Amphoteric oxides. Preparation, chemical properties, formation of medium and complex salts

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

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

The term amphotericity was coined at the beginning of the 19th century. At that time, chemical substances were separated based on their similar properties, manifested in chemical reactions.

What are amphoteric metals?

The list of amphoteric metals includes many items. Some of them can be confidently called amphoteric, some - presumably, others - conditionally. If we consider the issue on a large scale, then for brevity we can simply name the serial numbers of the above mentioned metals. These numbers are: 4.13, from 22 to 32, from 40 to 51, from 72 to 84, from 104 to 109. But there are metals that can be called basic. These include chromium, iron, aluminum and zinc. Strontium and beryllium complete the main group. The most common of all listed at the moment is aluminum. Its alloys have been used for many centuries in a wide variety of fields and applications. The metal has excellent anti-corrosion resistance and is easy to cast and various types of machining. In addition, the popularity of aluminum is complemented by such advantages as high thermal conductivity and good electrical conductivity.

Aluminum is an amphoteric metal that tends to exhibit chemical activity. The durability of this metal is determined by a strong oxide film and, under normal environmental conditions, during chemical reactions, aluminum acts as a reducing element. Such an amphoteric substance is capable of interacting with oxygen in the event of fragmentation of the metal into small particles. Such interaction requires the influence of high temperature conditions. A chemical reaction upon contact with an oxygen mass is accompanied by a huge release of thermal energy. At temperatures above 200 degrees, the interaction of reactions when combined with a substance such as sulfur forms aluminum sulfide. Amphoteric aluminum is not able to directly interact with hydrogen, and when this metal is mixed with other metal components, various alloys containing intermetallic compounds arise.

Iron is an amphoteric metal, which is one of the side subgroups of group 4 of the period in the system of elements of the chemical type. This element stands out as the most common component of the group of metallic substances in the components of the earth's crust. Iron is classified as a simple substance, among the distinctive properties of which are its malleability and silvery-white color. Such a metal has the ability to provoke an increased chemical reaction and quickly goes into the stage of corrosion when exposed to high temperatures. Iron placed in pure oxygen burns out completely, and when brought to a finely dispersed state it can spontaneously ignite in plain air. When exposed to air, a metallic substance quickly oxidizes due to excessive humidity, that is, it rusts. When burning in an oxygen mass, a kind of scale is formed, which is called iron oxide.

What elements are amphoteric?

Compounds that exhibit chemical duality are called amphoteric. They are divided into 3 groups:

1. Oxides – Cu2O, Cr2O, PbO2, PbO, SnO
2. Hydroxides - Al(OH)3, Fe(OH)3, Zn(OH)2.

The third group includes metals - aluminum, copper, iron, zinc, beryllium, lead, etc. They occupy a significant part in the periodic table of Mendeleev and are under serial numbers - 22–32, 40–51. Others go separately.

Representatives of metals:

1. Iron. Belongs to the group of amphoteric. It is a simple substance. Characteristic properties: silver-white color, malleability, versatility. If you place iron in pure oxygen, it will burn out completely, and if you make it finely dispersed, spontaneous combustion may occur in the open air. Iron oxide is formed when it burns in an environment saturated with pure oxygen. It is scale.
2. Aluminum. When exposed to air, it is coated with a durable oxide film that protects it from rust. If you crush it into small particles, it begins to interact with oxygen. Upon contact with oxygen, a large amount of heat is released. If you heat aluminum to 200°C, it begins to react with sulfur. As a result of this reaction, aluminum sulfide is formed.

Amphoteric metals are simple elements that are analogues of a group of metallic-type substances. Similarities can be seen in chemical and physical properties.

Where are amphoteric metals used?

Chemical properties of alkaline earth metals.
list of alkaline earth metals Areas of application:

1. Manufacturing of parts for seismic and speed sensors, clock mechanisms, torque.
2. Production of parts for equipment that will interact with aggressive factors.
3. Reinforcement of high pressure pipes.
4. Shipbuilding, aircraft construction.
5. Production of household appliances and tools. These include cutlery, tape measures, razor blades, and kitchen utensils.
6. Assembly of video recording equipment.

Every year more and more chemical compounds appear. Thanks to this, new amphoteric metals are discovered. They are called materials of the future, but their popularity is growing slowly. This is due to the high cost and small size of finished products.

https://youtube.com/watch?v=BZIhw3pQFQs

Amphoteric metal bases

GOST 9012-59. metals. Brinell hardness measurement method (with changes n 1, 2, 3, 4, 5)

Under normal conditions, these substances do not dissolve in water and can easily be classified as 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, sodium hydroxide must be introduced slowly and carefully, drop by drop, into a container with zinc chloride.

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

We will devote this lesson to the study of amphoteric oxides and hydroxides. Here we will talk about substances that have amphoteric (dual) properties and the characteristics of the chemical reactions that occur with them. But first, let’s repeat what acidic and basic oxides react with. Next we will consider examples of amphoteric oxides and hydroxides.

Topic: Introduction

Lesson: Amphoteric oxides and hydroxides

Amphoteric metals list - Metalist's Handbook

ChromiumAmphoteric chromium

Simple substances similar to metal elements in structure and a number of chemical and physical parameters are called amphoteric, i.e. these are those elements that exhibit chemical duality.

It should be noted that these are not metals themselves, but their salts or oxides.

For example, oxides of some metals can have two properties: under some conditions they can exhibit properties inherent in acids, while under others, they behave like alkalis.

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

The term amphotericity was coined at the beginning of the 19th century. At that time, chemical substances were separated based on their similar properties, manifested in chemical reactions.

Amphoteric metal bases

Metals

Under normal conditions, these substances do not dissolve in water and can easily be classified as 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, sodium hydroxide must be introduced slowly and carefully, drop by drop, into a container with zinc chloride.

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

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Historical reference

In the mid-19th century, 63 chemical elements were discovered . The original plan was to arrange the elements according to increasing atomic mass and divide them into groups. However, it was not possible to structure them, and the proposal of the chemist Nuland was not taken seriously due to attempts to connect chemistry and music.

In 1869, Dmitri Ivanovich Mendeleev first published his periodic table in the pages of the Journal of the Russian Chemical Society. He soon notified chemists around the world about his discovery. Mendeleev subsequently continued to refine and improve his table until it acquired its modern form. It was Mendeleev who managed to arrange the chemical elements in such a way that they changed not monotonously, but periodically. The theory was finally combined into the periodic law in 1871. Let's move on to considering non-metals and metals in the periodic table.

Interaction of oxides with acids

Basic and amphoteric oxides react with acids. In this case, salts and water are formed:

FeO + H2SO4 = FeSO4 + H2O

Non-salt-forming oxides do not react with acids at all, and acidic oxides do not react with acids in most cases.

When does an acidic oxide react with an acid?

When solving the multiple-choice part of the Unified State Exam, you should conditionally assume that acidic oxides do not react with either acidic oxides or acids, except in the following cases:

1) silicon dioxide, being an acidic oxide, reacts with hydrofluoric acid, dissolving in it. In particular, thanks to this reaction, glass can be dissolved in hydrofluoric acid. In the case of excess HF, the reaction equation has the form:

SiO2 + 6HF = H2[SiF6] + 2H2O,

and in case of HF deficiency:

SiO2 + 4HF = SiF4 + 2H2O

2) SO2, being an acidic oxide, easily reacts with hydrosulfide acid H2S according to the type of coproportionation:

S+4O2 + 2H2S-2 = 3S + 2H2O

3) Phosphorus (III) oxide P2O3 can react with oxidizing acids, which include concentrated sulfuric acid and nitric acid of any concentration. In this case, the oxidation state of phosphorus increases from +3 to +5:

4) Sulfur oxide (IV) SO2 can be oxidized with nitric acid taken in any concentration. In this case, the oxidation degree of sulfur increases from +4 to +6.

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 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 able to react with all higher acidic oxides corresponding to stable acids, for example, P2O5, SO3, N2O5, 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:

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 the salt Na[Al(OH)4], the salt Na3[Al(OH)6] 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

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

Receipt

A common method for preparing amphoteric hydroxides is precipitation with dilute alkali from solutions of salts of the corresponding amphoteric element, for example:

ZnSO4 + 2NaOH ⟶ Zn(OH)2↓ + Na2SO4{\displaystyle {\mathsf {ZnSO_{4}\ +\ 2NaOH\ \longrightarrow \ Zn(OH)_{2}\downarrow \ +\ Na_{2}SO_{ 4}}}}

In excess of alkali, the hydroxide precipitate will begin to dissolve:

Zn(OH)2+2NaOH→Na2Zn(OH)4{\displaystyle {\mathsf {Zn(OH)_{2}+2NaOH\rightarrow Na_{2}}}}

In some cases, during precipitation, not a hydroxide is formed, but an oxide hydrate of the corresponding element (for example, hydrates of iron(III), chromium(III), tin(II), etc. oxides). The chemical properties of such hydrates are for the most part similar to those of the corresponding hydroxides.

Grounds. Amphoteric hydroxides

Bases, their classification, properties, preparation

1. Which of the following substances are bases: LiOH, CH3COOH, Fe(OH)2, CH3NH2, H2SO3, Mg(OH)2?

Classification of bases

2.

Characterize calcium hydroxide Ca(OH)2 according to all classification criteria.

RECEIVING

3. Given substances: Fe(OH)2, Ca(OH)2, LiOH, Al(OH)3. Which substances are formed by the interaction of metals with water, and which by the action of an alkali on a salt solution?

Chemical properties of bases

1. Dissociation of bases with the formation of hydroxide ions OH-:

NaOH <=> Na+ + OH-LiOH <=> Li+ + OH-

2. Reaction with acids to form a salt (neutralization reaction):

Mg(OH)2 + 2HNO3 = Mg(NO3)2 + 2H2OMg(OH)2 + 2H+ = Mg2+ + 2H2O

3. Interaction of alkalis with acid oxides to form salt and water:

2NaOH + SiO2 = Na2SiO3 + H2O (when heated)Ca(OH)2 + CO2 = CaCO3↓ + H2O

4. The interaction of an alkali solution with solutions of various salts to form an insoluble base:

CuSO4 + 2NaOH = Cu(OH)2 ↓+ Na2SO4Cu2+ + 2OH- = Cu(OH)2 ↓

5. Decomposition of insoluble bases when heated to form metal oxide and water:

Cu(OH)2 = CuO + H2O (when heated)

6. Interaction of alkali solutions with some non-metals:

2NaOH + Cl2 = NaCl + NaClO + H2O (cold) 6NaOH + 3Cl2 = 5NaCl + NaClO3 + 3H2O (heated) 2NaOH + Si = Na2SiO3 + 2H2

1. Interaction of alkali with some metals (forming amphoteric compounds).

??? 4. Given substances: CaO, SO2, Ba(OH)2, HClO4, KCl, CuCl2.

a) Which of the following substances react with sodium hydroxide?

b) Write equations for possible reactions.

c) Which of the following reactions is a neutralization reaction?

5. What substances decompose when heated: Fe(OH)2, NaOH, Al(OH)3, Fe(OH)3, Ba(OH)2? Write equations for possible reactions.

6. Three test tubes contain solutions of sodium chloride, hydrochloric acid, and sodium hydroxide. How can you recognize these solutions chemically?

7. What mass of alkali NaOH must be in solution to react with 16 g of copper (II) sulfate to obtain a precipitate of copper (II) hydroxide?

Amphoteric hydroxides

Amphoteric hydroxides are hydroxides that, upon dissociation, simultaneously form both H+ cations and OH- hydroxide ions. Amphoteric hydroxides correspond to amphoteric oxides. For example, Al(OH)3, Zn(OH)2, Cr(OH)3, Be(OH)2 and others.

1) Interaction of amphoteric hydroxides with acids:

Al(OH)3 + 3HCl = AlCl3 + 3H2O

Al(OH)3 + 3H+ = Al3+ + 3H2O

2) Interaction of amphoteric hydroxides with alkali:

Al(OH)3 + NaOH = Na[Al(OH)4] (sodium tetrahydroxoaluminate)Zn(OH)2 + 2NaOH = Na2[Zn(OH)4] (sodium tetrahydroxoaluminate)

3) They exhibit the properties of insoluble bases - they decompose when heated to form oxide and water:

2Al(OH)3 → Al2O3 + 3H2O

??? 8. a) Give examples of reactions that prove the properties of zinc hydroxide.

b) In which of the following reactions does zinc hydroxide manifest itself as an acid?

c) In which of the following reactions does zinc hydroxide manifest itself as a base?

d) Write the reaction equation for the production of zinc hydroxide.

Abstract

Key words of the abstract: amphotericity, amphoteric oxides, amphoteric hydroxides,

Hydroxides are inorganic compounds formed by one or another element, oxygen and hydrogen.

Hydroxides necessarily contain hydrogen bonded to oxygen (O–H bond). In some hydroxides, hydrogen can directly combine with the atoms of the element.

For example, in H3PO3, two hydrogen atoms are bonded to oxygen atoms, and one is bonded to a phosphorus atom.

In some cases, hydroxides are products of hydration (combination with water) of the corresponding oxides, although many oxides do not interact directly with water. As a result of dehydration of hydroxides, the corresponding oxides are formed.

Hydroxides can be basic (bases), acidic (oxygen-containing acids), or amphoteric. Examples of basic hydroxides are NaOH, Ba(OH)2, Mg(OH)2. Examples of acid hydroxides are HClO4 (perchloric acid, higher chlorine hydroxide), H3PO4 (orthophosphoric acid, higher hydroxide of phosphorus), H2SO4 (sulfuric acid, higher sulfur hydroxide).

Graphic formulas of the listed hydroxides are given below. All hydroxides have an O–H bond:

THE CONCEPT OF AMPHOTERIC OXIDES AND HYDROXIDES

Amphotericity (from the Greek amphoteros - both) is the ability of chemical compounds to exhibit both acidic and basic properties, depending on the nature of the reagent with which the amphoteric substance enters into an acid-base interaction.

Amphoteric oxides and hydroxides are oxides and hydroxides that exhibit both basic and acidic properties. They react with both acids and bases. Amphoteric oxides correspond to amphoteric hydroxides, for example: BeO – Be(OH)2, Cr2O3 – Cr(OH)3.

Amphoteric hydroxides are practically insoluble in water. Their basic and acidic properties are weakly expressed; they are weak acids and weak bases.

Amphoteric oxides and hydroxides are, as a rule, metal oxides and hydroxides in which the oxidation state of the metal is +3, +4, sometimes +2.

Among the oxides of elements of the main subgroups, the following are amphoteric: BeO, Al2O3, SnO, SnO2, PbO, Sb2O3.

Amphoteric hydroxides are the following hydroxides of elements of the main subgroups: Be(OH)2, Al(OH)3, Pb(OH)2 and some others.

Oxides and hydroxides of elements of side subgroups , corresponding to high oxidation states, are usually acidic, for example: CrO3 (corresponds to H2CrO4), Mn2O7 (corresponds to HMnO4).

Lower oxides and hydroxides are characterized by a predominance of basic properties, for example: CrO and Cr(OH)2, MnO and Mn(OH)2.

Oxides and hydroxides, in which the oxidation states of elements are +3 and +4, are, as a rule, amphoteric: Cr2O3 and Cr(OH)3, Fe2O3 and Fe(OH)3. However, the last elements in the d-elements decades (for example, Zn) form amphoteric oxides and hydroxides even in low oxidation states, for example ZnO and Zn(OH)2.

CHEMICAL PROPERTIES OF AMPHOTERIC OXIDES AND HYDROXIDES

Let us consider the amphoteric properties of zinc oxide and hydroxide – ZnO and Zn(OH)2. Both substances react with acids:

ZnO + 2 HNO3 = Zn( NO3)2 + H2 O Zn( OH)2 + H2 SO4 = ZnSO4 + 2 H2 O

Zinc oxide and hydroxide also react with alkalis:

As a result of reactions with alkalis in solutions, complex ions are formed - tetrahydroxocincates (for example, Na2[Zn(OH)4]), and upon fusion - zincates (for example, Na2ZnO2). To create formulas for zincates, write the formula for zinc hydroxide Zn(OH)2 as the acid formula - H2ZnO2. The acidic residue of such an acid is the ZnO22– ion.

When these salts dissociate in solutions, ions are formed in which zinc is part of the anions, therefore they say that these compounds contain “zinc in anionic form.”

Let us prove the amphotericity of aluminum oxide and hydroxide Al(OH)3. It dissolves in acids to form salts, where aluminum is in cationic form:

Al(OH)3 + 3H+ = Al3+ + 3 H2 O

But aluminum hydroxide also interacts with alkalis. When fused with alkalis, aluminates (meta-aluminates) are formed:

Al(OH)3 + NaOH = NaAlO2 + 2 H2 O

To create aluminate formulas, write the formula for aluminum hydroxide Al(OH)3 as H3AlO3. Now imagine that 1 mole of water is split off from 1 mole of this compound. A compound HAlO2 is formed, in which the acidic residue is the AlO2– ion.

In the resulting salt, aluminum forms the AlO2– anion (aluminum in anionic form). In solutions, as a result of reaction with alkali, a complex tetrahydroxoaluminate ion [Al(OH)4]– is formed:

Al(OH)3 + NaOH = Na[ Al( OH)4] (sodium tetrahydroxyaluminate)

Some amphoteric oxides and hydroxides do not exhibit amphotericity under ordinary conditions and behave as basic compounds in solutions. The amphoteric nature of such compounds manifests itself under more severe conditions. For example, iron (III) hydroxide Fe(OH)3 easily reacts with acids:

Fe(OH)3 + 3H Cl = FeCl3 + 3H2O Fe(OH)3 + NaOH = NaFeO2 + 2H2O

A sign of this reaction will be the dissolution of the brown precipitate. But when adding an alkali solution, iron (III) hydroxide does not dissolve. Then, perhaps, consider it a base, insoluble in water? The fact is that in this case, salt is formed by melting with alkali.

Naturally, during melting, the hydroxide decomposes into oxide and water, and in reality the reaction takes place between the oxide and the alkali:

Fe2 O3 + 2NaOH = 2NaFeO2 + H2 O

Thus, iron (III) hydroxide can be classified as amphoteric hydroxides, although its basic properties predominate.

Lesson summary "Amphoteric oxides and hydroxides."

Next topic: "".

Interaction with simple substances

with oxygen

Without heating, beryllium and magnesium do not react with either atmospheric oxygen or pure oxygen due to the fact that they are covered with thin protective films consisting of BeO and MgO oxides, respectively. Their storage does not require any special methods of protection from air and moisture, unlike alkaline earth metals, which are stored under a layer of liquid inert to them, most often kerosene.

Be, Mg, Ca, Sr, when burned in oxygen, form oxides of the composition MeO, and Ba - a mixture of barium oxide (BaO) and barium peroxide (BaO2):

2Mg + O2 = 2MgO

2Ca + O2 = 2CaO

2Ba + O2 = 2BaO

Ba + O2 = BaO2

It should be noted that during the combustion of alkaline earth metals and magnesium in air, a side reaction of these metals with air nitrogen also occurs, as a result of which, in addition to metal compounds with oxygen, nitrides with the general formula Me3N2 are also formed.

with halogens

Beryllium reacts with halogens only at high temperatures, and the rest of the Group IIA metals - already at room temperature:

Mg + I2 = MgI2 – magnesium iodide

Ca + Br2 = CaBr2 – calcium bromide

Ba + Cl2 = BaCl2 – barium chloride

with non-metals of groups IV–VI

All metals of group IIA react when heated with all nonmetals of groups IV–VI, but depending on the position of the metal in the group, as well as the activity of the nonmetals, varying degrees of heating are required. Since beryllium is the most chemically inert among all Group IIA metals, its reactions with non-metals require a significantly higher temperature.

It should be noted that the reaction of metals with carbon can form carbides of different natures. There are carbides that belong to methanides and are conventionally considered derivatives of methane, in which all hydrogen atoms are replaced by metal. They, like methane, contain carbon in the -4 oxidation state, and when they are hydrolyzed or interact with non-oxidizing acids, one of the products is methane. There is also another type of carbides - acetylenides, which contain the C22- ion, which is actually a fragment of the acetylene molecule. Carbides such as acetylenides, upon hydrolysis or interaction with non-oxidizing acids, form acetylene as one of the reaction products. The type of carbide - methanide or acetylenide - obtained when a particular metal reacts with carbon depends on the size of the metal cation. Metal ions with a small radius usually form metanides, and larger ions form acetylenides. In the case of metals of the second group, methanide is obtained by the interaction of beryllium with carbon:

The remaining metals of group II A form acetylenides with carbon:

With silicon, group IIA metals form silicides - compounds of the Me2Si type, with nitrogen - nitrides (Me3N2), with phosphorus - phosphides (Me3P2):

with hydrogen

All alkaline earth metals react with hydrogen when heated. In order for magnesium to react with hydrogen, heating alone, as in the case of alkaline earth metals, is not enough; in addition to high temperature, increased hydrogen pressure is also required. Beryllium does not react with hydrogen under any conditions.

Application of oxides

Oxides are extremely common throughout the globe and are used both in everyday life and in industry. The most important oxide, hydrogen oxide, water, made life on Earth possible. Sulfur oxide SO3 is used to produce sulfuric acid, as well as for processing food products - this increases the shelf life of, for example, fruits.

Iron oxides are used to obtain paints and produce electrodes, although most iron oxides are reduced to metallic iron in metallurgy.

Calcium oxide, also known as quicklime, is used in construction. Zinc and titanium oxides are white and insoluble in water, which is why they have become a good material for the production of paints - white.

Silicon oxide SiO2 is the main component of glass. Chromium oxide Cr2O3 is used for the production of colored green glasses and ceramics, and due to its high strength properties, for polishing products (in the form of GOI paste).

Carbon monoxide CO2, which is emitted by all living organisms when they breathe, is used to extinguish fires, and also, in the form of dry ice, to cool something.

Receipt

A common method for preparing amphoteric hydroxides is precipitation with dilute alkali from solutions of salts of the corresponding amphoteric element, for example:

ZnSO4 + 2NaOH ⟶ Zn(OH)2↓ + Na2SO4{\displaystyle {\mathsf {ZnSO_{4}\ +\ 2NaOH\ \longrightarrow \ Zn(OH)_{2}\downarrow \ +\ Na_{2}SO_{ 4}}}}

In excess of alkali, the hydroxide precipitate will begin to dissolve:

Zn(OH)2+2NaOH→Na2Zn(OH)4{\displaystyle {\mathsf {Zn(OH)_{2}+2NaOH\rightarrow Na_{2}}}}

In some cases, during precipitation, not a hydroxide is formed, but an oxide hydrate of the corresponding element (for example, hydrates of iron(III), chromium(III), tin(II), etc. oxides). The chemical properties of such hydrates are for the most part similar to those of the corresponding hydroxides.

Which oxides react with water?

Of all the oxides, only the following react with water:

1) all active basic oxides (oxides of alkali metal and alkali metal);

2) all acid oxides, except silicon dioxide (SiO2);

those. From the above it follows that they definitely do not react with water:

1) all low-active basic oxides;

2) all amphoteric oxides;

3) non-salt-forming oxides (NO, N2O, CO, SiO).

Note:

Magnesium oxide reacts slowly with water when boiled. Without strong heating, the reaction of MgO with H2O does not occur.

The ability to determine which oxides can react with water even without the ability to write the corresponding reaction equations already allows you to get points for some questions in the test part of the Unified State Exam.

Now let's figure out how certain oxides react with water, i.e. Let's learn to write the corresponding reaction equations.

Active basic oxides react with water to form their corresponding hydroxides. Recall that the corresponding metal oxide is a hydroxide that contains the metal in the same oxidation state as the oxide. For example, when the active basic oxides K+12O and Ba+2O react with water, their corresponding hydroxides K+1OH and Ba+2(OH)2 are formed:

K2O + H2O = 2KOH – potassium hydroxide

BaO + H2O = Ba(OH)2 – barium hydroxide

All hydroxides corresponding to active basic oxides (alkaline metal and alkali metal oxides) belong to alkalis. Alkalis are all metal hydroxides that are highly soluble in water, as well as the slightly soluble calcium hydroxide Ca(OH)2 (as an exception).

The interaction of acidic oxides with water, as well as the reaction of active basic oxides with water, leads to the formation of the corresponding hydroxides. Only in the case of acidic oxides do they correspond not to basic ones, but to acidic hydroxides, more often called oxygen-containing acids. Let us recall that the corresponding acidic oxide is an oxygen-containing acid that contains an acid-forming element in the same oxidation state as in the oxide.

Thus, if we, for example, want to write down the equation for the interaction of the acidic oxide SO3 with water, first of all we must remember the basic sulfur-containing acids studied in the school curriculum. These are hydrogen sulfide H2S, sulfurous H2SO3 and sulfuric H2SO4 acids. Hydrogen sulfide acid H2S, as can be easily seen, does not contain oxygen, so its formation during the interaction of SO3 with water can be immediately excluded. Of the acids H2SO3 and H2SO4, only sulfuric acid H2SO4 contains sulfur in the oxidation state +6, as in SO3 oxide. Therefore, it is precisely this that will be formed in the reaction of SO3 with water:

H2O + SO3 = H2SO4

Similarly, the oxide N2O5, containing nitrogen in the oxidation state +5, reacting with water, forms nitric acid HNO3, but in no case nitrous HNO2, since in nitric acid the oxidation state of nitrogen, as in N2O5, is +5, and in nitrous — +3:

N+52O5 + H2O = 2HN+5O3

Exception:

Nitrogen (IV) oxide (NO2) is a non-metal oxide in the +4 oxidation state, i.e. in accordance with the algorithm described in the table at the very beginning of this chapter, it should be classified as acidic oxides. However, there is no acid that contains nitrogen in the +4 oxidation state.

In the case of NO2 oxide, it is generally accepted that two acids correspond to it at once, since its interaction with water leads to the simultaneous formation of two acids:

2NO2 + H2O = HNO2 + HNO3

Chrome, iron and copper

Hard metal of bluish-white color. The etymology of the word "chrome" originates from the Greek. χρῶμα - color, which is due to the wide variety of colors of chromium compounds. The mass fraction of this element in the earth's crust is 0.02% by mass.

Chromium has oxidation states of +2, +3 and +6. Compounds where chromium takes on an oxidation state of +2 have basic properties, +3 have amphoteric properties, and +6 have acidic properties.

In nature, chromium occurs in the form of the following compounds.

• Fe(CrO2)2 - chromium iron ore, chromite
• (Mg, Fe)Cr2O4 - magnochromite
• (Fe, Mg)(Cr, Al)2O4 - aluminum chromite

In industry, chromium is obtained by calcining chromium iron ore with carbon. Aluminothermy is also used to displace chromium from its oxide.

Already in air it reacts with oxygen: a film of chromium (III) oxide - Cr2O3 - is formed on the metal surface - passivation occurs. Reacts with non-metals when heated.

It flows in a hot state.

The reaction does not occur with cold concentrated sulfuric and nitric acid. It starts only when heated.

Reactions with salts of less active metals

Chromium is capable of displacing metals from salts that are to the right of it in the voltage series.

Chromium(II) compounds are basic in nature. Chromium (II) oxide is oxidized by atmospheric oxygen to a more stable form - chromium (III) oxide, and reacts with acids and acid oxides.

Chromium (II) hydroxide, as an insoluble hydroxide, easily decomposes when heated into the corresponding oxide and water, and reacts with acids and acid oxides.

These are the most stable compounds that are amphoteric in nature. These include chromium(III) oxide and chromium(III) hydroxide.

Chromium(III) oxide reacts with both alkalis and acids. In reactions with alkalis at normal temperature (in solution), complex salts are formed; upon calcination, mixed oxides are formed. With acids, chromium (III) oxide forms various salts.

H2O + NaOH + Cr2O3 → Na3[Cr(OH)6] (in solution, sodium hexahydroxochromate)

Cr2O3 + 2NaOH → (t°) 2NaCrO2 + H2O (calcination, sodium chromite)

Cr2O3 + HCl = CrCl3 + H2O (we maintain the oxidation state Cr +3)

Chromium(III) oxide reacts with more reactive metals (for example, during aluminothermy).

When oxidized, chromium(III) compounds produce chromium(VI) compounds (in an alkaline environment).

In this oxidation state, chromium exhibits acidic properties. These include chromium (VI) oxide - CrO3, and two acids that are in a state of equilibrium in solution: chromic acid - H2CrO4 and dichromic acid - H2Cr2O7.

It is fundamentally important to remember the color of chromates and dichromates (it is often given in tasks as a hint). Chromates color the solution yellow, and dichromates turn it orange.

Chromates transform into dichromates as the acidity of the medium increases (often in reactions with acids). The color of the solution changes from yellow to orange.

If alkali is added to an orange solution of dichromate, it will change its color to yellow - chromate is formed.

The decomposition of ammonium dichromate looks very impressive and is called “volcano”