrahedral groups). Boron-oxygen groups are the basis of not only island, but also more complex structures - chain, layered and 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 n H 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; 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.
Alkali metal, ammonium and T1(I) borates 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 SO 2 ;. 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 borates. 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 VO 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 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. To simple or binary EC halides n (n- most often an integer from 1 for monohalides to 7 for IF 7 and ReF 7, but can also be fractional, for example 7/6 for Bi 6 Cl 7) include, in particular, salts of hydrohalic acids and interhalogen compounds (for example, halofluorides). There are also mixed halides, polyhalides, hydrohalides, oxohalides, oxyhalides, hydroxohalides, thiohalides and complex halides. The oxidation 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 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 rows of connections EH n 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 ranks of MF n and MC1 n 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 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:
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 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 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 sea and ocean water, salt and underground brines. Some halides, for example NaCl, KC1, CaCl 2, are part of living organisms.
3. Carbonates (from Latin carbo, gender 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 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 2PbCO 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- oxidation state of the metal M), and in the form of crystalline hydrates M(NO 3) n x H 2 O ( X= 1-9). From aqueous solutions at a temperature close to room temperature, only alkali metal nitrates crystallize anhydrous, the rest - in the form of crystalline hydrates. The physicochemical properties of anhydrous and hydrated nitrate of the same metal can be very different.
Anhydrous crystalline compounds of d-element nitrates are colored. Conventionally, nitrates can be divided into compounds with a predominantly covalent type of bond (salts of Be, Cr, Zn, Fe and other transition metals) and with a predominantly ionic type of bond (salts of alkali and alkaline earth metals). Ionic nitrates are characterized by higher thermal stability, the predominance of crystal structures of higher symmetry (cubic) and the absence of splitting of the nitrate ion bands in the IR spectra. Covalent nitrates have 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, less - alkaline earth and nitrites. d-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 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 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 shenites 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. 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 = 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 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). 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 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 of 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 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 / 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 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
Filtration, separation of heterogeneous systems of 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 in 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. (cm. Fig.) 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 wash 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.
A salt can be defined as a compound that is formed by the reaction between an acid and a base, but is not water. This section will consider those properties of salts that are associated with ionic equilibria.
reactions of salts in water
Somewhat later it will be shown that solubility is a relative concept. However, for the purposes of the discussion ahead, we can roughly divide all salts into those that are soluble and those that are insoluble in water.
Some salts, when dissolved in water, form neutral solutions. Other salts form acidic or alkaline solutions. This is due to the occurrence of a reversible reaction between salt ions and water, as a result of which conjugate acids or bases are formed. Whether the salt solution turns out to be neutral, acidic or alkaline depends on the type of salt. In this sense, there are four types of salts.
Salts formed by strong acids and weak bases. Salts of this type, when dissolved in water, form an acidic solution. Let's take ammonium chloride NH4Cl as an example. When this salt is dissolved in water, the ammonium ion acts as
The excess amount of H3O+ ions formed in this process causes the acidic properties of the solution.
Salts formed by a weak acid and a strong base. Salts of this type, when dissolved in water, form an alkaline solution. As an example, let's take sodium acetate CH3COONa1. The acetate ion acts as a base, accepting a proton from water, which in this case acts as an acid:
The excess amount of OH- ions formed in this process determines the alkaline properties of the solution.
Salts formed by strong acids and strong bases. When salts of this type are dissolved in water, a neutral solution is formed. Let's take sodium chloride NaCl as an example. When dissolved in water, this salt is completely ionized, and, therefore, the concentration of Na+ ions turns out to be equal to the concentration of Cl- ions. Since neither one nor the other ion enters into acid-base reactions with water, an excess amount of H3O+ or OH ions does not form in the solution. Therefore, the solution turns out to be neutral.
Salts formed by weak acids and weak bases. An example of this type of salt is ammonium acetate. When dissolved in water, ammonium ion reacts with water as an acid, and acetate ion reacts with water as a base. Both of these reactions are described above. An aqueous solution of a salt formed by a weak acid and a weak base can be weakly acidic, weakly alkaline, or neutral, depending on the relative concentrations of the H3O+ and OH- ions formed as a result of the reactions of the salt's cations and anions with water. This depends on the relationship between the values of the dissociation constants of the cation and anion.
