Potassium, sodium or lithium may react with water. In this case, compounds related to hydroxides are found in the reaction products. The properties of these substances, the peculiarities of the occurrence of chemical processes in which bases participate, are determined by the presence of a hydroxyl group in their molecules. Thus, in electrolytic dissociation reactions, bases are split into metal ions and OH - anions. We will look at how bases interact with non-metal oxides, acids and salts in our article.
To correctly name the base, you need to add the word hydroxide to the name of the metal element. Let's give specific examples. Aluminum base belongs to amphoteric hydroxides, the properties of which we will consider in the article. The obligatory presence in the molecules of bases of a hydroxyl group associated with the metal cation by an ionic type of bond can be determined using indicators, for example, phenolphthalein. In an aqueous environment, an excess of OH - ions is determined by the change in color of the indicator solution: colorless phenolphthalein becomes crimson. If a metal exhibits multiple valencies, it can form multiple bases. For example, iron has two bases, in which it is equal to 2 or 3. The first compound is characterized by the characteristics of the second - amphoteric. Therefore, the properties of higher hydroxides differ from compounds in which the metal has a lower degree of valency.
Bases are solid substances that are resistant to heat. In relation to water, they are divided into soluble (alkalis) and insoluble. The first group is formed by chemically active metals - elements of the first and second groups. Substances that are insoluble in water consist of atoms of other metals whose activity is inferior to sodium, potassium or calcium. Examples of such compounds are iron or copper bases. The properties of hydroxides will depend on which group of substances they belong to. Thus, alkalis are thermally stable and do not decompose when heated, while bases insoluble in water are destroyed under the influence of high temperature, forming oxide and water. For example, copper base decomposes as follows:
Cu(OH) 2 = CuO + H 2 O
The interaction between two important groups of compounds - acids and bases - is called in chemistry a neutralization reaction. This name can be explained by the fact that chemically aggressive hydroxides and acids form neutral products - salts and water. Being, in fact, an exchange process between two complex substances, neutralization is characteristic of both alkalis and water-insoluble bases. Let us give the equation for the neutralization reaction between caustic potassium and chloride acid:
KOH + HCl = KCl + H2O
An important property of alkali metal bases is their ability to react with acidic oxides, resulting in salt and water. For example, by passing carbon dioxide through sodium hydroxide, you can obtain its carbonate and water:
2NaOH + CO 2 = Na 2 CO 3 + H 2 O
Ion exchange reactions include the interaction between alkalis and salts, which occurs with the formation of insoluble hydroxides or salts. Thus, by pouring the solution dropwise into a solution of copper sulfate, you can obtain a blue jelly-like precipitate. This is a copper base, insoluble in water:
CuSO 4 + 2NaOH = Cu(OH) 2 + Na 2 SO 4
The chemical properties of hydroxides, insoluble in water, differ from alkalis in that when slightly heated they lose water - they dehydrate, turning into the form of the corresponding basic oxide.
If an element or can react with both acids and alkalis, it is called amphoteric. These include, for example, zinc, aluminum and their bases. The properties of amphoteric hydroxides make it possible to write their molecular formulas both in the form of a hydroxo group and in the form of acids. Let us present several equations for the reactions of aluminum base with chloride acid and sodium hydroxide. They illustrate the special properties of hydroxides, which are amphoteric compounds. The second reaction occurs with the decomposition of alkali:
2Al(OH) 3 + 6HCl = 2AlCl 3 + 3H 2 O
Al(OH) 3 + NaOH = NaAlO 2 + 2H 2 O
The products of the processes will be water and salts: aluminum chloride and sodium aluminate. All amphoteric bases are insoluble in water. They are extracted as a result of the interaction of appropriate salts and alkalis.
In industries requiring large volumes of alkalis, they are obtained by electrolysis of salts containing cations of active metals of the first and second groups of the periodic table. The raw material for the extraction of, for example, sodium hydroxide is a solution of table salt. The reaction equation will be:
2NaCl + 2H 2 O = 2NaOH + H 2 + Cl 2
Bases of low-active metals are obtained in the laboratory by reacting alkalis with their salts. The reaction is an ion exchange type and ends with the precipitation of a base. A simple way to produce alkalis is a substitution reaction between the active metal and water. It is accompanied by heating of the reacting mixture and is of the exothermic type.
The properties of hydroxides are used in industry. Alkalies play a special role here. They are used as kerosene and gasoline purifiers, for making soap, processing natural leather, as well as in technologies for the production of artificial silk and paper.
Since d-metal oxides are insoluble in water, their hydroxides are obtained indirectly using exchange reactions between their salts and alkali solutions:
ZnCl 2 + 2NaOH = Zn(OH) 2 + 2NaCl;
MnCl 2 + 2NaOH = Mn(OH) 2 + 2NaCl (in the absence of oxygen);
FeSO 4 + 2KOH = Fe(OH) 2 + K 2 SO 4 (in the absence of oxygen).
Hydroxides of d-elements in lower oxidation states are weak bases; They are insoluble in water, but dissolve well in acids:
Cu(OH) 2 + 2HCl = CuCl 2 + H 2 O
Cu(OH) 2 + H 2 SO 4 = CuSO 4 + H 2 O
Hydroxides of d-elements in intermediate oxidation states and zinc hydroxide dissolve not only in acids, but also in excess alkali solutions with the formation of hydroxo complexes (i.e., they exhibit amphoteric properties), for example:
Zn(OH) 2 + H 2 SO 4 = ZnSO 4 + 2H 2 O;
Zn(OH) 2 + 2NaOH = Na 2;
Cr(OH) 3 + 3HNO 3 = Cr(NO 3) 3 + 3H 2 O;
Cr(OH) 3 + 3KOH = K 3.
In higher oxidation states, transition metals form hydroxides, which exhibit acidic properties or amphoteric properties with a predominance of acidic ones:
With an increase in the degree of oxidation of an element, the basic properties of oxides and hydroxides weaken, and the acidic properties increase.
Therefore, across the period from left to right, there is an increase in the acidic properties of d-metal hydroxides in higher oxidation states up to the Mn subgroup, then the acidic properties weaken:
Sc(OH) 3 - TiO 2 xH 2 O - V 2 O 5 xH 2 O - H 2 CrO 4 - HMnO 4
Strengthening acid properties
Fe(OH) 3 - Co(OH) 2 - Cu(OH) 2 - Zn(OH) 2
Slow weakening of acid properties
Let us consider the change in the properties of d-metal hydroxides in subgroups. From top to bottom in the subgroup, the basic properties of hydroxides of d-elements in higher oxidation states increase, while the acidic properties decrease. For example, for the sixth group of d-metals:
H 2 CrO 4 - sharp - MoO 3 H 2 O - weak - WO 3 H 2 O
Acid properties are reduced
Connections of d-elements in lower oxidation states they exhibit, mostly, reducing properties, especially in an alkaline environment. Therefore, for example, hydroxides Mn(+2), Cr(+2), Fe(+2) are very unstable and are quickly oxidized by atmospheric oxygen:
2Mn(OH)2 + O2 + 2H2O = 2Mn(OH)4;
4Cr(OH) 2 + O 2 + 2H 2 O = 4Cr(OH) 3
In order to convert cobalt (II) or nickel (II) hydroxide into Co(OH) 3 or Ni(OH) 3, it is necessary to use a stronger oxidizing agent - for example, hydrogen peroxide H 2 O 2 in an alkaline medium or bromine Br 2:
2Co(OH) 2 + H 2 O 2 = 2Co(OH) 3;
2 Ni(OH) 2 + Br 2 +2NaOH = 2 Ni(OH) 3 + 2NaBr
Derivatives of Ti(III), V(III), V(II), Cr (II) are easily oxidized in air, some salts can be oxidized even with water:
2Ti 2 (SO 4) 3 + O 2 + 2H 2 O = 4TiOSO 4 + 2H 2 SO 4;
2CrCl 2 + 2H 2 O = 2Cr(OH) Cl 2 + H 2
Compounds of d-elements in higher oxidation states (from +4 to +7) usually exhibit oxidizing properties. However, Ti(IV) and V(V) compounds are always stable and therefore have relatively weak oxidizing properties:
TiOSO 4 + Zn + H 2 SO 4 = Ti 2 (SO 4) 3 + ZnSO 4 + H 2 O;
Na 3 VO 4 + Zn + H 2 SO 4 = VOSO 4 + ZnSO 4 + H 2 O
Reduction occurs under harsh conditions - with atomic hydrogen at the moment of its release (Zn + 2H + = 2H + Zn 2+).
And chromium compounds in higher oxidation states are strong oxidizing agents, especially in an acidic environment:
K2Cr2O7 + 3SO2 + H2SO4 = Cr2(SO4)3 + K2SO4 + H2O;
2CrO 3 + C 2 H 5 OH = Cr 2 O 3 + CH 3 COH + H 2 O
Mn(VI), Mn(VII) and Fe(VI) compounds exhibit even stronger oxidizing properties:
2KMnO 4 + 6KI + 4H 2 O = 2MnO 2 + 3I 2 + 8KOH;
4K 2 FeO 4 + 10H 2 SO 4 = 2Fe 2 (SO 4) 3 + 3O 2 +10H 2 O+ 4K 2 SO 4
Thus, the oxidizing properties of compounds of d-elements in higher oxidation states increase across the period from left to right.
The oxidizing ability of compounds of d-elements in higher oxidation states in the subgroup from top to bottom weakens. For example, in the chromium subgroup: potassium bichromate K 2 Cr 2 O 7 interacts even with such a weak reducing agent as SO 2 . To reduce molybdate or tungstate ions, a very strong reducing agent is required, for example, a hydrochloric acid solution of tin (II) chloride:
K 2 Cr 2 O 7 + SO 2 + H 2 SO 4 = Cr 2 (SO 4) 3 + K 2 SO 4 + H 2 O
3 (NH 4) 2 MoO 4 + HSnCl 3 + 9HCl = MoO 3 MoO 5 + H 2 SnCl 6 + 4H 2 O + 6NH 4 Cl
The last reaction occurs when heated, and the oxidation state of the d-element decreases very slightly.
Compounds of d-metals in intermediate oxidation states exhibit redox duality. For example, iron (III) compounds, depending on the nature of the partner substance, can exhibit reducing agent properties:
2FeCl3 + Br2 + 16KOH = 2K2FeO4 + 6KBr + 6KCl +8H2O,
and oxidizing properties:
2FeCl 3 + 2KI = 2FeCl 2 + I 2 + 2KCl.
or = hydrogen + base (if the base is not soluble in water)
The reaction occurs only if
the metal is in the activity series up to hydrogen.
Base - a complex substance in which each metal atom is associated with one or more hydroxo groups.
in oxidation states +1 And +2 show basic properties ,
Fill out the table:
metals of the main subgroups I - III groups
Comparison Questions
I group
II group
2. Physical properties.
III group
Interaction:
a) with water
b) with acids
c) with acid oxides
d) with amphoteric oxides
d) with alkalis
5. Hydroxide formula.
6. Physical properties
Interaction:
a) action on indicators
b) with acids
c) with acid oxides
d) with salt solutions
e) with non-metals
e) with alkalis
h) attitude to heating
The properties of oxides and hydroxides in the period change from basic through amphoteric to acidic, because the positive oxidation state of elements increases.
Na 2 O , Mg +2 O , Al 2 O 3
basic amphoteric
Na +1 O N , Mg +2 (O N ) 2 , Al +3 (O N ) 3
alkali Weak Amphoteric
base hydroxide
In the main subgroups, the basic properties of oxides and hydroxides increase from top to bottom .
Metal compounds I A groups
Alkali metal oxides
General formula Meh 2 ABOUT
Physical properties: Solid, crystalline substances, highly soluble in water.
Li 2 O, Na 2 O - colorless, K 2 O, Rb 2 O - yellow, Cs 2 O - orange.
Methods of obtaining:
Oxidation of the metal produces only lithium oxide
4 Li + O 2 → 2 Li 2 O
(in other cases, peroxides or superoxides are obtained).
All oxides (except Li 2 O) are obtained by heating a mixture of peroxide (or superoxide) with an excess of metal:
Na 2 O 2 + 2Na → 2Na 2 O
KO 2 + 3K → 2K 2 O
Chemical properties
Typical basic oxides:
React with water, forming alkalis: Na 2 O + H 2 O →
2. React with acids, forming salt and water: Na 2 O + H Cl →
3. Interact with acid oxides, forming salts: Na 2 O + SO 3 →
4. Interact with amphoteric oxides, forming salts: Na 2 O + ZnO → Na 2 ZnO 2
Alkali metal hydroxides
General formula – MeOH
Physical properties: White crystalline substances, hygroscopic, highly soluble in water (with the release of heat). The solutions are soapy to the touch and very caustic.
NaOH – sodium hydroxide
KOH – caustic potassium
Strong bases - Alkalis. The main properties are enhanced in the following order:
LiOH → NaOH → KOH → RbOH → CsOH
Methods of obtaining:
1. Electrolysis of chloride solutions:
2NaCl + 2H2O → 2NaOH + H 2 + Cl 2
2. Exchange reactions between salt and base:
K 2 CO 3 + Ca(OH) 2 → CaCO 3 + 2KOH
3. Interaction of metals or their basic oxides (or peroxides and superoxides) with water:
2 Li + 2 H 2 O → 2 LiOH + H2
Li 2 O + H 2 O → 2 LiOH
Na 2 O 2 + 2 H 2 O → 2 NaOH + H 2 O 2
Chemical properties
1. Change the color of the indicators:
Litmus - blue
Phenolphthalein – to raspberry
Methyl orange - to yellow
2. Interact with all acids.
NaOH + HCl → NaCl + H2O
3. Interact with acid oxides.
2NaOH + SO 3 → Na 2 SO 4 + H 2 O
4. Interact with salt solutions if gas or sediment is formed.
2 NaOH + CuSO 4 → Cu(OH) 2 ↓ + Na 2 SO 4
5. Interact with some non-metals (sulfur, silicon, phosphorus)
2 NaOH + Si + H 2 O → Na 2 SiO 3 + 2H 2
6. Interact with amphoteric oxides and hydroxides
2 NaOH + Zn O + H 2 O → Na 2 [Zn (OH) 4 ]
2 NaOH + Zn (OH) 2 → Na 2 [Zn (OH) 4 ]
7. When heated, they do not decompose, except for LiOH.
II groups
Metal oxides II A groups
General formula MeO
Physical properties: Solid, white crystalline substances, slightly soluble in water.
Methods of obtaining:
Oxidation of metals (except Ba, which forms peroxide)
2Ca + O 2 → 2CaO
2) Thermal decomposition of nitrates or carbonates
CaCO 3 → CaO + CO 2
2Mg(NO 3) 2 → 2MgO + 4NO 2 + O 2
Chemical properties
BeO – amphoteric oxide
Oxides Mg, Ca, Sr, Ba – basic oxides
They interact with water (except BeO), forming alkalis (Mg (OH) 2 - weak base):
CaO + H 2 O →
2. React with acids, forming salt and water: CaO + H Cl →
3. Interact with acid oxides, forming salts: CaO + SO 3 →
4. BeO interacts with alkalis: BeO + 2 NaOH + H 2 O → Na 2 [Be (OH) 4 ]
Metal hydroxides II A groups
General formula – Me(OH) 2
Physical properties: White crystalline substances are less soluble in water than alkali metal hydroxides. Be(OH) 2 – insoluble in water.
The main properties are enhanced in the following order:
Be(OH) 2 → Mg (HE) 2 → Ca (HE) 2 → Sr (HE) 2 → B a (HE) 2
Methods of obtaining:
Reactions of alkaline earth metals or their oxides with water:
Ba + 2 H 2 O → Ba (OH) 2 + H 2
CaO (quicklime) + H 2 O → Ca (OH) 2 (slaked lime)
Chemical properties
Be(OH) 2 – amphoteric hydroxide
Mg (OH) 2 – weak base
Ca(OH) 2, Sr (OH) 2, Ba(OH) 2 - strong bases - alkalis.
Change the color of the indicators:
Litmus - blue
Phenolphthalein – to raspberry
Methyl orange - to yellow
2. React with acids, forming salt and water:
Be(OH) 2 + H 2 SO 4 →
3. Interact with acid oxides:
Ca(OH) 2 + SO 3 →
4. Interact with salt solutions if gas or sediment is formed:
Ba(OH) 2 + K 2 SO 4 →
Beryllium hydroxide reacts with alkalis:
Be(OH) 2 + 2 NaOH → Na 2 [Be(OH) 4 ]
When heated, they decompose: Ca(OH) 2 →
Compounds of metals of the main subgroup III groups
Aluminum connections
Aluminium oxide
Al 2 O 3
O = Al – O – Al = O
Physical properties: Alumina, corundum, colored – ruby (red), sapphire (blue).
Solid refractory (t° pl. = 2050 ° C) substance; exists in several crystal modifications.
Methods of obtaining:
Combustion of aluminum powder: 4 Al + 3 O 2 → 2 Al 2 O 3
Decomposition of aluminum hydroxide: 2 Al (OH) 3 → Al 2 O 3 + 3 H 2 O
Chemical properties
Al 2 O 3 - amphoteric oxide with predominant basic properties; does not react with water.
As a basic oxide: Al 2 O 3 + 6 HCl → 2 AlCl 3 + 3 H 2 O
As an acidic oxide: Al 2 O 3 + 2 NaOH + 3 H 2 O → 2 Na [Al (OH) 4 ]
2) Alloyed with alkalis or alkali metal carbonates:
Al 2 O 3 + Na 2 CO 3 → 2 NaAlO 2 (sodium aluminate) + CO 2
Al 2 O 3 + 2 NaOH → 2 NaAlO 2 + H 2 O
Aluminum hydroxide Al ( OH ) 3
Physical properties: white crystalline substance,
insoluble in water.
Methods of obtaining:
1) Precipitation from salt solutions with alkalis or ammonium hydroxide:
AlCl 3 + 3NaOH → Al(OH) 3 + 3NaCl
Al 2 (SO 4) 3 + 6NH 4 OH → 2Al(OH) 3 + 3(NH 4) 2 SO 4
Al 3+ + 3 OH ¯ → Al (OH) 3 (white gelatinous)
2) Weak acidification of aluminate solutions:
Na + CO 2 → Al(OH) 3 + NaHCO 3
Chemical properties
Al ( OH ) 3 - A mphoteric hydroxide :
1) Reacts with acids and alkali solutions:
As a base Al (OH) 3 + 3 HCl → AlCl 3 + 3 H 2 O
As acid Al (OH) 3 + NaOH → Na [Al (OH) 4 ]
(sodium tetrahydroxyaluminate)
When heated, it decomposes: 2 Al (OH) 3 → Al 2 O 3 + 3 H 2 O
Fill out the table: Comparative characteristics of oxides and hydroxides
metals of the main subgroups I - III groups
Comparison Questions
I group
II group
Oxidation state of Me in the oxide.
2. Physical properties.
III group
3. Chemical properties (compare).
4. Methods for producing oxides.
Interaction:
a) with water
b) with acids
c) with acid oxides
d) with amphoteric oxides
d) with alkalis
5. Hydroxide formula.
Oxidation state of Me in hydroxide.
6. Physical properties
7. Chemical properties (compare).
8. Methods for producing hydroxides.
Interaction:
a) action on indicators
b) with acids
c) with acid oxides
d) with salt solutions
e) with non-metals
e) with alkalis
g) with amphoteric oxides and hydroxides
h) attitude to heating
Bases are formed by metal atoms and a hydroxyl group (OH -), which is why they are called hydroxides.
1. In relation to to the water the grounds are divided into:
2. By interaction with others Chemically, hydroxides are divided into:
A number of exceptions:
See chemical properties
THINGS
_________________________________
simple complex
____/______ ______________/___________
metals nonmetals oxides hydroxides salts
K, Ba S, P P 2 O 5 H 2 SO 4 Cu(NO 3) 2
Na 2 O Ba(OH) 2 Na 2 CO 3
Let's consider the classification, chemical properties and methods of obtaining complex substances.
OXIDES
OXIDE is a complex substance consisting of two elements, one of which is oxygen, which is in the -2 oxidation state.
The exceptions are:
1) compounds of oxygen and fluorine - fluorides: for example, oxygen fluoride OF 2 (oxidation state of oxygen in this compound +2)
2) peroxides (compounds of some elements with oxygen in which there is a bond between oxygen atoms), for example:
hydrogen peroxide H 2 O 2 potassium peroxide K 2 O 2
Examples of oxides: calcium oxide - CaO, barium oxide - BaO. If an element forms several oxides, then the valence of the element is indicated in their names in parentheses, for example: sulfur oxide (IV) - SO 2, sulfur oxide (VI) - SO 3.
All oxides can be divided into two large groups: salt-forming (salt-forming) and non-salt-forming.
Salt-forming substances are divided into three groups: basic, amphoteric and acidic.
O OXIDES
_________________/__________________
salt-forming non-salt-forming
CO, N2O, NO
↓ ↓ ↓
basic amphoteric acid
(they (they correspond to
correspond, acids)
grounds)
CaO, Li 2 O ZnO, BeO, PbO P 2 O 5, Mn 2 O 7
Cr 2 O 3, Al 2 O 3
Non-metals form acidic oxides, for example: nitrogen oxide (V) - N 2 O 5, carbon monoxide (IV) - CO 2. Metals with a valence of less than three, as a rule, form basic oxides, for example: sodium oxide - Na 2 O, magnesium oxide - MgO; and with a valence of more than four - acidic oxides, for example, manganese (VII) oxide - Mn 2 O 7, tungsten (VI) oxide - WO 3.
Let's consider the chemical properties of acidic and basic oxides.
CHEMICAL PROPERTIES OF OXIDES
BASIC ACID
Interaction with water
The product of the reaction is:
base acid
(if, in the composition of the oxide P 2 O 5 + 3H 2 O à 2H 3 PO 4
includes active metal, SiO 2 +H 2 O ≠
Li, Na, K, Rb, Cs, Fr, Ba, Ca)
CaO + H 2 O à Ca(OH) 2
2. Interaction with each other, forming salts CuO + SO 3 à CuSO 4
3. Interaction with hydroxides:
With soluble acids, with soluble bases
As a result of the reaction, salt and water are formed
CuO + H 2 SO 4 àCuSO 4 + H 2 O CO 2 +Ca(OH) 2 àCaCO 3 + H 2 O
Less volatile oxides
Replaces more volatile ones
from their salts:
K 2 CO 3 + SiO 2 à K 2 SiO 3 + CO 2
Amphoteric oxides include: metal oxides with a valence of three, for example: aluminum oxide - Al 2 O 3, chromium (III) oxide - Cr 2 O 3, iron (III) oxide - Fe 2 O 3, as well as a few exceptions , in which the metal is divalent, for example: beryllium oxide BeO, zinc oxide ZnO, lead (II) oxide - PbO. .
Amphoteric oxides have a dual nature: they are simultaneously capable of reactions in which they enter as basic and as acidic oxides
Let us prove the amphoteric nature of aluminum oxide. Let us present the equations for the reactions of interaction with hydrochloric acid and alkali (in an aqueous solution and when heated). When aluminum oxide and hydrochloric acid interact, a salt is formed - aluminum chloride. In this case, aluminum oxide acts as the main oxide.
Al 2 O 3 + 6HCl à2AlCl 3 + 3H 2 O
as main
In an aqueous solution, a complex salt is formed -
sodium tetrahydroxyaluminate:
Al 2 O 3 + 2NaOH + 3H 2 Oà 2Na sodium tetrahydroxoaluminate
like acidic
When fused with alkalis, meta-aluminates are formed.
Let us imagine the molecule of aluminum hydroxide Al(OH) 3 in the form of an acid, i.e. In the first place we write all the hydrogen atoms, in the second the acid residue:
H 3 AlO 3 - aluminum acid
For trivalent metals, subtract 1 H 2 O from the acid formula, obtaining meta-aluminum acid:
- H 2 O
HAlO 2 - meta-aluminum acid
fusion
Al 2 O 3 +2 NaOHà 2NaAlO 2 + H 2 O sodium metaaluminate
like acidic
METHODS FOR OBTAINING OXIDES:
1. Interaction of simple substances with oxygen:
4Al + 3O 2 à 2Al 2 O 3
2. Combustion or roasting of complex substances:
CH 4 + 2O 2 à CO 2 + 2H 2 O
2ZnS + 3O 2 à 2SO 2 + 2ZnO
3. Decomposition when heating insoluble hydroxides:
Cu(OH) 2 à CuO + H 2 O H 2 SiO 3 à SiO 2 + H 2 O
4. Decomposition when heating medium and acidic salts:
CaCO 3 à CaO + CO 2
2КHCO 3 àK 2 CO 3 + CO 2 +H 2 O
4AgNO 3 à4Ag + 4NO 2 + O 2
HYDROXIDES
Hydroxides are divided into three groups: bases, acids and amphoteric hydroxides (showing properties of both bases and acids).
BASE is a complex substance consisting of metal atoms and one or more hydroxyl groups
(- HE).
For example: sodium hydroxide - NaOH, barium hydroxide Ba(OH) 2. The number of hydroxyl groups in the base molecule is equal to the valency of the metal.
ACID is a complex substance that consists of hydrogen atoms that can be replaced by metal atoms and an acidic residue.
For example: sulfuric acid - H 2 SO 4, phosphoric acid - H 3 PO 4.
The valency of the acid residue is determined by the number of hydrogen atoms. In chemical compounds, the valence of the acid residue is retained (see Table 1).
TABLE 1 FORMULAS OF SOME ACIDS AND
ACID RESIDUE
| Acid name | Formula | Acid residue | Valency of acid residue | Name of the salt formed by this acid |
| Fluorescent | HF | F | I | fluoride |
| Solyanaya | HCl | Cl | I | chloride |
| Hydrobromic | HBr | Br | I | bromide |
| Hydroiodic | HI | I | I | iodide |
| Nitrogen | HNO3 | NO 3 | I | nitrate |
| Nitrogenous | HNO2 | NO 2 | I | nitrite |
| Vinegar | CH 3 COOH | CH 3 COO | I | acetate |
| Sulfuric | H2SO4 | SO 4 | II | sulfate |
| Sulphurous | H2SO3 | SO 3 | II | sulfite |
| Hydrogen sulfide | H2S | S | II | sulfide |
| Coal | H2CO3 | CO3 | II | carbonate |
| Flint | H2SiO3 | SiO3 | II | silicate |
| Phosphorus | H3PO4 | PO 4 | III | phosphate |
Based on their solubility in water, hydroxides are divided into two groups: soluble (for example, KOH, H 2 SO 4) and insoluble (H 2 SiO 3, Cu(OH) 2). Bases that dissolve in water are called alkalis.
