Simple substances. Molecules are made up of atoms of the same type (atoms of the same element). In chemical reactions they cannot decompose to form other substances.
Complex substances (or chemical compounds). Molecules are made up of atoms different types(atoms of various chemical elements). In chemical reactions they decompose to form several other substances.
There is no sharp boundary between metals and non-metals, because There are simple substances that exhibit dual properties.
Allotropy
Allotropy- the ability of some chemical elements form several simple substances, differing in structure and properties.
C - diamond, graphite, carbine.
O - oxygen, ozone.
S - rhombic, monoclinic, plastic.
P - white, red, black.
The phenomenon of allotropy is caused by two reasons:
1) different number atoms in a molecule, such as oxygen O 2 and ozone O 3
2) the formation of various crystalline forms, for example diamond and graphite.
BASES
Reasons- complex substances in which metal atoms are connected to one or more hydroxyl groups (from the point of view of theory electrolytic dissociation, bases are complex substances, the dissociation of which in an aqueous solution produces metal cations (or NH 4 +) and hydroxide anions OH -).
Classification. Soluble in water (alkalies) and insoluble. Amphoteric bases also exhibit the properties of weak acids.
Receipt
1. Reactions of active metals (alkali and alkaline earth metals) with water:
2Na + 2H 2 O ® 2NaOH + H 2 -
Ca + 2H 2 O ® Ca(OH) 2 + H 2 -
2. Interaction of active metal oxides with water:
BaO + H 2 O ® Ba(OH) 2
3. Electrolysis of aqueous salt solutions
2NaCl + 2H 2 O ® 2NaOH + H 2 - + Cl 2 -
Chemical properties
| Alkalis | Insoluble bases |
| 1. Action on indicators. | |
| litmus - blue methyl orange - yellow phenolphthalein - raspberry |
-- |
| 2. Interaction with acid oxides. | |
| 2KOH + CO 2 ® K 2 CO 3 + H 2 O KOH + CO 2 ® KHCO 3 |
-- |
| 3. Interaction with acids (neutralization reaction) | |
| NaOH + HNO 3 ® NaNO 3 + H 2 O | Cu(OH) 2 + 2HCl ® CuCl 2 + 2H 2 O |
| 4. Exchange reaction with salts | |
| Ba(OH) 2 + K 2 SO 4 ® 2KOH + BaSO 4 ¯ 3KOH+Fe(NO 3) 3 ® Fe(OH) 3 ¯ + 3KNO 3 |
-- |
| 5. Thermal decomposition. | |
| -- | Cu(OH) 2 - t ° ® CuO + H 2 O |
OXIDES
Classification
Oxides- these are complex substances consisting of two elements, one of which is oxygen.
| OXIDES | |
| Non-salt-forming | CO, N2O, NO |
| Salt-forming | Basic - these are metal oxides in which the latter exhibit a small oxidation state +1, +2 Na 2 O; MgO; CuO |
| |
Amphoteric (usually for metals with oxidation state +3, +4). Amphoteric hydroxides correspond to them as hydrates ZnO; Al 2 O 3; Cr 2 O 3; SnO2 |
| |
Acidic - these are oxides of non-metals and metals with an oxidation state from +5 to +7 SO2; SO 3; P2O5; Mn 2 O 7; CrO3 |
| |
Basic oxides bases correspond acidic- acids, amphoteric- and those and others |
Receipt
1. Interaction of simple and complex substances with oxygen:
2Mg + O 2 ® 2MgO
4P + 5O 2 ® 2P 2 O 5
S + O 2 ® SO 2
2CO + O 2 ® 2CO 2
2CuS + 3O 2 ® 2CuO + 2SO 2
CH 4 + 2O 2 ® CO 2 + 2H 2 O
4NH 3 + 5O 2 - cat. ® 4NO + 6H 2 O
2. Decomposition of some oxygen-containing substances (bases, acids, salts) when heated:
Cu(OH) 2 - t ° ® CuO + H 2 O
(CuOH) 2 CO 3 - t ° ® 2CuO + CO 2 + H 2 O
2Pb(NO 3) 2 - t ° ® 2PbO + 4NO 2 + O 2
2HMnO 4 - t °;H 2 SO 4 (conc.) ® Mn 2 O 7 + H 2 O
Chemical properties
| Basic oxides | Acidic oxides |
| 1. Interaction with water | |
| The base is formed: Na 2 O + H 2 O ® 2NaOH CaO + H 2 O ® Ca(OH) 2 |
Acid is formed: SO 3 + H 2 O ® H 2 SO 4 P 2 O 5 + 3H 2 O ® 2H 3 PO 4 |
| 2. Interaction with acid or base: | |
| When reacting with acid salt and water are formed MgO + H 2 SO 4 - t ° ® MgSO 4 + H 2 O CuO + 2HCl - t ° ® CuCl 2 + H 2 O |
When reacting with a base salt and water are formed CO 2 + Ba(OH) 2 ® BaCO 3 + H 2 O SO 2 + 2NaOH ® Na 2 SO 3 + H 2 O |
| Amphoteric oxides interact | |
| with acids as bases: ZnO + H 2 SO 4 ® ZnSO 4 + H 2 O |
with bases as acidic: ZnO + 2NaOH ® Na 2 ZnO 2 + H 2 O (ZnO + 2NaOH + H 2 O ® Na 2) |
| 3. The interaction of basic and acidic oxides with each other leads to salts. | |
| Na 2 O + CO 2 ® Na 2 CO 3 | |
| 4. Reduction to simple substances: | |
| 3CuO + 2NH 3 ® 3Cu + N 2 + 3H 2 O P 2 O 5 + 5C ® 2P + 5CO |
|
8. Main classes of inorganic compounds
8.1. Oxides
Oxides are complex substances consisting of two elements, one of which is oxygen, which is in the -2 oxidation state. Examples of oxides are Al 2 O 3 - aluminum oxide, SiO 2 - silicon oxide, NO - nitrogen oxide (II).
According to the international nomenclature, the compounds in question are called oxides indicating the oxidation state of the element if this element forms several oxides. When writing the name, the oxidation state is indicated by Roman numerals in brackets, for example, FeO - iron (II) oxide, Fe 2 O 3 - iron (III) oxide, SO 2 - sulfur oxide (IV), SO 3 - sulfur oxide (VI). Very often in the literature there are also trivial names for oxides - red lead (Pb 3 O 4), laughing gas (N 2 O), iron scale (Fe 3 O 4) and many others.
Oxides are divided into salt-forming and non-salt-forming. Salt-forming oxides are usually divided into basic, amphoteric and acidic.
It should be distinguished from oxides peroxides, for example H 2 O 2, Na 2 O 2 and superoxides KO 2, CsO 2. In these compounds, the oxidation state of oxygen is absolute value less than two and can be fractional.
8.1.1. Basic oxides
Basic oxides are formed only by metals; bases correspond to them as hydrates. For example, CaO, FeO, CuO are basic oxides, since they correspond to the bases Ca(OH) 2, Fe(OH) 2, Cu(OH) 2.
Preparation of basic oxides
The main oxides are obtained:
oxidation of metals with oxygen:
4 Li + O 2 2 Li 2 O,
2 Mg + O 2 2 MgO;
during oxidation alkali metals with oxygen, only lithium forms Li 2 O. Sodium gives peroxide (Na 2 O 2), the rest - superoxides (KO 2, RbO 2, CsO 2).
decomposition when heated oxygen compounds: hydroxides, nitrates, carbonates:
2 Fe(OH) 3 Fe 2 O 3 + 3 H 2 O,
2 Cu(NO 3) 2 2 CuO + 4 NO 2 + O 2,
CaCO 3 CaO + CO 2 (except alkali metal carbonates).
roasting sulfides:
2 ZnS + 3 O 2 2 ZnO + 2 SO 2.
Chemical properties of basic oxides
Oxides of alkali and alkaline earth metals react directly with water:
Na 2 O + H 2 O 2 NaOH,
BaO + H 2 O Ba(OH) 2 .
Basic oxides react with acids to form salt and water, for example:
FeO + 2 HCl FeCl 2 + H 2 O;
Basic oxides also react with acidic oxides:
BaO + CO 2 BaCO 3 ;
Basic oxides can also enter into redox reactions:
Fe 2 O 3 + 3 C 2 Fe + 3 CO,
CuO + H 2 Cu + H 2 O.
8.1.2. Acidic oxides
Acidic oxides are formed by nonmetals (SO 2, SO 3, CO 2, P 4 O 10, etc.) or transition metals in high oxidation states (for example, CrO 3, Mn 2 O 7).
Acidic oxides are prepared by the same methods as basic oxides. For example:
4 FeS 2 + 11O 2 2 Fe 2 O 3 + 8 SO 2,
Zn 2 (OH) 2 CO 3 2 ZnO + CO 2 + H 2 O,
as well as the decomposition of acids:
H 2 SiO 3 SiO 2 + H 2 O.
Chemical properties of acid oxides
1. Some acid oxides form acids when reacting with water:
SO 3 + H 2 O H 2 SO 4,
N2O5 + H2O2HNO3.
Some acid oxides are anhydrides acids For example, SO 3 is sulfuric acid anhydride, SO 2 is sulfurous acid anhydride, CO 2 is anhydride carbonic acid, P 4 O 10 is an anhydride of three acids (metaphosphoric HPO 3, orthophosphoric H 3 PO 4, pyrophosphoric H 4 P 2 O 7).
2. Acidic oxides interact with basic ones, forming salts:
SO 3 + CaO CaSO 4 .
3. Acidic oxides react with bases, forming salt and water:
CO 2 + 2 NaOH Na 2 CO 3 + H 2 O.
4. Like other types of oxides, acidic oxides can enter into redox reactions:
CO 2 + 2 Mg C + 2 MgO,
SO 2 + 2 H 2 S 3 S + 2H 2 O.
8.1.3. Amphoteric oxides
Amphoteric oxides have dual properties, i.e. depending on the conditions, they exhibit basic or acid properties. These include: ZnO, Al 2 O 3, BeO, Cr 2 O 3, etc. Amphoteric oxides do not interact with water, but react with both acids and bases. For example:
ZnO + H 2 SO 4 ZnSO 4 + H 2,
ZnO+ 2 NaOH + H 2 O Na 2.
Amphoteric oxides can interact with both basic and acidic oxides:
BeO + Na 2 O Na 2 BeO 2
ZnO + SO 3 ZnSO 4
Amphoteric oxides, when fused with alkalis or alkali metal carbonates, form salts:
ZnO + 2 NaOH Na 2 ZnO 2 + H 2 O,
Al 2 O 3 + 2 NaOH 2 NaAlO 2 + H 2 O.
Al 2 O 3 + Na 2 CO 3 2 NaAlO 2 + CO 2
The physical properties of oxides are very diverse. All basic and amphoteric oxides, as well as some acidic oxides (SiO 2, P 4 O 10, etc.) are solids. Many acid oxides at ordinary temperatures are gases (SO 2, CO 2) or liquids (Cl 2 O 7, Mn 2 O 7).
Properties non-salt-forming oxides(CO, NO, N 2 O, etc.) will be described in the following sections, which are devoted to the chemistry of the corresponding elements.
It is worth mentioning mixed oxides(Pb 2 O 3, Pb 3 O 4, etc.), in which the same element (Pb) is in various degrees oxidation. These compounds can also be classified as salts: Pb +2 Pb +4 O 3, Pb 2 +2 Pb +4 O 2.
8.2. Reasons
The basis from the point of view of the theory of electrolytic dissociation are compounds, upon dissociation of which hydroxo groups OH are formed as anions – . Not only metal hydroxides can have the properties of bases, but also some other substances, for example, NH 3, the molecule of which can attach a proton:
NH 3 + H + NH 4 +
8.2.1. Nomenclature of bases
According to the international nomenclature, the bases are usually called hydroxides of elements: NaOH – sodium hydroxide, CsOH – cesium hydroxide.
If an element can form several bases, then its oxidation state is indicated in the names in parentheses with a Roman numeral. For example, Fe(OH) 2 is iron (II) hydroxide, Fe(OH) 3 is iron (III) hydroxide.
Most bases are slightly soluble in water. Bases that dissolve in water are called alkalis. Alkalis are, for example, NaOH, KOH, Ba(OH) 2.
8.2.2. Getting grounds
In a general way the formation of bases is an exchange reaction between salt and alkali:
Cu(NO 3) 2 + 2 KOH Cu(OH) 2 + 2KNO 3 ,
Na 2 CO 3 + Ba(OH) 2 BaCO 3 + 2 NaOH.
Alkalis are formed by the interaction of alkali and alkaline earth metals, as well as their oxides, with water:
2 Na + 2 H 2 O 2 NaOH + H 2,
BaO + H 2 O Ba(OH) 2 .
In industry, alkalis are usually obtained by electrolysis of aqueous solutions of chlorides:
2 KCl + 2 H 2 O 2 KOH + H 2 + Cl 2.
8.2.3. Properties of bases
Alkali solutions change the color of indicators: colorless phenolphthalein turns crimson, methyl orange turns yellow, litmus turns blue.
Most bases that are slightly soluble in water decompose easily when heated:
Cu(OH) 2 CuO + H 2 O.
Alkalis are thermally stable and melt without decomposition. The exception is lithium hydroxide, which also decomposes when heated:
2 LiOH Li 2 O + H 2 O.
Both alkalis and insoluble bases can react with acids (neutralization reaction):
NaOH + 2 HCl NaCl + H 2 O,
2 Fe(OH) 3 + 3H 2 SO 4 Fe 2 (SO 4) 3 + 6H 2 O.
The interaction of bases with acidic and amphoteric oxides is discussed in section 8.1.
Amphoteric hydroxides react with both acids and bases. For example:
Al(OH) 3 + 3 HCl AlCl 3 + H 2 O
Al(OH) 3 + 3 NaOH Na 3
IN aqueous solutions containing alkali, along with 3–, there are other ions, in particular, 2–, –, 3–, etc. Aluminum hydroxo complexes also contain H 2 O molecules, which are usually not indicated in the formulas.
TO amphoteric hydroxides include Zn(OH) 2, Al(OH) 3, Cr(OH) 3, Be(OH) 2, Pb(OH) 2, etc.
In conclusion, it should be noted the ability of alkalis to interact with some non-metals and oxides:
6 KOH + 3 S K 2 SO 3 + 2 K 2 S + 3 H 2 O,
6 NaOH + 3 Cl 2 5 NaCl + NaClO 3 + 3 H 2 O,
2 KOH + NO 2 KNO 2 + KNO 3 + H 2 O.
The above reactions are redox reactions and are discussed in Section 7.
8.3. Acids
From the point of view of the theory of electrolytic dissociation, acid - chemical compound upon dissociation in water, only H + ions are formed as cations. The concepts of acids and bases arising from Arrhenius' theory of electrolytic dissociation are applicable only to aqueous solutions. The study of processes occurring in non-aqueous media, without the participation of a solvent, required significant additions and led to the emergence of various theories acids and bases.
8.3.1. Classification and nomenclature of acids
There are oxygen-free (H 2 S, HBr, HCl) and oxygen-containing (H 3 PO 4, HNO 3, HClO 3) acids.
In the free state, carbonic (H 2 CO 3) and sulfurous (H 2 SO 3) acids are unstable. There are also strong (H 2 SO 4, HNO 3, HCl, HBr, HI, HClO 4, etc.) and weak (H 2 S, H 2 CO 3, HCN, H 2 SO 3, HClO, etc.) acids.
The number of hydrogen ions formed during the dissociation of the formula unit of an acid determines its basicity (see section 9.8).
The names of oxygen-containing acids are derived from the name of the non-metal with the addition of the ending -naya, -vaya if the oxidation state of the nonmetal is maximum. As the oxidation state decreases, the suffixes change in the following way: -ovate, -iste, -ovate.
Examples of the names of some oxygen-containing acids are given in Table 7.1.
Table 7.1.
Names of some oxygen-containing acids
Titles oxygen-free acids consist of the name of a non-metal with the addition of a connecting vowel O and words - hydrogen. For example:
HF – hydrofluoric acid,
HCl – hydrochloric acid,
H 2 S – hydrosulfide acid.
8.3.2. Obtaining acids
1. Most oxygen-containing acids are obtained by reacting acid oxides with water (see §8.1: acid oxides).
2. To obtain water-insoluble acids, an indirect method is used (by the action of the acid on the corresponding salt):
Na 2 SiO 3 + H 2 SO 4 H 2 SiO 3 + Na 2 SO 4 .
3. Some oxygen-free acids are obtained by direct combination of non-metals with hydrogen:
H 2 + Cl 2 2 HCl,
or by the exchange reaction between salt and acid:
NaCl + H 2 SO 4 (conc.) HCl + NaHSO 4
8.3.3. General properties of acids
Acids are liquids (H 2 SO 4, HNO 3, HCl, etc.) or solids(H 3 PO 4, H 3 BO 3, etc.).
Solutions strong acids can destroy tissue and skin.
Acid solutions change color indicators, which is used for their qualitative detection. Used as indicators litmus(V neutral environment– violet, in acidic – red, in alkaline – blue), methyl orange(in a neutral environment - orange, in an acidic environment - red, in an alkaline environment - yellow) and others.
The strength of oxygen-free acids, for example, in the series HCl – HBr – HI, increases with increasing radius of the anion, since the anion larger radius holds the proton weaker, thereby facilitating the dissociation of the acid. Thus, in the main subgroups periodic table from top to bottom, the strength of oxygen-free acids increases with increasing radius of the central atom.
On the contrary, in the series HClO – HClO 2 – HClO 3 – HClO 4, with a decrease in the radius of the C1 z+ cation and an increase in its charge, the strength of oxygen-containing compounds increases.
The most important chemical properties of acids are:
interaction with basic and amphoteric oxides, bases and salts:
CaO + H 2 SO 4 CaSO 4 + H 2 O,
ZnO + H 2 SO 4 ZnSO 4 + H 2 O,
2 Fe(OH) 3 + 3 H 2 SO 4 Fe 2 (SO 4) 3 + 6 H 2 O,
BaCl 2 + H 2 SO 4 BaSO 4 + 2HCl;
the interaction of an acid with a base is a neutralization reaction;
interaction with metals to form salts and release hydrogen:
Mg + 2 HCl MgCl 2 + H 2,
Fe + H 2 SO 4 (diluted) FeSO 4 + H 2.
Hydrogen from acids is not replaced by metals that are standard electrode potentials(in the voltage series) to the right of hydrogen. When metals react with concentrated sulfuric acid and nitric acid, hydrogen is usually not released.
8.3.4. Properties of concentrated sulfuric acid
Concentrated sulfuric acid in reactions with metals it can be reduced to SO 2, S or H 2 S. The composition of the reduction products is determined by the activity of the metal, the acid concentration and the temperature of the reacting system. At ordinary temperatures, concentrated H 2 SO 4 does not react with gold and platinum, and some metals (Fe, Cr, Al) passivated in concentrated sulfuric acid.
Low-active metals (standing in the series of standard electrode potentials to the right of hydrogen) reduce concentrated sulfuric acid to SO 2:
Cu + 2 H 2 SO 4 (conc) CuSO 4 + SO 2 + 2 H 2 O.
Active metals (Ca, Mg, Zn, etc.) reduce concentrated H 2 SO 4 to free sulfur or H 2 S:
3 Zn + 4 H 2 SO 4(conc) 3 ZnSO 4 + S + 4H 2 O,
4 Ca + 5 H 2 SO 4(conc) 4 CaSO 4 + H 2 S + 4H 2 O.
When sulfuric acid reacts with non-metals, SO 2 is formed:
C + H 2 SO 4 (conc) 2 SO 2 + CO 2 + 2H 2 O,
S + 2 H 2 SO 4 (conc) 3 SO 2 + 2H 2 O,
2 P + 5 H 2 SO 4 (conc) 5 SO 2 + 2 H 3 PO 4 + 2H 2 O.
When concentrated H 2 SO 4 interacts with compounds containing metal cations that are in the lowest oxidation state, further oxidation of these metals occurs:
2 FeO + 4 H 2 SO 4 (conc) Fe 2 (SO 4) 3 + SO 2 + 4 H 2 O.
8.3.5. Properties of nitric acid
Nitric acid oxidizes most elements to their highest degree oxidation. Interaction of HNO 3 of various concentrations with metals various activities presented in the following diagram:
Thus, when concentrated HNO 3 interacts with low-active metals, NO 2 is formed:
Ag + 2 HNO 3(conc) AgNO 3 + NO 2 + H 2 O.
When dilute HNO 3 acts on low-active metals, NO is released:
3 Cu + 8 HNO 3(dil) 3 Cu(NO 3) 2 + 2 NO + 4 H 2 O,
and in the case of active metals, NH 4 NO 3 is formed:
4 Ca + 10 HNO 3(dil) 4 Ca(NO 3) 2 + NH 4 NO 3 + 3 H 2 O.
When concentrated nitric acid reacts with non-metals, NO is usually formed:
3 C + 4 HNO 3(dil) 3 CO 2 + 4 NO + 2H 2 O,
P + 5 HNO 3(dil) + 2 H 2 O 3H 3 PO 4 + 5NO.
When concentrated HNO 3 interacts with compounds containing metal cations that are in the lowest oxidation state, further oxidation of these metals occurs:
FeO + 4 HNO 3(dil) Fe(NO 3) 3 + NO 2 + 2 H 2 O.
It should be noted that the products of the reduction of nitric acid with metals are both nitrogen and even hydrogen, and, as a rule, a mixture of substances is formed. How metal is more active and the lower the concentration, the deeper it is restored.
8.4. Salts
8.4.1. Classification and nomenclature of salts
From the point of view of the theory of electrolytic dissociation, salts are compounds whose dissociation produces metal cations and OH – anions (see section 9).
Distinguish following types salts: medium, sour, basic, double, mixed And complex.
8.4.1.1. Medium salts
IN medium salts all hydrogen atoms of the corresponding acid are replaced by metal atoms. The dissociation equation for the average salt Na 2 SO 4 in a dilute solution is written as follows:
Na 2 SO 4 2 Na + + SO 4 2– ,
(NH 4) 2 Cr 2 O 7 2 NH 4 + + Cr 2 O 7 2– ,
it is indicated that the degree of dissociation tends to unity (α 1)
8.4.1.2. Acid salts
IN acid salts the hydrogen atoms of the corresponding acid are not completely replaced by the metal. Acid salt obtained by the reaction of incomplete neutralization of the acid:
H 2 SO 3 + NaOH (deficiency) NaHSO 3 + H 2 O,
or when an average salt reacts with an excess of acid:
Ca 3 (PO 4) 2 + 4 H 3 PO 4 (excess) 3 Ca (H 2 PO 4) 2.
To convert an acid salt to a medium salt, add a base:
NaHSO 4 + NaOH Na 2 SO 4 + H 2 O,
Ca(HCO 3) 2 + Ca(OH) 2 CaCO 3 + 2H 2 O.
The dissociation of an acid salt can be expressed by the equation:
NaHCO 3 Na + + HCO 3 – ; α 1
The HCO 3 - anion will dissociate to a small extent:
HCO 3 – H + + CO 3 2– .
Acid salts are formed by polybasic acids.
8.4.1.3. Basic salts
Basic salts are the product of incomplete substitution of OH groups - bases with acidic residues:
Mg(OH) 2 + HC1 (insufficient) MgOHС1 + H 2 O.
Basic salts form bases containing two or more hydroxo groups.
To convert a basic salt to a medium salt, you need to add an acid:
Mg(OH)Cl + HCl MgCl 2 + H 2 O.
The dissociation of a basic salt is expressed by the equation:
Mg(OH)Cl MgOH + + Cl – ; α 1.
The MgOH+ cation undergoes further dissociation as a weak electrolyte:
MgOH + Mg 2+ + OH – .
8.4.1.4. Double and mixed salts
Double salts- These are salts consisting of two different cations and one anion. Examples of double salts are: potassium alum KAl(SO 4) 2 and sylvinite KCl · NaCl.
The dissociation of a double salt in a dilute solution can be expressed by the equation:
KAl(SO 4) 2 K + + Al 3+ + 2SO 4 2– ; α 1.
Mixed salts are salts consisting of one cation and two different anions. Chloride of lime CaOCl 2, for example, is a salt of hypochlorous (HC1O) and hydrochloric (HC1) acids.
8.4.1.5. Complex salts
Part complex salts includes a complex ion (in the formulas it consists of square brackets), consisting of a central atom - a complexing agent, surrounded by several particles - molecules or ions (ligands). In dilute solutions, the complex salt dissociates as follows.
CHEMISTRY. INORGANIC COMPOUNDS
Inorganic compounds include compounds of all chemical elements, with the exception of most carbon compounds.
Acids, bases and salts. Acids are compounds that dissociate in water to release hydrogen ions (H+). These ions determine the characteristic properties of strong acids: sour taste and ability to interact with bases. Bases are substances that dissociate in water to release hydroxide ions (OH-). Salts are ionic compounds formed by the interaction of acids and bases:
The nomenclature is not organic compounds.
The nomenclature of most common inorganic compounds is based on the following rules.
Elements. The names of metals usually end in -iy (for example, sodium, potassium, aluminum, magnesium). The exception is metals that have been known since antiquity and received their names at the same time. These are, for example, iron, copper, gold. The names of non-metals usually end in -or (chlorine, boron, phosphorus), -od (hydrogen, oxygen, iodine) or -one (argon, neon). Knowing the names of the elements and the most common ions and using the rules below, you can name almost any inorganic compound.
Acids. The names of acids whose molecules do not contain oxygen end in hydrogen, for example hydrochloric (HCl), hydrobromic (HBr), hydroiodic (HI). The names of oxygen-containing acids depend on the degree of oxidation of the central element. The name of the acid in which this element has a lower oxidation state ends in -, for example, nitrous (HNO2), sulfurous (H2SO3), and the name of the larger acid ends in -, for example, nitric (HNO3), sulfuric (H2SO4). Using chlorine as an example, let us consider the case when the element forms more than two oxygen-containing acids. Their names are formed as follows: hypochlorous acid, HClO; chloride, HClO2; hypochlorous, HClO3; chlorine, HClO4. The oxidation state of chlorine here is +1, +3, +5 and +7, respectively. Names of acids whose molecules contain different quantities water, differ from each other by the prefixes ortho-, hypo-, pyro- and meta- (in order of decreasing water content):
Positively charged ions. The names of these ions are formed as follows: after the word ion, indicate the name of the element and in Roman numerals the degree of its oxidation. For example, Cu2+ is a copper(II) ion, Cu+ is a copper(I) ion. The names of some positive ions end in -onium: ammonium, NH4+; hydronium, H3O+.
Negatively charged ions. The names of monoatomic negatively charged ions (and, accordingly, salts) obtained from oxygen-free acids end in -ide: chloride ion, Cl-; bromide ion, Br-. The names of ions (and, accordingly, salts) obtained from oxygen-containing acids, in which the central element has a lower oxidation state, end in -it: sulfite, SO32-; nitrite, NO2-; phosphite, PO33-; and the larger one - at -at: sulfate, SO42-; nitrate, NO3-; phosphate, PO43-. The names of ions obtained from partially neutralized acids are formed by adding the word acidic or the prefixes hydro- or bi- to the name of the ion: hydrocarbonate (bicarbonate), HCO3-; acid sulfate, HSO4-.
Salts and covalent compounds. For salts and covalent compounds use the names of the ions they contain: sodium chloride, NaCl; sodium hydroxide, NaOH. If an element can have several oxidation states, then after its name the degree of oxidation in this compound is indicated in Roman numerals: iron(II) sulfate, FeSO4; iron(III) sulfate, Fe2(SO4)3. If a compound is formed by two non-metals, then the prefixes di-, tri-, tetra-, penta-, etc. are used to indicate the number of their atoms. For example, carbon disulfide, CS2; phosphorus pentachloride, PCl5, etc.
Collier's Encyclopedia. - Open Society. 2000 .
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Inorganic substances are divided into four main classes: oxides, acids, bases, salts. Let us present the classes of connections known to us in the form of a single diagram:
The division of substances into classes is quite arbitrary. For example, we know that acids are divided into mono-, bi- and tribasic, but they are not usually separated into separate classes of compounds. Likewise, the strong and the strong are not separate classes. weak acids. The same is true for grounds. There is an important connection between classes, which is called genetic. This connection lies in the fact that substances of other classes can be obtained from substances of one class. There are two main ways of genetic connections between substances: one of them begins with metals, the other with non-metals. For example, calcium sulfate CaSO 4 can be obtained either from the calcium metal, or in another way - from the non-metal sulfur:
On the other hand, from salt one can again come to metal and non-metal:
At the same time, there are other ways of interconversion of compounds different classes. Thus, the genetic relationships between different classes of compounds are very diverse.
Oxides and their classification.
As we already know, oxides can be acidic or basic. This division forms the basis for their classification.
Most acidic oxides react well with water to produce acid. For example, the sour taste of simple carbonated water is explained by the formation of carbonic acid H 2 CO 3 from the acidic oxide CO 2:
CO 2 + H 2 O = H 2 CO 3 (carbonic acid)
In the simplest cases, the formula of the resulting acid can be easily obtained from the formula of the acid oxide by simple addition. For example:
The resulting silicic acid salt can be converted into silicic acid itself by adding another acid:
Na 2 SiO 3 + 2 HCl = H 2 SiO 3 + 2 NaCl
Thus, acid oxide always corresponds to a certain acid:
CO 2 (carbon monoxide) – H 2 CO 3 (carbonic acid);
SO 3 (sulfur oxide VI) – H 2 SO 4 (sulfuric acid);
SiO 2 (silicon oxide) – H 2 SiO 3 (silicic acid).
Since the reaction with bases is general For everyone acid oxides, they can be defined as follows:
Oxides that react with bases to form salt and water are called ACID OXIDES.
Acidic oxides are formed mainly by non-metals. There are only two metal oxides to remember that are also acidic. These are oxides of chromium and manganese, in which the metals have the greatest of all possible degree oxidation:
CrO 3 (chromium VI oxide) – H 2 CrO 4 (chromic acid);
Mn 2 O 7 (manganese oxide VII) – HmnO 4 (manganese acid).
Basic oxides are formed only by metals. Some of them react easily with water, giving the corresponding base:
Li 2 O + H 2 O = 2 LiOH (base – lithium hydroxide).
Another example is the well-known reaction of producing slaked lime from calcium oxide and water.
CaO + H 2 O = Ca(OH) 2 (base – calcium hydroxide).
There is, however a large number of insoluble basic oxides. They are classified specifically as basic oxides due to their reactions with acids:
ZnO + H 2 O = reaction does not occur (ZnO is insoluble in water);
ZnO + 2 HCl = ZnCl 2 (salt) + H 2 O
The last reaction is similar to the reaction neutralization between acid (HCl) and zinc hydroxide Zn(OH) 2, which could be obtained from ZnO if zinc oxide were dissolved in water:
= Zn(OH) 2
Zn(OH) 2 + 2 HCl = ZnCl 2 (salt) + H 2 O
Each basic oxide corresponds to a specific base:
MgO (magnesium oxide) – Mg(OH) 2 (magnesium hydroxide);
Fe 2 O 3 (iron III oxide) – Fe(OH) 3 (iron III hydroxide);
Na 2 O (sodium oxide) – NaOH (sodium hydroxide).
Thus, general property basic oxides is the ability to react with acids to form salt and water.
Oxides that react with ACIDS to form salt and water are called BASIC OXIDES.
Oxides of chromium and manganese, in which the metal has a lower oxidation state, are ordinary basic oxides (like the oxides of all other metals). Here are the hydroxides that correspond to them:
CrO (chromium II oxide) – Cr(OH) 2 (chromium II hydroxide);
MnO (manganese II oxide) – Mn(OH) 2 (manganese II hydroxide).
Chromium (II) compounds are extremely unstable and quickly transform into chromium (III) compounds. We have already become acquainted with the use of many interesting oxides in Chapter 6, “Oxygen.”
Acids. Classification of acids. Chemical properties.
All acids, regardless of their origin, have a common property - they contain reactive hydrogen atoms. In this regard, acids can be given following definition:
Acid is compound, the molecule of which contains one or more hydrogen atoms and an acid residue.
The properties of acids are determined by the fact that they are able to replace hydrogen atoms in their molecules with metal atoms. For example:
Using sulfuric acid as an example, let us consider its formation from the acidic oxide SO 3, and then the reaction of sulfuric acid with magnesium. We know the valences of all elements participating in the reaction, so we will write the compounds in the form of structural formulas:
These examples make it easy to see the connection between acid oxide SO 3, acid H 2 SO 4 and salt MgSO 4. One is “born” from the other, and the sulfur atom and oxygen atoms move from a compound of one class (acid oxide) to compounds of other classes (acid, salt).
Acids are classified according to the following criteria: a) by the presence or absence of oxygen in the molecule b) by the number of hydrogen atoms. According to the first sign, acids are divided into oxygen-containing And oxygen-free
Table 4.1. Classification of acids by composition.
According to the number of hydrogen atoms that can be replaced by a metal, all acids are divided into monobasic(with one hydrogen atom), dibasic(with 2 hydrogen atoms) and tribasic(with 3 hydrogen atoms), as shown in table. 4.2:
Table 4.2. Classification of acids according to the number of hydrogen atoms.
The term “monobasic acid” arose because to neutralize one molecule of such an acid, one base is required, i.e. one molecule of any simple base such as NaOH or KOH:
HNO 3 + NaOH = NaNO 3 + H 2 O
HCl + KOH = KCl + H2O
A dibasic acid requires two bases for its neutralization, and a tribasic acid requires three bases:
H 2 SO 4 + 2 NaOH = Na 2 SO 4 + 2 H 2 O
H 3 PO 4 + 3 NaOH = Na 3 PO 4 + 3 H 2 O
Let's consider the most important chemical properties of acids.
1. Effect of acid solutions on indicators . Almost all acids (except silicic acid) are highly soluble in water. Solutions of acids in water change the color of special substances - indicators. The presence of acid is determined by the color of the indicators. Indicator litmus painted with solutions acids V red color, indicator methyl orange– also in red.
Indicators are substances of complex structure. In solutions of bases and in neutral solutions they have a different color than in solutions of acids.
2. Interaction of acids with bases . This reaction is called a neutralization reaction. An acid reacts with a base to form a salt, in which the acidic residue is always found unchanged. The second product of the neutralization reaction is necessarily water. For example:
|
base | ||||||
For neutralization reactions, it is sufficient that at least one of the reactants is soluble in water. Since almost all acids are soluble in water, they enter into neutralization reactions not only with soluble but also with insoluble bases. The exception is silicic acid, which is poorly soluble in water and therefore can only react with soluble bases - such as NaOH and KOH:
H 2 SiO 3 + 2 NaOH = Na 2 SiO 3 + 2H 2 O
3. Interaction of acids with basic oxides . Since basic oxides are the closest relatives of bases, acids also enter into neutralization reactions with them:
As in the case of reactions with bases, acids form salt and water with basic oxides. The salt contains the acid residue of the acid that was used in the neutralization reaction.
For example, phosphoric acid is used to clean iron from rust (iron oxides). Phosphoric acid, removing its oxide from the surface of the metal, reacts very slowly with the iron itself. Iron oxide turns into a soluble salt, FePO 4, which is washed off with water along with acid residues.
4. Interaction of acids with metals . For the interaction of acids with a metal, certain conditions must be met (unlike the reactions of acids with bases and basic oxides, which almost always occur).
Firstly, the metal must be sufficiently active (reactive) with respect to acids. For example, gold, silver, mercury and some other metals do not react with acids. Metals such as sodium, calcium, zinc react very actively, releasing hydrogen gas and a large amount of heat.
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not formed | ||||||
According to their reactivity towards acids, all metals are located in metal activity series(Table 4-3). On the left are the most active metals, on the right are the inactive ones. The further to the left a metal is in the activity series, the more intensely it interacts with acids.
Table 4.3. Metal activity series.
Secondly, the acid must be sufficient strong, to react even with the metal from the left side of the table. 4-3. The strength of an acid is understood as its ability to release hydrogen ions H +.
For example, plant acids (malic, citric, oxalic, etc.) are weak acids and react very slowly with metals such as zinc, chromium, iron, nickel, tin, lead (although they are capable of reacting with bases and metal oxides).
On the other hand, such strong acids such as sulfuric or hydrochloric (hydrochloric) are capable of reacting with all metals from the left side of the table. 4.3.
In this regard, there is another classification of acids - by strength. In Table 4.4, in each column the acid strength decreases from top to bottom.
Table 4.4. Classification of acids into strong and weak acids.
It should be remembered that there is one important exception in the reactions of acids with metals. When metals interact with nitric acid hydrogen is not released. This is due to the fact that Nitric acid contains in its molecule a strong oxidizing agent - nitrogen in the oxidation state +5. Therefore, the more active oxidizing agent N +5 reacts with metals first, and not H +, as in other acids. Any hydrogen released in some quantity is immediately oxidized and is not released as a gas. The same is observed for reactions concentrated sulfuric acid, in the molecule of which sulfur S +6 also acts as the main oxidizing agent. The composition of the products in these redox reactions depends on many factors: metal activity, acid concentration, temperature. For example:
Cu + 4 HNO 3 (conc.) = Cu(NO 3) 2 + 2 NO 2 + 2 H 2 O
3 Cu + 8HNO 3 (diluted) = 3 Cu(NO 3) 2 + 2 NO + 4 H 2 O
8 K + 5 H 2 SO 4 (conc.) = 4 K 2 SO 4 + H 2 S + 4 H 2 O
3 Zn + 4 H 2 SO 4 (conc.) = 3 ZnSO 4 + S + 4 H 2 O
There are metals that react with diluted acids, but does not react with concentrated ones (i.e. anhydrous) acids – sulfuric acid and nitric acid.
These metals - Al, Fe, Cr, Ni and some others - upon contact with anhydrous acids, are immediately covered with oxidation products (passivated). Oxidation products that form strong films can be dissolved in aqueous solutions of acids, but are insoluble in concentrated acids.
This circumstance is used in industry. For example, concentrated sulfuric acid is stored and transported in iron barrels.
To the group inorganic substances include all substances that are inherently opposite to organic. That is, this means that there is no carbon in the composition of inorganic substances. The exceptions are carbides, cyanides, carbonates and carbon monoxide.
All inorganic substances are divided into two large groups:
Simple substances
Complex substances.
Simple substances
These are substances consisting of atoms of one element.
Divided into two large groups:
Metals,
Non-metals.
Metals
Metals - the group is called simple bodies, having known characteristic properties, which in their typical representatives sharply distinguish metals from other chemical elements.
IN physically These are for the most part solid bodies at ordinary temperatures, opaque (in a thick layer), possessing a certain shine, malleable, malleable, good conductors of heat and electricity, and so on. IN chemically They are characterized by the ability to form basic oxides with oxygen, which, when combined with acids, give salts.
Metals include: iron, copper, zinc, calcium, potassium, aluminum, gold, silver, sodium, tin, beryllium, etc.
Nonmetals
Nonmetals are a group of simple bodies that have known characteristic properties that sharply distinguish nonmetals from other chemical elements.
In physical terms, these are various solid bodies: solid, liquid and gaseous.
Non-metals include: hydrogen, oxygen, nitrogen, phosphorus, sulfur, carbon, argon, neon, etc.
Complex substances
These are substances consisting of atoms of two or more elements. Divided into four large groups:
Reasons
Acids
Oxides
Oxides are compounds of various chemical elements with oxygen.
Depending on the chemical properties, they are distinguished:
Salt-forming oxides,
Non-salt-forming oxides.
Salt-forming oxides are oxides that produce salts when reacting with other elements. They are divided into 3 groups:
Basic oxides (sodium oxide Na2O, copper oxide (II) CuO),
Acidic oxides (sulfur oxide SO3, nitrogen oxide NO2),
Amphoteric oxides (oxide zinc ZnO, aluminum oxide Al2O3)
Non-salt-forming oxides are oxides that do not produce salts when interacting with other elements. They usually break down into gas and water.
Example: carbon monoxide CO, nitric oxide NO.
Reasons
These are substances whose molecules consist of metal molecules and a hydrox group - OH. Bases are formed by the interaction of a number of metals (sodium, potassium) or some oxides (calcium oxide CaO) with water.
Example: NaOH, Ca(OH)2, Al(OH)3, Fe(OH)3.
Acids
A group of compounds with a known, fairly specific chemical function. This function is expressed in such typical representatives of this group as sulfuric acid H2SO4, nitric acid HNO3, hydrochloric acid HCl and others.
There are a large number of classifications of acids, including special interest represent two - by oxygen content and by belonging to the class of chemical compounds.
Classification of acids according to oxygen content:
Oxygen-free (HCl, H2S, HBr)
Classification of acids according to their class of chemical compounds:
Inorganic (HBr,HCl, H2S, HNO3, H2SO4),
Organic (HCOOH, CH3COOH).
Salts
It is a chemical compound formed by the interaction of an acid and a base.
Example: NaCl, KNO3, CuSO4, Ca3(PO4)2.