Bases (basic hydroxides). Chemical properties of hydroxides: bases, acids, amphoteric hydroxides How metal hydroxides interact

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.

Nomenclature and structure of the molecule

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 base molecules of a hydroxyl group associated with a 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.

Physical characteristics

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

Chemical properties of hydroxides

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.

Bases exhibiting dual properties

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.

Methods of preparation and use

In industries that require 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.

Before discussing the chemical properties of bases and amphoteric hydroxides, let's clearly define what they are?

1) Bases or basic hydroxides include metal hydroxides in the oxidation state +1 or +2, i.e. the formulas of which are written either as MeOH or Me(OH) 2. However, there are exceptions. Thus, the hydroxides Zn(OH) 2, Be(OH) 2, Pb(OH) 2, Sn(OH) 2 are not bases.

2) Amphoteric hydroxides include metal hydroxides in the oxidation state +3, +4, as well as, as exceptions, the hydroxides Zn(OH) 2, Be(OH) 2, Pb(OH) 2, Sn(OH) 2. Metal hydroxides in the oxidation state +4 are not found in Unified State Examination tasks, so they will not be considered.

Chemical properties of bases

All grounds are divided into:

Let us remember that beryllium and magnesium are not alkaline earth metals.

In addition to being soluble in water, alkalis also dissociate very well in aqueous solutions, while insoluble bases have a low degree of dissociation.

This difference in solubility and ability to dissociate between alkalis and insoluble hydroxides leads, in turn, to noticeable differences in their chemical properties. So, in particular, alkalis are more chemically active compounds and are often able to enter into reactions that insoluble bases do not.

Interaction of bases with acids

Alkalis react with absolutely all acids, even very weak and insoluble ones. For example:

Insoluble bases react with almost all soluble acids, but do not react with insoluble silicic acid:

It should be noted that both strong and weak bases with the general formula of the form Me(OH) 2 can form basic salts when there is a lack of acid, for example:

Interaction with acid oxides

Alkalis react with all acidic oxides, forming salts and often water:

Insoluble bases are capable of reacting with all higher acid oxides corresponding to stable acids, for example, P 2 O 5, SO 3, N 2 O 5, to form medium salts:

Insoluble bases of the type Me(OH) 2 react in the presence of water with carbon dioxide exclusively to form basic salts. For example:

Cu(OH) 2 + CO 2 = (CuOH) 2 CO 3 + H 2 O

Due to its exceptional inertness, only the strongest bases, alkalis, react with silicon dioxide. In this case, normal salts are formed. The reaction does not occur with insoluble bases. For example:

Interaction of bases with amphoteric oxides and hydroxides

All alkalis react with amphoteric oxides and hydroxides. If the reaction is carried out by fusing an amphoteric oxide or hydroxide with a solid alkali, this reaction leads to the formation of hydrogen-free salts:

If aqueous solutions of alkalis are used, then hydroxo complex salts are formed:

In the case of aluminum, under the action of an excess of concentrated alkali, instead of Na salt, Na 3 salt is formed:

Interaction of bases with salts

Any base reacts with any salt only if two conditions are met simultaneously:

1) solubility of the starting compounds;

2) the presence of precipitate or gas among the reaction products

For example:

Thermal stability of substrates

All alkalis, except Ca(OH) 2, are resistant to heat and melt without decomposition.

All insoluble bases, as well as slightly soluble Ca(OH) 2, decompose when heated. The highest decomposition temperature of calcium hydroxide is about 1000 o C:

Insoluble hydroxides have much lower decomposition temperatures. For example, copper (II) hydroxide decomposes already at temperatures above 70 o C:

Chemical properties of amphoteric hydroxides

Interaction of amphoteric hydroxides with acids

Amphoteric hydroxides react with strong acids:

Amphoteric metal hydroxides in the oxidation state +3, i.e. type Me(OH) 3, do not react with acids such as H 2 S, H 2 SO 3 and H 2 CO 3 due to the fact that the salts that could be formed as a result of such reactions are subject to irreversible hydrolysis to the original amphoteric hydroxide and corresponding acid:

Interaction of amphoteric hydroxides with acid oxides

Amphoteric hydroxides react with higher oxides, which correspond to stable acids (SO 3, P 2 O 5, N 2 O 5):

Amphoteric metal hydroxides in the oxidation state +3, i.e. type Me(OH) 3, do not react with acidic oxides SO 2 and CO 2.

Interaction of amphoteric hydroxides with bases

Of the bases, amphoteric hydroxides react only with alkalis. In this case, if an aqueous solution of alkali is used, then hydroxo complex salts are formed:

And when amphoteric hydroxides are fused with solid alkalis, their anhydrous analogues are obtained:

Interaction of amphoteric hydroxides with basic oxides

Amphoteric hydroxides react when fused with oxides of alkali and alkaline earth metals:

Thermal decomposition of amphoteric hydroxides

All amphoteric hydroxides are insoluble in water and, like any insoluble hydroxides, decompose when heated into the corresponding oxide and water.

Hydroxides can be thought of as the product of the addition (real or mental) of water to the corresponding oxides. Hydroxides are divided into bases, acids, and amphoteric hydroxides. Bases have the general composition M(OH)x, acids have the general composition HxCo. In molecules of oxygen-containing acids, the replaced hydrogen atoms are connected to the central element through oxygen atoms. In molecules of oxygen-free acids, hydrogen atoms are attached directly to a non-metal atom. Amphoteric hydroxides include primarily hydroxides of aluminum, beryllium and zinc, as well as hydroxides of many transition metals in intermediate oxidation states.
Based on solubility in water, soluble bases are distinguished - alkalis (formed by alkali and alkaline earth metals). The bases formed by other metals do not dissolve in water. Most inorganic acids are soluble in water. The only water-insoluble inorganic acids include silicic acid H2SiO3. Amphoteric hydroxides do not dissolve in water.

Chemical properties of bases.

All bases, both soluble and insoluble, have a common characteristic property - to form salts.
Let's consider the chemical properties of soluble bases (alkalis):
1. When dissolved in water, they dissociate to form a metal cation and a hydroxide anion. Change the color of the indicators: violet litmus - to blue, phenolphthalein - to crimson, methyl orange - to yellow, universal indicator paper - to blue.
2. Interaction with acid oxides:
alkali + acid oxide = salt.
3. Interaction with acids:
alkali + acid = salt + water.
The reaction between an acid and alkali is called a neutralization reaction.
4. Interaction with amphoteric hydroxides:
alkali + amphoteric hydroxide = salt (+ water)
5. Interaction with salts (subject to the solubility of the original salt and the formation of a precipitate or gas as a result of the reaction.
Let's consider the chemical properties of insoluble bases:
1. Interaction with acids:
base + acid = salt + water.
Polyacid bases are capable of forming not only intermediate, but also basic salts.
2. Heat decomposition:
base = metal oxide + water.

Chemical properties of acids.

All acids have a common characteristic property - the formation of salts when replacing hydrogen cations with metal/ammonium cations.
Let's consider the chemical properties of water-soluble acids:
1. When dissolved in water, they dissociate to form hydrogen cations and an acid residue anion. Change the color of the indicators to red (pink), with the exception of phenolphthalein (does not react to acids, remains colorless).
2. Interaction with metals in the activity series to the left of hydrogen (subject to the formation of a soluble salt):
acid + metal = salt + hydrogen.
When interacting with metals, the exceptions are oxidizing acids - nitric and concentrated sulfuric acids. Firstly, they also react with some metals that are to the right of hydrogen in the activity series. Secondly, the reaction with metals never releases hydrogen, but produces a salt of the corresponding acid, water and the reduction products of nitrogen or sulfur, respectively.
3. Interaction with bases/amphoteric hydroxides:
acid + base = salt + water.
4. Interaction with ammonia:
acid + ammonia = ammonium salt
5. Interaction with salts (subject to the formation of gas or sediment):
acid + salt = salt + acid.
Polybasic acids are capable of forming not only intermediate, but also acidic salts.
Insoluble silicic acid does not change the color of indicators (a very weak acid), but is capable of reacting with alkali solutions with slight heating:
1. Interaction of silicic acid with alkali solution:
silicic acid + alkali = salt + water.
2. Decomposition (during long-term storage or heating)
silicic acid = silicon(IV) oxide + water.

Chemical properties of amphoteric hydroxides.

Amphoteric hydroxides are capable of forming two series of salts, since when reacting with alkalis they exhibit the properties of an acid, and when reacting with acids they exhibit the properties of a base.
Let's consider the chemical properties of amphoteric hydroxides:
1. Interaction with alkalis:
amphoteric hydroxide + alkali = salt (+ water).
2. Interaction with acids:
amphoteric hydroxide + acid = salt + water.

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

base acid salt

Cu(OH) 2 + H 2 SO 4 = CuSO 4 + 2H 2 O,

base acid salt

2NaOH + PbO = Na 2 PbO 2 + H 2 O,

base amphoteric salt

2NaOH + Pb(OH) 2 = Na 2 PbO 2 + 2H 2 O,

base amphoteric salt

hydroxide

2H 3 PO 4 + 3Na 2 O = 2Na 3 PO 4 + 3H 2 O,

acid basic salt

H 2 SO 4 + SnO = SnSO 4 + H 2 O,

acid amphoteric salt

H 2 SO 4 + Sn(OH) 2 = SnSO 4 + 2H 2 O.

acid amphoteric salt

hydroxide

Amphoteric hydroxides exhibit the following basic properties in reactions with acids:

2Al(OH) 3 + 3H 2 SO 4 = Al 2 (SO 4) 3 + 6H 2 O,

with alkalis (bases) – acidic properties:

H 3 AlO 3 + 3NaOH = Na 3 AlO 3 + 3H 2 O,

or H 3 AlO 3 + NaOH = NaAlO 2 + 2H 2 O.

Bases and acids react with salts to form a precipitate or weak electrolyte. Weak acids - H 3 PO 4, H 2 CO 3, H 2 SO 3, H 2 SiO 3 and others.

2NaOH + NiSO 4 = Ni(OH) 2  + Na 2 SO 4,

base salt

3H 2 SO 4 + 2Na 3 PO 4 = 2H 3 PO 4 + 3Na 2 SO 4

acid salt

Oxygen-free acids undergo the same reactions as the previously discussed oxygen-containing acids.

Example. Make up the formulas of hydroxides corresponding to the oxides: a) FeO; b) N 2 O 3; c) Cr 2 O 3. Name the connections.

Solution

a) FeO is a basic oxide, therefore, the corresponding hydroxide is a base; in the base formula, the number of hydroxyl groups (OH) is equal to the oxidation state of the metal atom; the formula of iron (II) hydroxide is Fe(OH) 2.

b) N 2 O 3 is an acidic oxide, therefore the corresponding hydroxide is an acid. The acid formula can be obtained based on the representation of the acid as a hydrate of the corresponding oxide:

N2O3. H 2 O = (H 2 N 2 O 4) = 2HNO 2 – nitrous acid.

c) Cr 2 O 3 is an amphoteric oxide, therefore, the corresponding hydroxide is amphoteric. Amphoteric hydroxides are written in the form of bases - Cr(OH) 3 - chromium (III) hydroxide.

Salts

Salts- substances that consist of basic and acidic residues. Thus, the salt CuSO 4 consists of a main residue - the metal cation Cu 2+ and an acid residue - SO 4 2 .

According to traditional nomenclature, the names of salts of oxygen acids are composed as follows: the ending - is added to the root of the Latin name of the central atom of the acidic residue - at(at higher oxidation states of the central atom) or – it(for a lower oxidation state) and then - the remainder of the base in the genitive case, for example: Na 3 PO 4 - sodium phosphate, BaSO 4 - barium sulfate, BaSO 3 - barium sulfite. The names of salts of oxygen-free acids are formed by adding the suffix - to the root of the Latin name of the non-metal. eid and the Russian name of the metal (residue from the base), for example CaS - calcium sulfide.

Medium salts do not contain in its composition, hydrogen ions and hydroxo groups that can be replaced by metal, for example CuCl 2, Na 2 CO 3 and others.

Chemical properties of salts

Medium salts enter into exchange reactions with alkalis, acids, and salts. For examples of appropriate reactions, see above.

Acid salts contain the acid residue contains a hydrogen ion, for example NaHCO 3, CaHPO 4, NaH 2 PO 4, etc. In the name of an acid salt, the hydrogen ion is denoted by the prefix hydro-, before which the number of hydrogen atoms in the salt molecule is indicated if it is greater than one. For example, the names of the salts of the above composition are, respectively, sodium bicarbonate, calcium hydrogen phosphate, sodium dihydrogen phosphate.

Acid salts are obtained

interaction between the base and polybasic acid with excess acid:

Ca(OH) 2 + H 3 PO 4 = CaHPO 4 + 2H 2 O;

the interaction of the average salt of a polybasic acid and the corresponding acid or a stronger acid taken in deficiency:

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

Na 3 PO 4 + HCl = Na 2 HPO 4 + NaCl.

Basic salts contain the base residue contains a hydroxo group, for example CuOHNO 3, Fe(OH) 2 Cl. In the name of the main salt, the hydroxo group is designated by the prefix hydroxo-, for example, the names of the above salts are respectively: copper (II) hydroxonitrate, iron (III) dihydroxychloride.

Basic salts are obtained

the interaction of a polyacid (containing more than one hydroxo group) base and acid with an excess of base:

Cu(OH) 2 + HNO 3 = CuOHNO 3 + H 2 O;

the interaction of a salt formed by a polyacid base and a base taken in deficiency:

FeCl 3 + NaOH = FeOHCl 2  + NaCl,

FeCl 3 + 2NaOH = Fe(OH) 2 Cl + 2NaCl.

Acidic and basic salts have all the properties of salts. In reactions with alkalis, acidic salts, and with acids, basic salts turn into intermediate salts.

Na 2 HPO 4 + NaOH = Na 3 PO 4 + H 2 O,

Na 2 HPO 4 + 2HCl = H 3 PO 4 + 2NaCl,

FeOHCl 2 + HCl = FeCl 3 + H 2 O,

FeOHCl 2 + 2NaOH = Fe(OH) 3  + 2NaCl.

Example 1. Make up the formulas of all salts that can be formed by the base Mg(OH) 2 and the acid H 2 SO 4.

Solution

We compose salt formulas from possible basic and acidic residues, observing the rule of electrical neutrality. Possible basic residues are Mg 2+ and MgOH +, acidic residues are SO 4 2- and HSO 4 . The charges of complex basic and acidic residues are equal to the sum of the oxidation states of their constituent atoms. Using a combination of basic and acidic residues, we compose the formulas of possible salts: MgSO 4 - average salt - magnesium sulfate; Mg(HSO 4) 2 – acid salt – magnesium hydrogen sulfate; (MgOH) 2 SO 4 – the main salt is magnesium hydroxysulfate.

Example 2. Write the reactions of salt formation during the interaction of oxides

a) PbO and N 2 O 5; b) PbO and Na 2 O.

Solution

In reactions between oxides, salts are formed, the basic residues of which are formed from basic oxides, the acid residues from acidic oxides.

a) In the reaction with the acidic oxide N 2 O 5, the amphoteric oxide PbO exhibits the properties of a basic oxide, therefore, the main residue of the resulting salt is Pb 2+ (the charge of the lead cation is equal to the oxidation state of lead in the oxide), the acid residue is NO 3 - (acid residue corresponding to a given acidic nitric oxide). Reaction equation

PbO + N 2 O 5 = Pb(NO 3) 2.

b) In the reaction with the basic oxide Na 2 O, the amphoteric oxide PbO exhibits the properties of an acidic oxide; the acidic residue of the resulting salt (PbO 2 2 ) is found from the acid form of the corresponding amphoteric hydroxide Pb(OH) 2 = H 2 PbO 2. Reaction equation

Main classes of inorganic compounds

*( Dear students! To study this topic and complete test tasks, as visual material, you must have a table of the Periodic Table of Elements, a table of the solubility of compounds and a series of metal stresses.

All substances are divided into simple, consisting of atoms of one element, and complex, consisting of atoms of two or more elements. Complex substances are usually divided into organic, which includes almost all carbon compounds (except for the simplest ones, such as CO, CO 2, H 2 CO 3, HCN) and inorganic. The most important classes of inorganic compounds include:

a) oxides - binary compounds of an element with oxygen;

b) hydroxides, which are divided into basic (bases), acidic (acids) and amphoteric;

Before proceeding with the characterization of classes of inorganic compounds, it is necessary to consider the concepts of valency and oxidation state.

Valence and oxidation state

Valence characterizes the ability of an atom to form chemical bonds. Quantitatively valence is the number of bonds that an atom of a given element forms in a molecule. In accordance with modern ideas about the structure of atoms and chemical bonds, atoms of elements are capable of donating, gaining electrons and forming common electron pairs. Assuming that each chemical bond is formed by a pair of electrons, valence can be defined as the number of electron pairs by which an atom is bonded to other atoms. Valence has no sign.

Oxidation state (CO) - This conventional charge of an atom in a molecule, calculated from the assumption that the molecule consists of ions.

Ions- These are positively and negatively charged particles of matter. Positively charged ions are called cations, negative - anions. Ions can be simple, for example Cl-(consist of one atom) or complex, for example SO 4 2-(consist of several atoms).

If the molecules of substances consist of ions, then we can conditionally assume that there is a purely electrostatic connection between the atoms in the molecule. This means that regardless of the nature of the chemical bond in the molecule, the atoms of the more electronegative element attract electrons from the less electronegative atom.

Oxidation state usually indicated by Roman numerals with a “+” or “-” sign before the number (e.g., +III), and the charge of an ion is indicated by an Arabic numeral with a “+” or “-” sign behind the number (e.g., 2-).

Rules for determining the oxidation state of an element in a compound:

1. The CO of an atom in a simple substance is zero, for example, O 2 0, C 0, Na 0.

2. CO of fluorine is always equal to -I, because it is the most electronegative element.

3. Hydrogen CO is equal to +I in compounds with non-metals (H 2 S, NH 3) and -I in compounds with active metals (LiH, CaH 2).

4. CO of oxygen in all compounds is equal to -II (except for hydrogen peroxide H 2 O 2 and its derivatives, where the oxidation state of oxygen is -I, and ОF 2, where oxygen exhibits CO +II).

5. Metal atoms always have a positive oxidation state equal to or less than their group number in the Periodic Table. For the first three groups, the CO of metals coincides with the group number, with the exception of copper and gold, for which the more stable oxidation states are +II and +III, respectively.

6. The highest (maximum) positive CO of an element is equal to the number of the group in which it is located (for example, P is in the V group A subgroup and has CO +V). This rule applies to elements of both main and secondary subgroups. An exception is for elements I B and VIII A and B subgroups, as well as for fluorine and oxygen.

7. Negative (minimal) CO is characteristic only for elements of the main subgroups IV A - VII A, and it is equal to the group number minus 8.

8. The sum of CO of all atoms in a molecule is zero, and in a complex ion it is equal to the charge of this ion.

Example: Calculate the oxidation state of chromium in the compound K 2 Cr 2 O 7 .

Solution: Let us denote the CO of chromium as X. Knowing the CO of oxygen, equal to -II, and the CO of potassium +I (by the number of the group in which potassium is located), we create the equation:

K 2 +I Cr 2 X O 7 -II

1 2 + X·2 + (-2)·7 = 0

Having solved the equation, we get x = 6. Therefore, the CO of the chromium atom is equal to +VI.

Oxides

Oxides are compounds of elements with oxygen. The oxidation state of oxygen in oxides is II.

Compilation of oxide formulas

The formula of any oxide will be E 2 O x, where X- the degree of oxidation of the element forming the oxide (even indices should be reduced by two, for example, they write not S 2 O 6, but SO 3). To compile the oxide formula, you need to know in which group of the Periodic Table the element is located. The maximum CO of an element is equal to the group number. In accordance with this, the formula of the higher oxide of any element, depending on the group number, will look like:

Exercise: Make up formulas for higher oxides of manganese and phosphorus.

Solution: Manganese is located in the VII B subgroup of the Periodic Table, which means its highest CO is +VII. The formula of the higher oxide will be Mn 2 O 7.

Phosphorus is located in the V A subgroup, hence the formula of its higher oxide is P 2 O 5.

If the element is not in the highest oxidation state, it is necessary to know this oxidation state. For example, sulfur, being in the VI A subgroup, may have an oxide in which it exhibits a CO equal to +IV. The formula for sulfur oxide (+IV) will be SO 2.

Nomenclature of oxides

According to the International Nomenclature (IUPAC), the name of oxides is formed from the word “oxide” and the name of the element in the genitive case.

For example: CaO - oxide of (what?) calcium

H 2 O - hydrogen oxide

SiO 2 - silicon oxide

The CO of the element forming the oxide may not be indicated if it exhibits only one CO, for example:

Al 2 O 3 - aluminum oxide;

MgO - magnesium oxide

If an element has several oxidation states, they must be indicated:

CuO - copper (II) oxide, Cu 2 O - copper (I) oxide

N 2 O 3 - nitric oxide (III), NO - nitric oxide (II)

The old names of oxides, indicating the number of oxygen atoms in the oxide, have been preserved and are often used. In this case, Greek numerals are used - mono-, di-, tri-, tetra-, penta-, hexa-, etc.

For example:

SO 2 - sulfur dioxide, SO 3 - sulfur trioxide

NO - nitrogen monoxide

In the technical literature, as well as in industry, trivial or technical names of oxides are widely used, for example:

CaO - quicklime, Al 2 O 3 - alumina

CO 2 - carbon dioxide, CO - carbon monoxide

SiO 2 - silica, SO 2 - sulfur dioxide

Methods for obtaining oxides

a) Direct interaction of the element with oxygen under proper conditions:

Al + O 2 → Al 2 O 3 ;(~ 700 °C)

Cu + O 2 → CuO(< 200 °С)

S + O 2 → SO 2

This method cannot produce oxides of inert gases, halogens, and “noble” metals.

b) Thermal decomposition of bases (except for alkali and alkaline earth metal bases):

Cu(OH) 2 → CuO + H 2 O(> 200 °C)

Fe(OH) 3 → Fe 2 O 3 + H 2 O(~ 500-700 °C)

c) Thermal decomposition of some acids:

H 2 SiO 3 → SiO 2 + H 2 O (1000°)

H 2 CO 3 → CO 2 + H 2 O (boiling)

d) Thermal decomposition of salts:

CaCO 3 → CaO + CO 2 (900° C)

FeCO 3 → FeO + CO 2 (490°)

Oxides classification

Based on their chemical properties, oxides are divided into salt-forming and non-salt-forming.

Non-salt-forming(indifferent) oxides form neither acids nor bases (do not react with acids, bases, or water). These include: carbon monoxide (II) - CO, nitrogen oxide (I) - N 2 O, nitrogen oxide (II) - NO and some others.

Salt-forming oxides are divided into basic, acidic and amphoteric.

Main are those oxides that correspond to hydroxides, called reasons. These are oxides of most metals in the lowest oxidation state (Li 2 O, Na 2 O, MgO, CaO, Ag 2 O, Cu 2 O, CdO, FeO, NiO, V 2 O 3, etc.).

By adding (directly or indirectly) water, basic oxides form basic hydroxides (bases). For example, copper (II) oxide - CuO corresponds to copper (II) hydroxide - Cu(OH) 2, and BaO oxide - barium hydroxide - Ba(OH) 2.

It is important to remember that the CO of the element in the oxide and its corresponding hydroxide is the same!

Basic oxides react with acids or acidic oxides to form salts.

Acidic are those oxides that correspond to acidic hydroxides, called acids. Acidic oxides form nonmetals and some metals in higher oxidation states (N 2 O 5, SO 3, SiO 2, CrO 3, Mn 2 O 7, etc.).

By adding water (directly or indirectly), acid oxides form acids. For example, nitrogen oxide (III) - N 2 O 3 corresponds to nitrous acid HNO 2, chromium oxide (VI) - CrO 3 - chromic acid H 2 CrO 4.

Acidic oxides react with bases or basic oxides to form salts.

Acidic oxides can be considered as products of the “removal” of water from acids and called anhydrides (i.e. anhydrous). For example, SO 3 is sulfuric acid anhydride H 2 SO 4 (or simply sulfuric anhydride), P 2 O 5 is orthophosphoric anhydride H 3 PO 4 (or simply phosphoric anhydride).

It is important to remember that the CO of an element in the oxide and its corresponding acid, as well as in the anion of this acid, is the same!

Amphoteric are those oxides that can correspond to both acids and bases. These include BeO, ZnO, Al 2 O 3, SnO, SnO 2, Cr 2 O 3 and oxides of some other metals in intermediate oxidation states. The acidic and basic properties of these oxides are expressed to varying degrees. For example, in aluminum and zinc oxides, the acidic and basic properties are expressed approximately equally, in Fe 2 O 3 the basic properties predominate, and in PbO 2 the acidic properties predominate.

Amphoteric oxides form salts when reacting with both acids and bases.

Chemical properties of oxides

The chemical properties of oxides (and their corresponding hydroxides) follow the principle of acid-base interaction, according to which compounds exhibiting acidic properties react with compounds having basic properties.

Basic oxides interact:

a) with acids:

CuO + H 2 SO 4 → H 2 O + CuSO 4 ;

BaO + H 3 PO 4 → H 2 O + Ba 3 (PO 4) 2;

b) with acid oxides:

CuO + SO 2 → CuSO 3;

BaO + N 2 O 5 → Ba(NO 3) 2;

c) oxides of alkali and alkaline earth metals can be dissolved in water:

Na 2 O + H 2 O → NaOH;

BaO + H 2 O → Ba(OH) 2.

Acidic oxides interact:

a) with reasons:

N 2 O 3 + NaOH → H 2 O + NaNO 2;

CO 2 + Fe(OH) 2 → H 2 O + FeCO 3 ;

b) with basic oxides:

SO 2 + CaO → CaSO 3;

SiO 2 + Na 2 O → Na 2 SiO 3;

c) can (but not all) dissolve in water:

SO 3 + H 2 O → H 2 SO 4;

P 2 O 3 + H 2 O → H 3 PO 3 .

Amphoteric oxides can interact:

a) with acids:

ZnO + H 2 SO 4 → H 2 O + ZnSO 4 ;

Al 2 O 3 + H 2 SO 4 → H 2 O + Al 2 (SO 4) 3;

b) with acid oxides:

ZnO + SO 3 → ZnSO 4;

Al 2 O 3 + SO 3 → Al 2 (SO 4) 3;

c) with reasons:

ZnO + NaOH + H 2 O → Na 2;

Al 2 O 3 + NaOH + H 2 O → Na 3;

d) with basic oxides:

ZnO + Na 2 O → Na 2 ZnO 2 ;

Al 2 O 3 + Na 2 O → NaAlO 2.

In the first two cases, amphoteric oxides exhibit the properties of basic oxides, in the last two cases - the properties of acidic oxides.

Hydroxides

Hydroxides are oxide hydrates with the general formula m E 2 O X· n H2O( n And m- small integers, X- valency of the element). Hydroxides differ from oxides in composition only by the presence of water in their molecule. According to their chemical properties, hydroxides are divided into basic(bases), acidic(acids) and amphoteric.

Bases (basic hydroxides)

The basis called a compound of an element with one, two, three and less often four hydroxyl groups with the general formula E(OH) X. The elements are always metals of the main or secondary subgroups.

Soluble bases- these are electrolytes that dissociate in an aqueous solution (break up into ions) to form anions of the hydroxyl group OH ‾ and a metal cation. For example:

KOH = K + + OH ‾ ;

Ba(OH) 2 = Ba 2+ + 2OH ‾

Due to the presence of hydroxyl ions OH ‾ in an aqueous solution, bases exhibit an alkaline reaction of the medium.

Drawing up a base formula

To compose the base formula, you need to write the symbol of the metal and, knowing its oxidation state, assign the corresponding number of hydroxyl groups next to it. For example: the Mg +II ion corresponds to the base Mg(OH) 2, the Fe +III ion corresponds to the base Fe(OH) 3, etc. For the first three groups of the main subgroups of the Periodic Table, the oxidation state of metals is equal to the group number, so the base formula will be EOH (for metals of the I A subgroup), E(OH) 2 (for metals of the II A subgroup), E(OH) 3 (for metals of the III A subgroups). For other groups (mainly side subgroups), it is necessary to know the oxidation state of the element, because it may not match the group number.

Nomenclature of bases

The names of the bases are formed from the word “hydroxide” and the name of the element in the genitive case, followed by Roman numerals in parentheses indicating the oxidation state of the element, if necessary. For example: KOH - potassium hydroxide, Fe(OH) 2 - iron (II) hydroxide, Fe(OH) 3 - iron (III) hydroxide, etc.

There are technical names for some bases: NaOH - sodium hydroxide, KOH - potassium hydroxide, Ca(OH) 2 - slaked lime.

Methods for obtaining bases

a) Dissolution of basic oxides in water (only oxides of alkali and alkaline earth metals are soluble in water):

Na 2 O + H 2 O → NaOH;

CaO + H 2 O → Ca(OH) 2;

b) Interaction of alkali and alkaline earth metals with water:

Na + H 2 O → H 2 + NaOH;

Ca + H 2 O → H 2 + Ca(OH) 2;

c) Displacement of a weak base from a salt by a strong base:

NaOH + CuSO 4 → Cu(OH) 2 ↓ + Na 2 SO 4;

Ba(OH) 2 + FeCl 3 → Fe(OH) 3 ↓ + BaCl 2.

Classification of bases

a) Based on the number of hydroxyl groups, bases are divided into single- and polyacid: EON, E(OH) 2, E(OH) 3, E(OH) 4. Index X in the base formula, E(OH) x is called the “acidity” of the base.

b) Reasons may be soluble And insoluble in water. Most bases are insoluble in water. Bases that are highly soluble in water form elements of the I A subgroup - Li, Na, K, Rb, Cs, Fr (alkali metals). They're called alkalis. In addition, ammonia hydrate NH 3 ·H 2 O, or ammonium hydroxide NH 4 OH, is a soluble base, but it is not an alkali. The hydroxides of Ca, Sr, Ba (alkaline earth metals) have less solubility, and their solubility increases in the group from top to bottom: Ba(OH) 2 is the most soluble base.

c) Based on their ability to dissociate into ions in solution, bases are divided into strong And weak. Strong bases are hydroxides of alkali and alkaline earth metals - they completely dissociate into ions. The remaining bases are bases of medium strength or weak. Ammonia hydrate is also a weak base.

Chemical properties of bases

Grounds interact with compounds exhibiting acidic properties:

a) React with acids to form salt and water. This reaction is called reaction neutralization:

Ca(OH) 2 + H 2 SO 4 → CaSO 4 + H 2 O;

b) Interact with acidic or amphoteric oxides (these reactions can also be classified as neutralization reactions or acid-base interactions):

Cu(OH) 2 + SO 2 → H 2 O + CuSO 4 ;

NaOH + ZnO → Na 2 ZnO 2 + H 2 O;

c) Interact with acidic salts (acid salts contain a hydrogen atom in the acid anion);

Ca(OH) 2 + Ca(HCO 3) 2 → CaCO 3 + H 2 O;

NaOH + Ca(HSO 4) 2 → CaSO 4 + Na 2 SO 4 + H 2 O;

d) Strong bases can displace weak ones from salts:

NaOH + MnCl 2 → Mn(OH) 2 ↓ + NaCl;

Ba(OH) 2 + Mg(NO 3) 2 → Mg(OH) 2 ↓ + Ba(NO 3) 2;

e) water-insoluble bases decompose into oxide and water when heated.