Amphoteric metals are simple substances that are structurally, chemically and similar to the metal group of elements. Metals themselves cannot exhibit amphoteric properties, unlike their compounds. For example, the oxides and hydroxides of some metals have a dual chemical nature - in some conditions they behave like acids, while in others they have the properties of alkalis.

The main amphoteric metals are aluminum, zinc, chromium, and iron. Beryllium and strontium can be attributed to the same group of elements.

amphoteric?

For the first time this property was discovered quite a long time ago. And the term "amphoteric elements" was introduced into science in 1814 by the famous chemists L. Tenard and J. Gay-Lussac. In those days, it was customary to divide chemical compounds into groups that corresponded to their basic properties during reactions.

However, the group of oxides and bases had dual abilities. Under some conditions, such substances behaved like alkalis, while in others, on the contrary, they acted like acids. This is how the term "amphoteric" was born. For such, the behavior during the acid-base reaction depends on the reaction conditions, the nature of the reactants involved, and the properties of the solvent.

Interestingly, under natural conditions, amphoteric metals can interact with both alkali and acid. For example, during the reaction of aluminum with aluminum sulfate is formed. And when the same metal reacts with concentrated alkali, a complex salt is formed.

Amphoteric bases and their main properties

Under normal conditions, these are solids. They are practically insoluble in water and are considered rather weak electrolytes.

The main method for obtaining such bases is the reaction of a metal salt with a small amount of alkali. The precipitation reaction must be carried out slowly and carefully. For example, when receiving zinc hydroxide, caustic soda is carefully added in drops to a test tube with zinc chloride. Each time you need to gently shake the container to see the white precipitate of metal at the bottom of the dish.

With acids and amphoteric substances react as bases. For example, the reaction of zinc hydroxide with hydrochloric acid produces zinc chloride.

But during reactions with bases, amphoteric bases behave like acids.

In addition, when strongly heated, they decompose to form the corresponding amphoteric oxide and water.

The most common amphoteric metals: a brief description

Zinc belongs to the group of amphoteric elements. And although alloys of this substance were widely used in ancient civilizations, it was only in 1746 that they could isolate it in its pure form.

Pure metal is a rather brittle bluish substance. Zinc rapidly oxidizes in air - its surface tarnishes and becomes covered with a thin film of oxide.

In nature, zinc exists mainly in the form of minerals - zincites, smithsonites, calamites. The most famous substance is zinc blende, which consists of zinc sulfide. The largest deposits of this mineral are in Bolivia and Australia.

Aluminum Today it is considered the most common metal on the planet. Its alloys have been used for many centuries, and in 1825 the substance was isolated in its pure form.

Pure aluminum is a light, silver-colored metal. It is easy to machine and cast. This element has high electrical and thermal conductivity. In addition, this metal is resistant to corrosion. The fact is that its surface is covered with a thin, but very resistant oxide film.

Today, aluminum is widely used in industry.

Bases, amphoteric hydroxides

Bases are complex substances consisting of metal atoms and one or more hydroxo groups (-OH). The general formula is Me + y (OH) y, where y is the number of hydroxo groups equal to the oxidation state of the metal Me. The table shows the classification of bases.

Properties of alkali hydroxides of alkali and alkaline earth metals

1. Aqueous solutions of alkalis are soapy to the touch, change the color of indicators: litmus - blue, phenolphthalein - raspberry.

2. Aqueous solutions dissociate:

3. Interact with acids, entering into an exchange reaction:

Polyacid bases can give intermediate and basic salts:

4. Interact with acid oxides, forming medium and acid salts, depending on the basicity of the acid corresponding to this oxide:

5. Interact with amphoteric oxides and hydroxides:

a) fusion:

b) in solutions:

6. React with water-soluble salts if a precipitate or gas is formed:

Insoluble bases (Cr (OH) 2, Mn (OH) 2, etc.) interact with acids and decompose when heated:

Amphoteric hydroxides

Compounds are called amphoteric, which, depending on the conditions, can be both donors of hydrogen cations and exhibit acidic properties, and their acceptors, i.e., exhibit basic properties.

Chemical properties of amphoteric compounds

1. Interacting with strong acids, they reveal the main properties:

Zn(OH) 2 + 2HCl = ZnCl 2 + 2H 2 O

2. Interacting with alkalis - strong bases, they exhibit acidic properties:

Zn (OH) 2 + 2NaOH \u003d Na 2 ( complex salt)

Al (OH) 3 + NaOH \u003d Na ( complex salt)

Compounds are called complex in which at least one covalent bond was formed by the donor-acceptor mechanism.

The general method for obtaining bases is based on exchange reactions, by which both insoluble and soluble bases can be obtained.

CuSO 4 + 2KOH \u003d Cu (OH) 2 ↓ + K 2 SO 4

K 2 CO 3 + Ba (OH) 2 \u003d 2 KOH + BaCO 3 ↓

When soluble bases are obtained by this method, an insoluble salt precipitates.

When obtaining water-insoluble bases with amphoteric properties, an excess of alkali should be avoided, since dissolution of the amphoteric base may occur, for example:

AlCl 3 + 4KOH \u003d K [Al (OH) 4] + 3KSl

In such cases, ammonium hydroxide is used to obtain hydroxides, in which amphoteric hydroxides do not dissolve:

AlCl 3 + 3NH 3 + ZH 2 O \u003d Al (OH) 3 ↓ + 3NH 4 Cl

Hydroxides of silver and mercury decompose so easily that when you try to obtain them by an exchange reaction, instead of hydroxides, oxides precipitate:

2AgNO 3 + 2KOH \u003d Ag 2 O ↓ + H 2 O + 2KNO 3

In industry, alkalis are usually obtained by electrolysis of aqueous solutions of chlorides.

2NaCl + 2H 2 O → ϟ → 2NaOH + H 2 + Cl 2

Alkalis can also be obtained by reacting alkali and alkaline earth metals or their oxides with water.

2Li + 2H 2 O \u003d 2LiOH + H 2

SrO + H 2 O \u003d Sr (OH) 2

acids

Acids are called complex substances, the molecules of which consist of hydrogen atoms that can be replaced by metal atoms, and acid residues. Under normal conditions, acids can be solid (phosphoric H 3 PO 4; silicon H 2 SiO 3) and liquid (sulfuric acid H 2 SO 4 will be a pure liquid).

Gases such as hydrogen chloride HCl, hydrogen bromide HBr, hydrogen sulfide H 2 S form the corresponding acids in aqueous solutions. The number of hydrogen ions formed by each acid molecule during dissociation determines the charge of the acid residue (anion) and the basicity of the acid.

According to protolytic theory of acids and bases, proposed simultaneously by the Danish chemist Bronsted and the English chemist Lowry, an acid is a substance splitting off with this reaction protons, A basis- a substance capable of receive protons.

acid → base + H +

Based on these ideas, it is clear basic properties of ammonia, which, due to the presence of a lone electron pair at the nitrogen atom, effectively accepts a proton when interacting with acids, forming an ammonium ion through a donor-acceptor bond.

HNO 3 + NH 3 ⇆ NH 4 + + NO 3 -

acid base acid base

A more general definition of acids and bases proposed by the American chemist G. Lewis. He suggested that acid-base interactions are quite do not necessarily occur with protone transfer. In the determination of acids and bases according to Lewis, the main role in chemical reactions is given to electronic steam.

Cations, anions, or neutral molecules that can accept one or more pairs of electrons are called Lewis acids.

For example, aluminum fluoride AlF 3 is an acid, since it is able to accept an electron pair when interacting with ammonia.

AlF 3 + :NH 3 ⇆ :

Cations, anions or neutral molecules capable of donating electron pairs are called Lewis bases (ammonia is a base).

The Lewis definition covers all acid-base processes that have been considered by the previously proposed theories. The table compares the definitions of acids and bases currently in use.

Nomenclature of acids

Since there are different definitions of acids, their classification and nomenclature are rather arbitrary.

According to the number of hydrogen atoms capable of splitting off in an aqueous solution, acids are divided into monobasic(e.g. HF, HNO 2), dibasic(H 2 CO 3 , H 2 SO 4) and tribasic(H 3 RO 4).

According to the composition of the acid is divided into anoxic(HCl, H 2 S) and oxygen-containing(HClO 4 , HNO 3).

Usually names of oxygenated acids derived from the name of a non-metal with the addition of the endings -kai, -way, if the oxidation state of the non-metal is equal to the group number. As the oxidation state decreases, the suffixes change (in order of decreasing metal oxidation state): - oval, ististaya, - ovate:

If we consider the polarity of the hydrogen-non-metal bond within a period, we can easily relate the polarity of this bond to the position of the element in the Periodic Table. From metal atoms that easily lose valence electrons, hydrogen atoms accept these electrons, forming a stable two-electron shell like the shell of a helium atom, and give ionic metal hydrides.

In hydrogen compounds of elements of groups III-IV of the Periodic system, boron, aluminum, carbon, silicon form covalent, weakly polar bonds with hydrogen atoms that are not prone to dissociation. For elements of groups V-VII of the Periodic system, within a period, the polarity of the non-metal-hydrogen bond increases with the charge of the atom, but the distribution of charges in the resulting dipole is different than in hydrogen compounds of elements that tend to donate electrons. Atoms of non-metals, in which several electrons are needed to complete the electron shell, pull towards themselves (polarize) a pair of bond electrons the stronger, the greater the charge of the nucleus. Therefore, in the series CH 4 - NH 3 - H 2 O - HF or SiH 4 - PH 3 - H 2 S - Hcl, bonds with hydrogen atoms, while remaining covalent, become more polar, and the hydrogen atom in the dipole of the element-hydrogen bond becomes more electropositive. If polar molecules are in a polar solvent, the process of electrolytic dissociation can occur.

Let us discuss the behavior of oxygen-containing acids in aqueous solutions. These acids have an H-O-E bond and, naturally, the O-E bond affects the polarity of the H-O bond. Therefore, these acids dissociate, as a rule, more easily than water.

H 2 SO 3 + H 2 O ⇆ H s O + + HSO 3

HNO 3 + H 2 O ⇆ H s O + + NO 3

Let's look at a few examples properties of oxygenated acids, formed by elements that are capable of exhibiting different oxidation states. It is known that hypochlorous acid HClO very weak hydrochloric acid HClO 2 also weak but stronger than hypochlorous, hypochlorous acid HclO 3 strong. Perchloric acid HClO 4 is one of the the strongest inorganic acids.

Dissociation according to the acidic type (with the elimination of the H ion) requires breaking the O-H bond. How can one explain the decrease in the strength of this bond in the series HClO - HClO 2 - HClO 3 - HClO 4? In this series, the number of oxygen atoms associated with the central chlorine atom increases. Each time a new bond of oxygen with chlorine is formed, an electron density is drawn away from the chlorine atom, and hence from the single O-Cl bond. As a result, the electron density partially leaves the О-Н bond, which is weakened because of this.

Such a pattern - enhancement of acidic properties with an increase in the degree of oxidation of the central atom - characteristic not only for chlorine, but also for other elements. For example, nitric acid HNO 3 , in which the nitrogen oxidation state is +5, is stronger than nitrous acid HNO 2 (nitrogen oxidation state is +3); sulfuric acid H 2 SO 4 (S +6) is stronger than sulfurous acid H 2 SO 3 (S +4).

Obtaining acids

1. Anoxic acids can be obtained in the direct combination of non-metals with hydrogen.

H 2 + Cl 2 → 2HCl,

H 2 + S ⇆ H 2 S

2. Some oxygenated acids can be obtained interaction of acid oxides with water.

3. Both anoxic and oxygenated acids can be obtained according to exchange reactions between salts and other acids.

BaBr 2 + H 2 SO 4 \u003d BaSO 4 ↓ + 2HBr

CuSO 4 + H 2 S \u003d H 2 SO 4 + CuS ↓

FeS + H 2 SO 4 (pa zb) \u003d H 2 S + FeSO 4

NaCl (T) + H 2 SO 4 (conc) = HCl + NaHSO 4

AgNO 3 + HCl = AgCl↓ + HNO 3

CaCO 3 + 2HBr \u003d CaBr 2 + CO 2 + H 2 O

4. Some acids can be obtained using redox reactions.

H 2 O 2 + SO 2 \u003d H 2 SO 4

3P + 5HNO 3 + 2H 2 O \u003d ZH 3 PO 4 + 5NO 2

Sour taste, action on indicators, electrical conductivity, interaction with metals, basic and amphoteric oxides, bases and salts, formation of esters with alcohols - these properties are common to inorganic and organic acids.

can be divided into two types of reactions:

1) are common For acids the reactions are associated with the formation of hydronium ion H 3 O + in aqueous solutions;

2) specific(i.e. characteristic) reactions specific acids.

The hydrogen ion can enter into redox reactions, reducing to hydrogen, as well as in a compound reaction with negatively charged or neutral particles having lone pairs of electrons, i.e. in acid-base reactions.

The general properties of acids include the reactions of acids with metals in the series of voltages up to hydrogen, for example:

Zn + 2Н + = Zn 2+ + Н 2

Acid-base reactions include reactions with basic oxides and bases, as well as with medium, basic, and sometimes acidic salts.

2 CO 3 + 4HBr \u003d 2CuBr 2 + CO 2 + 3H 2 O

Mg (HCO 3) 2 + 2HCl \u003d MgCl 2 + 2CO 2 + 2H 2 O

2KHSO 3 + H 2 SO 4 \u003d K 2 SO 4 + 2SO 2 + 2H 2 O

Note that polybasic acids dissociate stepwise, and at each next step, dissociation is more difficult, therefore, with an excess of acid, acidic salts are most often formed, rather than medium ones.

Ca 3 (PO 4) 2 + 4H 3 PO 4 \u003d 3Ca (H 2 PO 4) 2

Na 2 S + H 3 PO 4 = Na 2 HPO 4 + H 2 S

NaOH + H 3 PO 4 = NaH 2 PO 4 + H 2 O

KOH + H 2 S \u003d KHS + H 2 O

At first glance, the formation of acidic salts may seem surprising. monobasic hydrofluoric (hydrofluoric) acid. However, this fact can be explained. Unlike all other hydrohalic acids, hydrofluoric acid is partially polymerized in solutions (due to the formation of hydrogen bonds) and different particles (HF) X can be present in it, namely H 2 F 2, H 3 F 3, etc.

A special case of acid-base balance - reactions of acids and bases with indicators that change color depending on the acidity of the solution. Indicators are used in qualitative analysis to detect acids and bases in solutions.

The most commonly used indicators are litmus(V neutral environment purple, V sour - red, V alkaline - blue), methyl orange(V sour environment red, V neutral - orange, V alkaline - yellow), phenolphthalein(V highly alkaline environment crimson red, V neutral and acidic - colorless).

Specific Properties different acids can be of two types: first, the reactions leading to the formation insoluble salts, and, secondly, redox transformations. If the reactions associated with the presence of an H + ion in them are common to all acids (qualitative reactions for detecting acids), specific reactions are used as qualitative reactions for individual acids:

Ag + + Cl - = AgCl (white precipitate)

Ba 2+ + SO 4 2- \u003d BaSO 4 (white precipitate)

3Ag + + PO 4 3 - = Ag 3 PO 4 (yellow precipitate)

Some specific reactions of acids are due to their redox properties.

Anoxic acids in aqueous solution can only oxidize.

2KMnO 4 + 16HCl = 5Cl 2 + 2KSl + 2MnCl 2 + 8H 2 O

H 2 S + Br 2 \u003d S + 2HBg

Oxygen-containing acids can only be oxidized if the central atom in them is in a lower or intermediate oxidation state, as, for example, in sulfurous acid:

H 2 SO 3 + Cl 2 + H 2 O \u003d H 2 SO 4 + 2HCl

Many oxygen-containing acids, in which the central atom has the maximum oxidation state (S +6, N +5, Cr +6), exhibit the properties of strong oxidizing agents. Concentrated H 2 SO 4 is a strong oxidizing agent.

Cu + 2H 2 SO 4 (conc) = CuSO 4 + SO 2 + 2H 2 O

Pb + 4HNO 3 \u003d Pb (NO 3) 2 + 2NO 2 + 2H 2 O

C + 2H 2 SO 4 (conc) = CO 2 + 2SO 2 + 2H 2 O

It should be remembered that:

• Acid solutions react with metals that are in the electrochemical series of voltages to the left of hydrogen, subject to a number of conditions, the most important of which is the formation of a soluble salt as a result of the reaction. The interaction of HNO 3 and H 2 SO 4 (conc.) with metals proceeds differently.

Concentrated sulfuric acid in the cold passivates aluminum, iron, chromium.

• In water, acids dissociate into hydrogen cations and anions of acid residues, for example:

• Inorganic and organic acids interact with basic and amphoteric oxides, provided that a soluble salt is formed:

• Both those and other acids react with bases. Polybasic acids can form both medium and acidic salts (these are neutralization reactions):

• The reaction between acids and salts occurs only if a precipitate or gas is formed:

The interaction of H 3 PO 4 with limestone will stop due to the formation of the last insoluble precipitate Ca 3 (PO 4) 2 on the surface.

The features of the properties of nitric HNO 3 and concentrated sulfuric H 2 SO 4 (conc.) acids are due to the fact that when they interact with simple substances (metals and non-metals), not H + cations, but nitrate and sulfate ions will act as oxidizing agents. It is logical to expect that as a result of such reactions, not hydrogen H 2 is formed, but other substances are obtained: necessarily salt and water, as well as one of the products of the reduction of nitrate or sulfate ions, depending on the concentration of acids, the position of the metal in a series of voltages and reaction conditions (temperature, metal fineness, etc.).

These features of the chemical behavior of HNO 3 and H 2 SO 4 (conc.) clearly illustrate the thesis of the theory of chemical structure about the mutual influence of atoms in the molecules of substances.

The concepts of volatility and stability (stability) are often confused. Volatile acids are called acids, the molecules of which easily pass into a gaseous state, that is, they evaporate. For example, hydrochloric acid is a volatile but persistent, stable acid. The volatility of unstable acids cannot be judged. For example, non-volatile, insoluble silicic acid decomposes into water and SiO 2 . Aqueous solutions of hydrochloric, nitric, sulfuric, phosphoric and a number of other acids are colorless. An aqueous solution of chromic acid H 2 CrO 4 is yellow, permanganic acid HMnO 4 is raspberry.

Reference material for passing the test:

Mendeleev table

Solubility table

Simple substances similar to metallic elements in structure and a number of chemical and physical parameters are called amphoteric, i.e. these are the elements that exhibit chemical duality. It should be noted that these are not the metals themselves, but their salts or oxides. For example, oxides of some metals can have two properties, under some conditions they can exhibit the properties inherent in acids, in others, they behave like alkalis.

The main amphoteric metals include aluminum, zinc, chromium and some others.

The term amphoteric was introduced into circulation at the beginning of the 19th century. At that time, chemicals were separated on the basis of their similar properties, manifested in chemical reactions.

What are amphoteric metals

The list of metals that can be classified as amphoteric is quite large. Moreover, some of them can be called amphoteric, and some - conditionally.

Let's list the serial numbers of the substances under which they are located in the Periodic Table. The list includes groups 22 to 32, 40 to 51 and many more. For example, chromium, iron and a number of others can rightfully be called basic, and strontium and beryllium can also be attributed to the latter.

By the way, aluminum is considered the brightest representative of amphora metals.

It is its alloys that have been used for a long time in almost all industries. It is used to make elements of aircraft fuselages, car bodies, and kitchen utensils. It has become indispensable in the electrical industry and in the production of equipment for heating networks. Unlike many other metals, aluminum is constantly reactive. The oxide film that covers the surface of the metal resists oxidative processes. Under normal conditions, and in certain types of chemical reactions, aluminum can act as a reducing element.

This metal is able to interact with oxygen if it is crushed into many small particles. This type of operation requires the use of high temperatures. The reaction is accompanied by the release of a large amount of thermal energy. When the temperature rises to 200 ºC, aluminum reacts with sulfur. The thing is that aluminum, not always, under normal conditions, can react with hydrogen. Meanwhile, when it is mixed with other metals, different alloys can occur.

Another pronounced amphoteric metal is iron. This element has the number 26 and is located between cobalt and manganese. Iron is the most common element found in the earth's crust. Iron can be classified as a simple element, having a silvery white color and malleable, of course, when exposed to high temperatures. Can quickly begin to corrode at high temperatures. Iron, if placed in pure oxygen, completely burns out and can ignite in the open air.

Such a metal has the ability to quickly go into the stage of corrosion when exposed to high temperatures. Iron placed in pure oxygen completely burns out. Being in the air, a metallic substance quickly oxidizes due to excessive moisture, that is, it rusts. When burning in an oxygen mass, a kind of scale is formed, which is called iron oxide.

Properties of amphoteric metals

They are defined by the very concept of amphotericity. In the typical state, that is, at normal temperature and humidity, most metals are solids. None of the metals can be dissolved in water. Alkaline bases appear only after certain chemical reactions. In the course of the reaction, metal salts interact. It should be noted that safety rules require special care when carrying out this reaction.

The combination of amphoteric substances with oxides or acids themselves is the first to show the reaction that is inherent in bases. At the same time, if they are combined with bases, acidic properties will appear.

Heating amphoteric hydroxides causes them to decompose into water and oxide. In other words, the properties of amphoteric substances are very wide and require careful study, which can be carried out during a chemical reaction.

The properties of amphoteric elements can be understood by comparing them with the parameters of traditional materials. For example, most metals have a low ionization potential and this allows them to act as reducing agents in chemical processes.

Amphoteric - can show both reducing and oxidizing characteristics. However, there are compounds that are characterized by a negative level of oxidation.

Absolutely all known metals have the ability to form hydroxides and oxides.

All metals have the ability to form basic hydroxides and oxides. By the way, metals can enter into an oxidation reaction only with certain acids. For example, the reaction with nitric acid can proceed in different ways.

Amphoteric substances related to simple ones have clear differences in structure and features. Belonging to a certain class can be determined at a glance for some substances, so it is immediately clear that copper is a metal, but bromine is not.

How to distinguish metal from non-metal

The main difference is that metals donate electrons that are in an external electron cloud. Non-metals actively attract them.

All metals are good conductors of heat and electricity, non-metals are deprived of such an opportunity.

Bases of amphoteric metals

Under normal conditions, these substances do not dissolve in water and can be safely attributed to weak electrolytes. Such substances are obtained after the reaction of metal salts and alkali. These reactions are quite dangerous for those who produce them, and therefore, for example, to obtain zinc hydroxide, caustic soda must be slowly and carefully introduced into a container with zinc chloride, drop by drop.

At the same time, amphoteric - interact with acids as bases. That is, when performing a reaction between hydrochloric acid and zinc hydroxide, zinc chloride will appear. And when interacting with bases, they behave like acids.

13.1. Definitions

The most important classes of inorganic substances traditionally include simple substances (metals and non-metals), oxides (acidic, basic and amphoteric), hydroxides (part of acids, bases, amphoteric hydroxides) and salts. Substances belonging to the same class have similar chemical properties. But you already know that when distinguishing these classes, different classification features are used.
In this section, we will finally formulate the definitions of all the most important classes of chemical substances and see how these classes are distinguished.
Let's start with simple substances (classification according to the number of elements that make up the substance). They are usually divided into metals And nonmetals(Fig. 13.1- A).
You already know the definition of "metal".

From this definition it can be seen that the main feature that allows us to divide simple substances into metals and non-metals is the type of chemical bond.

In most non-metals, the bonds are covalent. But there are also noble gases (simple substances of the elements of group VIIIA), whose atoms in the solid and liquid state are connected only by intermolecular bonds. Hence the definition.

According to the chemical properties among metals, a group of so-called amphoteric metals. This name reflects the ability of these metals to react with both acids and alkalis (as amphoteric oxides or hydroxides) (Fig. 13.1- b).
In addition, due to chemical inertness among metals, noble metals. These include gold, ruthenium, rhodium, palladium, osmium, iridium, platinum. By tradition, a somewhat more reactive silver is also classified as a noble metal, but such inert metals as tantalum, niobium, and some others are not included. There are other classifications of metals, for example, in metallurgy, all metals are divided into black and colored relating iron and its alloys to ferrous metals.
From complex substances the most important are, first of all, oxides(see §2.5), but since their classification takes into account the acid-base properties of these compounds, we first recall what acids And grounds.

Thus, we isolate acids and bases from the total mass of compounds using two features: composition and chemical properties.
Based on their composition, acids are divided into oxygen-containing (oxoacids) And anoxic(Fig. 13.2).

It should be remembered that oxygen-containing acids in their structure are hydroxides.

Note. Traditionally, for oxygen-free acids, the word "acid" is used when it comes to a solution of the corresponding individual substance, for example: the substance HCl is called hydrogen chloride, and its aqueous solution is called hydrochloric or hydrochloric acid.

Now back to oxides. We referred oxides to the group acidic or major by how they react with water (or by whether they are made from acids or bases). But not all oxides react with water, but most of them react with acids or alkalis, so it is better to classify oxides by this property.

There are several oxides that under normal conditions do not react with either acids or alkalis. Such oxides are called non-salt-forming. This is, for example, CO, SiO, N 2 O, NO, MnO 2. In contrast to them, the remaining oxides are called salt-forming(Fig. 13.3).

As you know, most acids and bases are hydroxide. According to the ability of hydroxides to react with both acids and alkalis, among them (as well as among oxides) they distinguish amphoteric hydroxides(Fig. 13.4).

Now we have to define salts. The term "salt" has been used for a long time. With the development of science, its meaning has been repeatedly changed, expanded and refined. In the modern sense, a salt is an ionic compound, but traditionally salts do not include ionic oxides (since they are called basic oxides), ionic hydroxides (bases), as well as ionic hydrides, carbides, nitrides, etc. Therefore, we can simply say, What

It is possible to give another, more precise, definition of salts.

In giving this definition, oxonium salts are usually classified as both salts and acids.
Salts are classified according to their composition into sour, medium And main(Fig. 13.5).

That is, the anions of acid salts include hydrogen atoms bound by covalent bonds with other atoms of the anions and capable of breaking away under the action of bases.

Basic salts usually have a very complex composition and are often insoluble in water. A typical example of a basic salt is the mineral malachite Cu 2 (OH) 2 CO 3 .

As you can see, the most important classes of chemicals are distinguished according to different classification criteria. But no matter how we distinguish a class of substances, all substances of this class have common chemical properties.

In this chapter, you will learn about the most characteristic chemical properties of the substances representing these classes and the most important methods for obtaining them.

metal
1. Where in the natural system of elements are the elements that form metals, and where are the elements that form non-metals?
2. Write the formulas for five metals and five non-metals.
3.Compose the structural formulas of the following compounds:
(H 3 O) Cl, (H 3 O) 2 SO 4, HCl, H 2 S, H 2 SO 4, H 3 PO 4, H 2 CO 3, Ba (OH) 2, RbOH.
4. Which oxides correspond to the following hydroxides:
H 2 SO 4 , Ca (OH) 2 , H 3 PO 4 , Al (OH) 3 , HNO 3 , LiOH?
What is the nature (acidic or basic) of each of these oxides?
5. Find salts among the following substances. Make up their structural formulas.
KNO 2 , Al 2 O 3 , Al 2 S 3 , HCN, CS 2 , H 2 S, K 2 , SiCl 4 , CaSO 4 , AlPO 4
6. Make the structural formulas of the following acid salts:
NaHSO 4 , KHSO 3 , NaHCO 3 , Ca(H 2 PO 4) 2 , CaHPO 4 .

13.2. Metals

In crystals of metals and in their melts, atomic cores are connected by a single electron cloud of a metallic bond. Like a single atom of an element that forms a metal, a metal crystal has the ability to donate electrons. The tendency of a metal to give up electrons depends on its structure and, above all, on the size of atoms: the larger the atomic cores (that is, the larger the ionic radii), the easier the metal gives up electrons.
Metals are simple substances, so the oxidation state of the atoms in them is 0. Entering into reactions, metals almost always change the oxidation state of their atoms. The atoms of metals, having no tendency to accept electrons, can only give them away or socialize them. The electronegativity of these atoms is low, therefore, even when they form covalent bonds, metal atoms acquire a positive oxidation state. Therefore, all metals, to one degree or another, exhibit restorative properties. They react:
1) C non-metals(but not all and not with all):
4Li + O 2 \u003d 2Li 2 O,
3Mg + N 2 \u003d Mg 3 N 2 (when heated),
Fe + S = FeS (when heated).
The most active metals easily react with halogens and oxygen, and only lithium and magnesium react with very strong nitrogen molecules.
Reacting with oxygen, most metals form oxides, and the most active ones form peroxides (Na 2 O 2 , BaO 2) and other more complex compounds.
2) C oxides less active metals:
2Ca + MnO 2 \u003d 2CaO + Mn (when heated),
2Al + Fe 2 O 3 \u003d Al 2 O 3 + 2Fe (with preheating).
The possibility of these reactions occurring is determined by the general rule (RWRs proceed in the direction of the formation of weaker oxidizing and reducing agents) and depends not only on the activity of the metal (a more active, that is, more easily giving up its electrons, metal restores a less active one), but also on the energy of the oxide crystal lattice ( the reaction proceeds in the direction of the formation of a more "strong" oxide).
3) C acid solutions(§ 12.2):
Mg + 2H 3 O \u003d Mg 2B + H 2 + 2H 2 O, Fe + 2H 3 O \u003d Fe 2 + H 2 + 2H 2 O,
Mg + H 2 SO 4p \u003d MgSO 4p + H 2, Fe + 2HCl p \u003d FeCl 2p + H 2.
In this case, the possibility of a reaction is easily determined by the series of voltages (the reaction proceeds if the metal in the series of voltages is to the left of hydrogen).
4) C salt solutions(§ 12.2):

Fe + Cu 2 \u003d Fe 2 + Cu, Cu + 2Ag \u003d Cu 2 + 2Ag,
Fe + CuSO 4p = Cu + FeSO 4p, Cu + 2AgNO 3p = 2Ag + Cu(NO 3) 2p.
A series of voltages is also used here to determine whether a reaction can proceed.
5) In addition, the most active metals (alkali and alkaline earth) react with water (§ 11.4):
2Na + 2H 2 O \u003d 2Na + H 2 + 2OH, Ca + 2H 2 O \u003d Ca 2 + H 2 + 2OH,
2Na + 2H 2 O \u003d 2NaOH p + H 2, Ca + 2H 2 O \u003d Ca (OH) 2p + H 2.
In the second reaction, the formation of a Ca(OH) 2 precipitate is possible.
Most metals in industry get, restoring their oxides:
Fe 2 O 3 + 3CO = 2Fe + 3CO 2 (at high temperature),
MnO 2 + 2C = Mn + 2CO (at high temperature).
In the laboratory, hydrogen is often used for this:

The most active metals, both in industry and in the laboratory, are obtained by electrolysis (§ 9.9).
In the laboratory, less reactive metals can be reduced from solutions of their salts with more reactive metals (see § 12.2 for restrictions).

1. Why do metals not tend to exhibit oxidizing properties?
2. What does the chemical activity of metals primarily depend on?
3. Perform transformations
a) Li Li 2 O LiOH LiCl; b) NaCl Na Na 2 O 2;
c) FeO Fe FeS Fe 2 O 3; d) CuCl 2 Cu(OH) 2 CuO Cu CuBr 2 .
4. Restore the left parts of the equations:
a) ... = H 2 O + Cu;
b) ... = 3CO + 2Fe;
c) ... = 2Cr + Al 2 O 3
. Chemical properties of metals.

13.3. non-metals

Unlike metals, non-metals are very different from each other in their properties - both physical and chemical, and even in the type of structure. But, apart from the noble gases, in all non-metals the bond between atoms is covalent.
The atoms that make up non-metals have a tendency to attach electrons, but, forming simple substances, they cannot "satisfy" this tendency. Therefore, non-metals (to one degree or another) have a tendency to attach electrons, that is, they can show oxidizing properties. The oxidative activity of non-metals depends, on the one hand, on the size of the atoms (the smaller the atoms, the more active the substance), and on the other hand, on the strength of covalent bonds in a simple substance (the stronger the bonds, the less active the substance). In the formation of ionic compounds, the atoms of non-metals really add "extra" electrons, and in the formation of compounds with covalent bonds, they only shift common electron pairs in their direction. In both cases, the degree of oxidation decreases.
Non-metals can oxidize:
1) metals(substances more or less inclined to donate electrons):
3F 2 + 2Al \u003d 2AlF 3,
O 2 + 2Mg \u003d 2MgO (with preheating),
S + Fe = FeS (when heated),
2C + Ca \u003d CaC 2 (when heated).
2) other non-metals(less likely to accept electrons):
2F 2 + C \u003d CF 4 (when heated),
O 2 + S = SO 2 (with preheating),
S + H 2 \u003d H 2 S (when heated),
3) many complex substances:
4F 2 + CH 4 \u003d CF 4 + 4HF,
3O 2 + 4NH 3 \u003d 2N 2 + 6H 2 O (when heated),
Cl 2 + 2HBr = Br 2 + 2HCl.
Here, the possibility of the reaction proceeding is determined primarily by the strength of bonds in the reactants and reaction products and can be determined by calculating G.
The strongest oxidizing agent is fluorine. Oxygen and chlorine are slightly inferior to it (pay attention to their position in the system of elements).
Boron, graphite (and diamond), silicon and other simple substances formed by elements adjacent to the boundary between metals and non-metals exhibit oxidizing properties to a much lesser extent. Atoms of these elements are less likely to accept electrons. It is these substances (especially graphite and hydrogen) that are capable of exhibiting restorative properties:
2C + MnO 2 \u003d Mn + 2CO,
4H 2 + Fe 3 O 4 \u003d 3Fe + 4H 2 O.
You will study the rest of the chemical properties of non-metals in the following sections when you get acquainted with the chemistry of individual elements (as was the case with oxygen and hydrogen). There you will also learn how to obtain these substances.

1. Which of the following substances are non-metals: Be, C, Ne, Pt, Si, Sn, Se, Cs, Sc, Ar, Ra?
2. Give examples of non-metals, which under normal conditions are a) gases, b) liquids, c) solids.
3. Give examples of a) molecular and b) non-molecular simple substances.
4. Give three examples of chemical reactions in which a) chlorine and b) hydrogen exhibit oxidizing properties.
5. Give three examples of chemical reactions that are not in the text of the paragraph, in which hydrogen exhibits reducing properties.
6. Perform transformations:
a) P 4 P 4 O 10 H 3 PO 4; b) H 2 NaH H 2; c) Cl 2 NaCl Cl 2 .
Chemical properties of non-metals.

13.4. Basic oxides

You already know that all basic oxides are solid non-molecular substances with ionic bonds.
The main oxides are:
a) oxides of alkali and alkaline earth elements,
b) oxides of some other elements that form metals in lower oxidation states, for example: CrO, MnO, FeO, Ag 2 O, etc.

They include singly charged, doubly charged (very rarely triply charged cations) and oxide ions. The most characteristic Chemical properties basic oxides are precisely associated with the presence of doubly charged oxide ions (very strong base particles) in them. The chemical activity of basic oxides depends primarily on the strength of the ionic bond in their crystals.
1) All basic oxides react with solutions of strong acids (§ 12.5):
Li 2 O + 2H 3 O \u003d 2Li + 3H 2 O, NiO + 2H 3 O \u003d Ni 2 + 3H 2 O,
Li 2 O + 2HCl p \u003d 2LiCl p + H 2 O, NiO + H 2 SO 4p \u003d NiSO 4p + H 2 O.
In the first case, in addition to the reaction with oxonium ions, the reaction with water also proceeds, but since its rate is much lower, it can be neglected, especially since the same products are still obtained in the end.
The ability to react with a weak acid solution is determined both by the strength of the acid (the stronger the acid, the more active it is) and the strength of the bond in the oxide (the weaker the bond, the more active the oxide).
2) Oxides of alkali and alkaline earth metals react with water (§ 11.4):
Li 2 O + H 2 O \u003d 2Li + 2OH BaO + H 2 O \u003d Ba 2 + 2OH
Li 2 O + H 2 O \u003d 2LiOH p, BaO + H 2 O \u003d Ba (OH) 2p.
3) In addition, basic oxides react with acidic oxides:
BaO + CO 2 \u003d BaCO 3,
FeO + SO 3 \u003d FeSO 4,
Na 2 O + N 2 O 5 \u003d 2NaNO 3.
Depending on the chemical activity of those and other oxides, the reactions can proceed at ordinary temperature or when heated.
What is the reason for such reactions? Let us consider the reaction of formation of BaCO 3 from BaO and CO 2 . The reaction proceeds spontaneously, and the entropy in this reaction decreases (from two substances, solid and gaseous, one crystalline substance is formed), therefore, the reaction is exothermic. In exothermic reactions, the energy of formed bonds is greater than the energy of breaking bonds, therefore, the bond energy in BaCO 3 is greater than in the initial BaO and CO 2 . Both in the starting substances and in the products of the reaction, there are two types of chemical bonds: ionic and covalent. The ionic bond energy (lattice energy) in BaO is slightly higher than in BaCO 3 (the size of the carbonate ion is larger than that of the oxide ion), therefore, the energy of the O 2 + CO 2 system is greater than the energy of CO 3 2 .

+ Q

In other words, the CO 3 2 ion is more stable than the O 2 ion and the CO 2 molecule taken separately. And the greater stability of the carbonate ion (its lower internal energy) is associated with the charge distribution of this ion (– 2 e) by three oxygen atoms of the carbonate ion instead of one in the oxide ion (see also § 13.11).
4) Many basic oxides can be reduced to metal with a more active reducing metal or non-metal:
MnO + Ca = Mn + CaO (when heated),
FeO + H 2 \u003d Fe + H 2 O (when heated).
The possibility of such reactions occurring depends not only on the activity of the reducing agent, but also on the strength of the bonds in the initial and resulting oxide.
general way to get almost all basic oxides is the oxidation of the corresponding metal with oxygen. Oxides of sodium, potassium and some other very active metals (under these conditions they form peroxides and more complex compounds), as well as gold, silver, platinum and other very inactive metals (these metals do not react with oxygen) cannot be obtained in this way. Basic oxides can be obtained by thermal decomposition of the corresponding hydroxides, as well as some salts (for example, carbonates). So, magnesium oxide can be obtained in all three ways:
2Mg + O 2 \u003d 2MgO,
Mg (OH) 2 \u003d MgO + H 2 O,
MgCO 3 \u003d MgO + CO 2.

1. Make up the reaction equations:
a) Li 2 O + CO 2 b) Na 2 O + N 2 O 5 c) CaO + SO 3
d) Ag 2 O + HNO 3 e) MnO + HCl f) MgO + H 2 SO 4
2. Make up the equations of the reactions that occur during the implementation of the following transformations:
a) Mg MgO MgSO 4 b) Na 2 O Na 2 SO 3 NaCl
c) CoO Co CoCl 2 d) Fe Fe 3 O 4 FeO
3. A portion of nickel weighing 8.85 g was calcined in a stream of oxygen to obtain nickel(II) oxide, then treated with an excess of hydrochloric acid. A solution of sodium sulfide was added to the resulting solution until the precipitation ceased. Determine the mass of this sediment.
Chemical properties of basic oxides.

13.5. Acid oxides

All acid oxides are substances with covalent bond.
Acid oxides include:
a) oxides of elements that form non-metals,
b) some oxides of elements that form metals, if the metals in these oxides are in higher oxidation states, for example, CrO 3, Mn 2 O 7.
Among acid oxides there are substances that are gases at room temperature (for example: CO 2, N 2 O 3, SO 2, SeO 2), liquids (for example, Mn 2 O 7) and solids (for example: B 2 O 3, SiO 2, N 2 O 5, P 4 O 6, P 4 O 10, SO 3, I 2 O 5, CrO 3). Most acid oxides are molecular substances (exceptions are B 2 O 3, SiO 2, solid SO 3, CrO 3 and some others; there are also non-molecular modifications of P 2 O 5). But non-molecular acid oxides also become molecular upon transition to the gaseous state.
Acid oxides are characterized by the following Chemical properties.
1) All acidic oxides react with strong bases as with solid ones:
CO 2 + Ca (OH) 2 \u003d CaCO 3 + H 2 O
SiO 2 + 2KOH \u003d K 2 SiO 3 + H 2 O (when heated),
and with alkali solutions (§ 12.8):
SO 3 + 2OH \u003d SO 4 2 + H 2 O, N 2 O 5 + 2OH \u003d 2NO 3 + H 2 O,
SO 3 + 2NaOH p \u003d Na 2 SO 4p + H 2 O, N 2 O 5 + 2KOH p \u003d 2KNO 3p + H 2 O.
The reason for the occurrence of reactions with solid hydroxides is the same as with oxides (see § 13.4).
The most active acid oxides (SO 3 , CrO 3 , N 2 O 5 , Cl 2 O 7 ) can also react with insoluble (weak) bases.
2) Acidic oxides react with basic oxides (§ 13.4):
CO 2 + CaO = CaCO 3
P 4 O 10 + 6FeO = 2Fe 3 (PO 4) 2 (on heating)
3) Many acidic oxides react with water (§11.4).
N 2 O 3 + H 2 O \u003d 2HNO 2 SO 2 + H 2 O \u003d H 2 SO 3 (more correct writing of the formula of sulfurous acid -SO 2. H 2 O
N 2 O 5 + H 2 O \u003d 2HNO 3 SO 3 + H 2 O \u003d H 2 SO 4
Many acidic oxides can be received by oxidation with oxygen (combustion in oxygen or in air) of the corresponding simple substances (C gr, S 8, P 4, P cr, B, Se, but not N 2 and not halogens):
C + O 2 \u003d CO 2,
S 8 + 8O 2 \u003d 8SO 2,
or when decomposing the corresponding acids:
H 2 SO 4 \u003d SO 3 + H 2 O (with strong heating),
H 2 SiO 3 \u003d SiO 2 + H 2 O (when dried in air),
H 2 CO 3 \u003d CO 2 + H 2 O (at room temperature in solution),
H 2 SO 3 \u003d SO 2 + H 2 O (at room temperature in solution).
The instability of carbonic and sulfurous acids makes it possible to obtain CO 2 and SO 2 under the action of strong acids on carbonates Na 2 CO 3 + 2HCl p \u003d 2NaCl p + CO 2 + H 2 O
(the reaction proceeds both in solution and with solid Na 2 CO 3), and sulfites
K 2 SO 3tv + H 2 SO 4conc \u003d K 2 SO 4 + SO 2 + H 2 O (if there is a lot of water, sulfur dioxide is not released as a gas).

Amphoteric compounds

Chemistry is always a unity of opposites.

Look at the periodic table.

Some elements (almost all metals exhibiting oxidation states +1 and +2) form main oxides and hydroxides. For example, potassium forms the oxide K 2 O, and the hydroxide KOH. They exhibit basic properties, such as interacting with acids.

K2O + HCl → KCl + H2O

Some elements (most non-metals and metals with oxidation states +5, +6, +7) form acidic oxides and hydroxides. Acid hydroxides are oxygen-containing acids, they are called hydroxides because there is a hydroxyl group in the structure, for example, sulfur forms acid oxide SO 3 and acid hydroxide H 2 SO 4 (sulfuric acid):

Such compounds exhibit acidic properties, for example, they react with bases:

H2SO4 + 2KOH → K2SO4 + 2H2O

And there are elements that form such oxides and hydroxides that exhibit both acidic and basic properties. This phenomenon is called amphoteric . Such oxides and hydroxides will be the focus of our attention in this article. All amphoteric oxides and hydroxides are solids, insoluble in water.

First, how do you determine if an oxide or hydroxide is amphoteric? There is a rule, a little conditional, but you can still use it:

Amphoteric hydroxides and oxides are formed by metals, in oxidation states +3 and +4, For example (Al 2 O 3 , Al(Oh) 3 , Fe 2 O 3 , Fe(Oh) 3)

And four exceptions:metalsZn , Be , Pb , sn form the following oxides and hydroxides:ZnO , Zn ( Oh ) 2 , BeO , Be ( Oh ) 2 , PbO , Pb ( Oh ) 2 , SNO , sn ( Oh ) 2 , in which they exhibit an oxidation state of +2, but despite this, these compounds exhibit amphoteric properties .

The most common amphoteric oxides (and their corresponding hydroxides): ZnO, Zn(OH) 2 , BeO, Be(OH) 2 , PbO, Pb(OH) 2 , SnO, Sn(OH) 2 , Al 2 O 3 , Al (OH) 3 , Fe 2 O 3 , Fe(OH) 3 , Cr 2 O 3 , Cr(OH) 3 .

The properties of amphoteric compounds are not difficult to remember: they interact with acids and alkalis.

• with interaction with acids, everything is simple; in these reactions, amphoteric compounds behave like basic ones:

Al 2 O 3 + 6HCl → 2AlCl 3 + 3H 2 O

ZnO + H 2 SO 4 → ZnSO 4 + H 2 O

BeO + HNO 3 → Be(NO 3 ) 2 + H 2 O

Hydroxides react in the same way:

Fe(OH) 3 + 3HCl → FeCl 3 + 3H 2 O

Pb(OH) 2 + 2HCl → PbCl 2 + 2H 2 O

• With interaction with alkalis it is a little more difficult. In these reactions, amphoteric compounds behave like acids, and the reaction products can be different, it all depends on the conditions.

Either the reaction takes place in solution, or the reactants are taken as solids and fused.

Interaction of basic compounds with amphoteric compounds during fusion.

Let's take zinc hydroxide as an example. As mentioned earlier, amphoteric compounds interacting with basic ones behave like acids. So we write zinc hydroxide Zn (OH) 2 as an acid. The acid has hydrogen in front, let's take it out: H 2 ZnO 2. And the reaction of alkali with hydroxide will proceed as if it were an acid. "Acid residue" ZnO 2 2-divalent:

2K Oh(TV) + H 2 ZnO 2 (solid) (t, fusion) → K 2 ZnO 2 + 2 H 2 O

The resulting substance K 2 ZnO 2 is called potassium metazincate (or simply potassium zincate). This substance is a salt of potassium and the hypothetical "zinc acid" H 2 ZnO 2 (it is not entirely correct to call such compounds salts, but for our own convenience we will forget about it). Only zinc hydroxide is written like this: H 2 ZnO 2 is not good. We write as usual Zn (OH) 2, but we mean (for our own convenience) that this is an "acid":

2KOH (solid) + Zn (OH) 2 (solid) (t, fusion) → K 2 ZnO 2 + 2H 2 O

With hydroxides, in which there are 2 OH groups, everything will be the same as with zinc:

Be (OH) 2 (solid.) + 2NaOH (solid.) (t, fusion) → 2H 2 O + Na 2 BeO 2 (sodium metaberyllate, or beryllate)

Pb (OH) 2 (solid.) + 2NaOH (solid.) (t, fusion) → 2H 2 O + Na 2 PbO 2 (sodium metaplumbate, or plumbate)

With amphoteric hydroxides with three OH groups (Al (OH) 3, Cr (OH) 3, Fe (OH) 3) a little differently.

Let's take aluminum hydroxide as an example: Al (OH) 3, write it in the form of an acid: H 3 AlO 3, but we don’t leave it in this form, but take out the water from there:

H 3 AlO 3 - H 2 O → HAlO 2 + H 2 O.

Here we are working with this “acid” (HAlO 2):

HAlO 2 + KOH → H 2 O + KAlO 2 (potassium metaaluminate, or simply aluminate)

But aluminum hydroxide cannot be written like this HAlO 2, we write it down as usual, but we mean “acid” there:

Al (OH) 3 (solid.) + KOH (solid.) (t, fusion) → 2H 2 O + KAlO 2 (potassium metaaluminate)

The same is true for chromium hydroxide:

Cr(OH) 3 → H 3 CrO 3 → HCrO 2

Cr (OH) 3 (solid.) + KOH (solid.) (t, fusion) → 2H 2 O + KCrO 2 (potassium metachromate,

BUT NOT CHROMATE, chromates are salts of chromic acid).

With hydroxides containing four OH groups, it is exactly the same: we bring hydrogen forward and remove water:

Sn(OH) 4 → H 4 SnO 4 → H 2 SnO 3

Pb(OH) 4 → H 4 PbO 4 → H 2 PbO 3

It should be remembered that lead and tin form two amphoteric hydroxides each: with an oxidation state of +2 (Sn (OH) 2, Pb (OH) 2), and +4 (Sn (OH) 4, Pb (OH) 4).

And these hydroxides will form different "salts":

Oxidation state
Hydroxide Formula	Sn(OH)2	Pb (OH) 2	Sn(OH)4	Pb(OH)4
Formula of hydroxide as acid	H2SnO2	H2PbO2	H2SnO3	H2PbO3
Salt (potassium)	K2SnO2	K 2 PbO 2	K2SnO3	K2PbO3
Salt name			metastannat	metablumbAT

The same principles as in the names of ordinary "salts", the element in the highest degree of oxidation - the suffix AT, in the intermediate - IT.

Such "salts" (metachromates, metaaluminates, metaberyllates, metazincates, etc.) are obtained not only as a result of the interaction of alkalis and amphoteric hydroxides. These compounds are always formed when a strongly basic "world" and an amphoteric one (by fusion) come into contact. That is, just like amphoteric hydroxides with alkalis, both amphoteric oxides and metal salts forming amphoteric oxides (salts of weak acids) will react. And instead of alkali, you can take a strongly basic oxide, and a salt of a metal that forms an alkali (salt of a weak acid).

Interactions:

Remember, the reactions below take place during fusion.

Amphoteric oxide with strongly basic oxide:

ZnO (solid) + K 2 O (solid) (t, fusion) → K 2 ZnO 2 (potassium metazincate, or simply potassium zincate)

Amphoteric oxide with alkali:

ZnO (solid) + 2KOH (solid) (t, fusion) → K 2 ZnO 2 + H 2 O

Amphoteric oxide with a salt of a weak acid and an alkali-forming metal:

ZnO (solid) + K 2 CO 3 (solid) (t, fusion) → K 2 ZnO 2 + CO 2

Amphoteric hydroxide with strongly basic oxide:

Zn (OH) 2 (solid) + K 2 O (solid) (t, fusion) → K 2 ZnO 2 + H 2 O

Amphoteric hydroxide with alkali:

Zn (OH) 2 (solid) + 2KOH (solid) (t, fusion) → K 2 ZnO 2 + 2H 2 O

Amphoteric hydroxide with a salt of a weak acid and an alkali-forming metal:

Zn (OH) 2 (solid) + K 2 CO 3 (solid) (t, fusion) → K 2 ZnO 2 + CO 2 + H 2 O

Salts of a weak acid and a metal that forms an amphoteric compound with a strongly basic oxide:

ZnCO 3 (solid) + K 2 O (solid) (t, fusion) → K 2 ZnO 2 + CO 2

Salts of a weak acid and a metal that forms an amphoteric compound with an alkali:

ZnCO 3 (solid) + 2KOH (solid) (t, fusion) → K 2 ZnO 2 + CO 2 + H 2 O

Salts of a weak acid and a metal that forms an amphoteric compound with a salt of a weak acid and a metal that forms an alkali:

ZnCO 3 (solid) + K 2 CO 3 (solid) (t, fusion) → K 2 ZnO 2 + 2CO 2

Below is information on salts of amphoteric hydroxides, the most common in the exam are marked in red.

	Hydroxide	Acid hydroxide	acid residue		Salt name
BeO	Be(OH) 2	H 2 BeO 2	BeO 2 2-	K 2 BeO 2	Metaberyllate (beryllate)
ZnO	Zn(OH) 2	H 2 ZnO 2	ZnO 2 2-	K 2 ZnO 2	Metazincate (zincate)
Al 2 O 3	Al(OH) 3	HAlO 2	AlO 2 —	KALO 2	Metaaluminate (aluminate)
Fe2O3	Fe(OH)3	HFeO 2	FeO 2 -	KFeO 2	Metaferrate (BUT NOT FERRATE)
	Sn(OH)2	H2SnO2	SnO 2 2-	K2SnO2
	Pb(OH)2	H2PbO2	PbO 2 2-	K 2 PbO 2
SnO 2	Sn(OH)4	H2SnO3	SnO 3 2-	K2SnO3	MetastannAT (stannate)
PbO2	Pb(OH)4	H2PbO3	PbO 3 2-	K2PbO3	MetablumbAT (plumbat)
Cr2O3	Cr(OH)3	HCrO 2	CrO2 -	KCrO 2	Metachromat (BUT NOT CHROMATE)

Interaction of amphoteric compounds with alkali solutions (here only alkalis).

In the Unified State Examination, this is called "the dissolution of aluminum hydroxide (zinc, beryllium, etc.) alkali." This is due to the ability of metals in the composition of amphoteric hydroxides in the presence of an excess of hydroxide ions (in an alkaline medium) to attach these ions to themselves. A particle is formed with a metal (aluminum, beryllium, etc.) in the center, which is surrounded by hydroxide ions. This particle becomes negatively charged (anion) due to hydroxide ions, and this ion will be called hydroxoaluminate, hydroxozincate, hydroxoberyllate, etc. Moreover, the process can proceed in different ways, the metal can be surrounded by a different number of hydroxide ions.

We will consider two cases: when the metal is surrounded four hydroxide ions, and when it is surrounded six hydroxide ions.

Let us write down the abbreviated ionic equation of these processes:

Al(OH) 3 + OH - → Al(OH) 4 -

The resulting ion is called the tetrahydroxoaluminate ion. The prefix "tetra" is added because there are four hydroxide ions. The tetrahydroxoaluminate ion has a - charge, since aluminum carries a 3+ charge, and four hydroxide ions 4-, in total it turns out -.

Al (OH) 3 + 3OH - → Al (OH) 6 3-

The ion formed in this reaction is called the hexahydroxoaluminate ion. The prefix "hexo-" is added because there are six hydroxide ions.

It is necessary to add a prefix indicating the amount of hydroxide ions. Because if you just write "hydroxoaluminate", it is not clear which ion you mean: Al (OH) 4 - or Al (OH) 6 3-.

When alkali reacts with amphoteric hydroxide, a salt is formed in solution. The cation of which is an alkali cation, and the anion is a complex ion, the formation of which we considered earlier. The anion is in square brackets.

Al (OH) 3 + KOH → K (potassium tetrahydroxoaluminate)

Al (OH) 3 + 3KOH → K 3 (potassium hexahydroxoaluminate)

What exactly (hexa- or tetra-) salt you write as a product does not matter. Even in the USE answers it is written: “... K 3 (the formation of K is acceptable". The main thing is not to forget to make sure that all indices are correctly affixed. Keep track of the charges, and keep in mind that their sum should be equal to zero.

In addition to amphoteric hydroxides, amphoteric oxides react with alkalis. The product will be the same. Only if you write the reaction like this:

Al 2 O 3 + NaOH → Na

Al 2 O 3 + NaOH → Na 3

But these reactions will not equalize. It is necessary to add water to the left side, because interaction occurs in solution, there is enough water there, and everything will equalize:

Al 2 O 3 + 2NaOH + 3H 2 O → 2Na

Al 2 O 3 + 6NaOH + 3H 2 O → 2Na 3

In addition to amphoteric oxides and hydroxides, some especially active metals interact with alkali solutions, which form amphoteric compounds. Namely, it is: aluminum, zinc and beryllium. To equalize, the left also needs water. And, in addition, the main difference between these processes is the release of hydrogen:

2Al + 2NaOH + 6H 2 O → 2Na + 3H 2

2Al + 6NaOH + 6H 2 O → 2Na 3 + 3H 2

The table below shows the most common examples of the properties of amphoteric compounds in the exam:

Amphoteric substance		Salt name
Al2O3 Al(OH)3		Sodium tetrahydroxoaluminate	Al(OH) 3 + NaOH → Na Al 2 O 3 + 2NaOH + 3H 2 O → 2Na 2Al + 2NaOH + 6H 2 O → 2Na + 3H 2
	Na 3	Sodium hexahydroxoaluminate	Al(OH) 3 + 3NaOH → Na 3 Al 2 O 3 + 6NaOH + 3H 2 O → 2Na 3 2Al + 6NaOH + 6H 2 O → 2Na 3 + 3H 2
Zn(OH) 2	K2	Sodium tetrahydroxozincate	Zn(OH) 2 + 2NaOH → Na 2 ZnO + 2NaOH + H 2 O → Na 2 Zn + 2NaOH + 2H 2 O → Na 2 + H 2
	K4	Sodium hexahydroxozincate	Zn(OH) 2 + 4NaOH → Na 4 ZnO + 4NaOH + H 2 O → Na 4 Zn + 4NaOH + 2H 2 O → Na 4 + H 2
Be(OH)2	Li 2	Lithium tetrahydroxoberyllate	Be(OH) 2 + 2LiOH → Li 2 BeO + 2LiOH + H 2 O → Li 2 Be + 2LiOH + 2H 2 O → Li 2 + H 2
	Li 4	Lithium hexahydroxoberyllate	Be(OH) 2 + 4LiOH → Li 4 BeO + 4LiOH + H 2 O → Li 4 Be + 4LiOH + 2H 2 O → Li 4 + H 2
Cr2O3 Cr(OH)3		Sodium tetrahydroxochromate	Cr(OH) 3 + NaOH → Na Cr 2 O 3 + 2NaOH + 3H 2 O → 2Na
	Na 3	Sodium hexahydroxochromate	Cr(OH) 3 + 3NaOH → Na 3 Cr 2 O 3 + 6NaOH + 3H 2 O → 2Na 3
Fe2O3 Fe(OH)3		Sodium tetrahydroxoferrate	Fe(OH) 3 + NaOH → Na Fe 2 O 3 + 2NaOH + 3H 2 O → 2Na
	Na 3	Sodium hexahydroxoferrate	Fe(OH) 3 + 3NaOH → Na 3 Fe 2 O 3 + 6NaOH + 3H 2 O → 2Na 3

The salts obtained in these interactions react with acids, forming two other salts (salts of a given acid and two metals):

2Na 3 + 6H 2 SO 4 → 3Na 2 SO 4 + Al 2 (SO 4 ) 3 + 12H 2 O

That's all! Nothing complicated. The main thing is not to confuse, remember what is formed during fusion, what is in solution. Very often, tasks on this issue come across in B parts.