Oxides are a type of compounds widespread in nature, which can be observed even in everyday life, in everyday life. An example is sand, water, rust, lime, carbon dioxide, a number of natural dyes. The ore of many valuable metals is oxide by its nature, which is why it is of great interest for scientific and industrial research.

The combination of chemical elements with oxygen is called oxides. As a rule, they are formed when any substances are heated in air. Distinguish between acidic and basic oxides. Metals form basic oxides, while non-metals form acidic ones. With the exception of oxides of chromium and manganese, which are also acidic. This article discusses the representative of the main oxides - CuO (II).

CuO(II)

Copper, heated in air at a temperature of 400–500 °C, gradually covered with a black coating, which chemists call divalent copper oxide, or CuO (II). The described phenomenon is represented in the following equation:

2 Cu + O 2 → 2 CuO

The term "bivalent" indicates the ability of an atom to react with other elements through two chemical bonds.

Interesting fact! Copper, being in various compounds, can be with different valencies and a different color. For example: copper oxides are bright red (Cu2O) and brown-black (CuO) in color. And copper hydroxides acquire yellow (CuOH) and blue (Cu (OH) 2) colors. A classic example of the phenomenon when quantity turns into quality.

Cu2O is sometimes also called nitrous oxide, copper (I) oxide, and CuO is oxide, copper (II) oxide. There is also copper (III) oxide - Cu2O3.

In geology, the oxide of divalent (or bivalent) copper is commonly called tenorite, its other name is melaconite. The name tenorite comes from the name of the outstanding Italian professor of botany Michele Tenore, (1780-1861). Melakonite is considered a synonym for the name tenorite and is translated into Russian as copper black or black copper ore. In one case or another, we are talking about a brown-black crystalline mineral that decomposes when calcined and melts only at an excess pressure of oxygen, insoluble in water and does not react with it.

We emphasize the main parameters of the named mineral.

Chemical formula: CuO

Its molecule consists from a Cu atom with a molecular weight of 64 a. e. m. and an O atom, molecular weight 16 a.m. e. m., where a. e. m. - atomic mass unit, it is also a dalton, 1 a. mu \u003d 1.660 540 2 (10) × 10 -27 kg \u003d 1.660 540 2 (10) × 10 -24 g. Accordingly, the molecular weight of the compound is: 64 + 16 \u003d 80 a. eat.

Crystal cell: monoclinic system. What does this type of crystal symmetry axes mean when two axes intersect at an oblique angle and have different lengths, and the third axis is located at an angle of 90 ° with respect to them.

Density – 6.51 g/cm3. For comparison, the density of pure gold is 19.32 g / cm³, and the density of table salt is 2.16 g / cm 3.

Melts at 1447 °C, under oxygen pressure.

Decomposes upon incandescence up to 1100 °C and is converted to copper (I) oxide:

4CuO = 2Cu2O + O2.

It does not react with water and does not dissolve in it..

But it reacts with an aqueous solution of ammonia, with the formation of tetraamminecopper (II) hydroxide: CuO + 4NH3 + H2O = (OH) 2.

In an acidic environment, it forms sulfate and water: CuO + H2SO4 = CuSO4 + H2O.

Reacting with alkali, it creates cuprate: CuO + 2 NaOH → Na2CuO2 + H2O.

Reaction CuO NaOH

Formed:

• by calcining copper (II) hydroxide at a temperature of 200 ° C: Cu (OH) 2 \u003d CuO + H2O;
• during the oxidation of metallic copper in air at a temperature of 400–500 °C: 2Cu + O2 = 2CuO;
• during high-temperature processing of malachite: (CuOH)₂CO₃ -> 2CuO + CO₂ + H₂O.

Reduced to metallic copper -

• in reaction with hydrogen: CuO + H2 = Cu + H2O;
• with carbon monoxide (carbon monoxide): CuO + CO = Cu + CO2;
• with active metal: CuO + Mg = Cu + MgO.

toxic. According to the degree of adverse effects on the human body, it is classified as a substance of the second hazard class. Causes irritation of the mucous membranes of the eyes, skin, respiratory tract and gastrointestinal system. When interacting with him, it is necessary to use such protective equipment as rubber gloves, respirators, goggles, overalls.

The substance is explosive and flammable.

Applied in industry, as a mineral component of feed, in pyrotechnics, in the production of catalysts for chemical reactions, as a coloring pigment for glass, enamels, and ceramics.

The oxidizing properties of copper oxide (II) are most often used in laboratory studies, when elemental analysis is required related to the study of organic materials for the presence of hydrogen and carbon in them.

It is important that CuO (II) is quite widespread in nature as the mineral tenerite, in other words, it is a natural ore compound from which copper can be obtained.

Latin name Cuprum and the corresponding symbol Cu comes from the name of the island of Cyprus. It was from there, through the Mediterranean Sea, that the ancient Romans and Greeks exported this valuable metal.

Copper is one of the seven most common metals in the world and has been in the service of man since ancient times. However, in its original, metallic state, it is quite rare. This is a soft, easy-to-work metal, characterized by a high density, a very high-quality conductor of current and heat. In terms of electrical conductivity, it is second only to silver, while it is a cheaper material. Widely used in the form of wire and thin sheet products.

Chemical compounds of copper are different increased biological activity. In animal and plant organisms, they are involved in the synthesis of chlorophyll, therefore they are considered a very valuable component in the composition of mineral fertilizers.

Copper is also needed in the human diet. Its deficiency in the body can lead to various blood diseases.

Video

From the video you will learn what copper oxide is.

Copper (Cu) belongs to the d-elements and is located in the IB group of the periodic table of D.I. Mendeleev. The electronic configuration of the copper atom in the ground state is written as 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 1 instead of the expected formula 1s 2 2s 2 2p 6 3s 2 3p 6 3d 9 4s 2 . In other words, in the case of a copper atom, the so-called “electron jump” from the 4s sublevel to the 3d sublevel is observed. For copper, in addition to zero, oxidation states +1 and +2 are possible. The +1 oxidation state is prone to disproportionation and is stable only in insoluble compounds such as CuI, CuCl, Cu 2 O, etc., as well as in complex compounds, for example, Cl and OH. Copper compounds in the +1 oxidation state do not have a specific color. So, copper (I) oxide, depending on the size of the crystals, can be dark red (large crystals) and yellow (small crystals), CuCl and CuI are white, and Cu 2 S is black-blue. More chemically stable is the oxidation state of copper, equal to +2. Salts containing copper in a given oxidation state are blue and blue-green in color.

Copper is a very soft, malleable and ductile metal with high electrical and thermal conductivity. The color of metallic copper is red-pink. Copper is in the activity series of metals to the right of hydrogen, i.e. refers to low-active metals.

with oxygen

Under normal conditions, copper does not interact with oxygen. Heat is required for the reaction between them to proceed. Depending on the excess or lack of oxygen and temperature conditions, it can form copper (II) oxide and copper (I) oxide:

with sulfur

The reaction of sulfur with copper, depending on the conditions of carrying out, can lead to the formation of both copper (I) sulfide and copper (II) sulfide. When a mixture of powdered Cu and S is heated to a temperature of 300-400 ° C, copper (I) sulfide is formed:

With a lack of sulfur and the reaction is carried out at a temperature of more than 400 ° C, copper (II) sulfide is formed. However, a simpler way to obtain copper (II) sulfide from simple substances is the interaction of copper with sulfur dissolved in carbon disulfide:

This reaction proceeds at room temperature.

with halogens

Copper reacts with fluorine, chlorine and bromine, forming halides with the general formula CuHal 2, where Hal is F, Cl or Br:

Cu + Br 2 = CuBr 2

In the case of iodine, the weakest oxidizing agent among halogens, copper (I) iodide is formed:

Copper does not interact with hydrogen, nitrogen, carbon and silicon.

with non-oxidizing acids

Almost all acids are non-oxidizing acids, except for concentrated sulfuric acid and nitric acid of any concentration. Since non-oxidizing acids are able to oxidize only metals that are in the activity series up to hydrogen; this means that copper does not react with such acids.

with oxidizing acids

- concentrated sulfuric acid

Copper reacts with concentrated sulfuric acid both when heated and at room temperature. When heated, the reaction proceeds in accordance with the equation:

Since copper is not a strong reducing agent, sulfur is reduced in this reaction only to the +4 oxidation state (in SO 2).

- with dilute nitric acid

The reaction of copper with dilute HNO 3 leads to the formation of copper (II) nitrate and nitrogen monoxide:

3Cu + 8HNO 3 (diff.) = 3Cu(NO 3) 2 + 2NO + 4H 2 O

- with concentrated nitric acid

Concentrated HNO 3 readily reacts with copper under normal conditions. The difference between the reaction of copper with concentrated nitric acid and the interaction with dilute nitric acid lies in the product of nitrogen reduction. In the case of concentrated HNO 3, nitrogen is reduced to a lesser extent: instead of nitric oxide (II), nitric oxide (IV) is formed, which is associated with greater competition between nitric acid molecules in concentrated acid for the electrons of the reducing agent (Cu):

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

with non-metal oxides

Copper reacts with some non-metal oxides. For example, with oxides such as NO 2 , NO, N 2 O, copper is oxidized to copper (II) oxide, and nitrogen is reduced to oxidation state 0, i.e. a simple substance N 2 is formed:

In the case of sulfur dioxide, instead of a simple substance (sulfur), copper (I) sulfide is formed. This is due to the fact that copper with sulfur, unlike nitrogen, reacts:

with metal oxides

When sintering metallic copper with copper oxide (II) at a temperature of 1000-2000 ° C, copper oxide (I) can be obtained:

Also, metallic copper can reduce iron (III) oxide upon calcination to iron (II) oxide:

with metal salts

Copper displaces less active metals (to the right of it in the activity series) from solutions of their salts:

Cu + 2AgNO 3 \u003d Cu (NO 3) 2 + 2Ag ↓

An interesting reaction also takes place, in which copper is dissolved in a salt of a more active metal - iron in the +3 oxidation state. However, there are no contradictions, because copper does not displace iron from its salt, but only restores it from the +3 oxidation state to the +2 oxidation state:

Fe 2 (SO 4) 3 + Cu \u003d CuSO 4 + 2FeSO 4

Cu + 2FeCl 3 = CuCl 2 + 2FeCl 2

The latter reaction is used in the production of microcircuits at the stage of etching of copper boards.

Corrosion of copper

Copper corrodes over time when exposed to moisture, carbon dioxide and atmospheric oxygen:

2Cu + H 2 O + CO 2 + O 2 \u003d (CuOH) 2 CO 3

As a result of this reaction, copper products are covered with a loose blue-green coating of copper (II) hydroxocarbonate.

Chemical properties of zinc

Zinc Zn is in the IIB group of the IVth period. Electronic configuration of valence orbitals of atoms of a chemical element in the ground state 3d 10 4s 2 . For zinc, only one single oxidation state is possible, equal to +2. Zinc oxide ZnO and zinc hydroxide Zn(OH) 2 have pronounced amphoteric properties.

Zinc tarnishes when stored in air, becoming covered with a thin layer of ZnO oxide. Oxidation proceeds especially easily at high humidity and in the presence of carbon dioxide due to the reaction:

2Zn + H 2 O + O 2 + CO 2 → Zn 2 (OH) 2 CO 3

Zinc vapor burns in air, and a thin strip of zinc, after glowing in a burner flame, burns in it with a greenish flame:

When heated, metallic zinc also interacts with halogens, sulfur, phosphorus:

Zinc does not directly react with hydrogen, nitrogen, carbon, silicon and boron.

Zinc reacts with non-oxidizing acids to release hydrogen:

Zn + H 2 SO 4 (20%) → ZnSO 4 + H 2

Zn + 2HCl → ZnCl 2 + H 2

Industrial zinc is especially easily soluble in acids, since it contains impurities of other less active metals, in particular, cadmium and copper. High-purity zinc is resistant to acids for certain reasons. To speed up the reaction, a sample of high purity zinc is brought into contact with copper, or a small amount of copper salt is added to the acid solution.

At a temperature of 800-900 o C (red heat), metallic zinc, being in a molten state, interacts with superheated water vapor, releasing hydrogen from it:

Zn + H 2 O \u003d ZnO + H 2

Zinc also reacts with oxidizing acids: concentrated sulfuric and nitric.

Zinc as an active metal can form sulfur dioxide, elemental sulfur and even hydrogen sulfide with concentrated sulfuric acid.

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

The composition of the products of nitric acid reduction is determined by the concentration of the solution:

Zn + 4HNO 3 (conc.) = Zn(NO 3) 2 + 2NO 2 + 2H 2 O

3Zn + 8HNO 3 (40%) = 3Zn(NO 3) 2 + 2NO + 4H 2 O

4Zn + 10HNO 3 (20%) = 4Zn (NO 3) 2 + N 2 O + 5H 2 O

5Zn + 12HNO 3 (6%) = 5Zn(NO 3) 2 + N 2 + 6H 2 O

4Zn + 10HNO 3 (0.5%) = 4Zn(NO 3) 2 + NH 4 NO 3 + 3H 2 O

The direction of the process is also affected by the temperature, the amount of acid, the purity of the metal, and the reaction time.

Zinc reacts with alkali solutions to form tetrahydroxozincates and hydrogen:

Zn + 2NaOH + 2H 2 O \u003d Na 2 + H 2

Zn + Ba (OH) 2 + 2H 2 O \u003d Ba + H 2

With anhydrous alkalis, zinc, when fused, forms zincates and hydrogen:

In a highly alkaline environment, zinc is an extremely strong reducing agent, capable of reducing nitrogen in nitrates and nitrites to ammonia:

4Zn + NaNO 3 + 7NaOH + 6H 2 O → 4Na 2 + NH 3

Due to complexation, zinc slowly dissolves in an ammonia solution, reducing hydrogen:

Zn + 4NH 3 H 2 O → (OH) 2 + H 2 + 2H 2 O

Zinc also restores less active metals (to the right of it in the activity series) from aqueous solutions of their salts:

Zn + CuCl 2 \u003d Cu + ZnCl 2

Zn + FeSO 4 \u003d Fe + ZnSO 4

Chemical properties of chromium

Chromium is an element of the VIB group of the periodic table. The electronic configuration of the chromium atom is written as 1s 2 2s 2 2p 6 3s 2 3p 6 3d 5 4s 1, i.e. in the case of chromium, as well as in the case of the copper atom, the so-called "electron slip" is observed

The most frequently exhibited oxidation states of chromium are +2, +3 and +6. They should be remembered, and within the framework of the USE program in chemistry, we can assume that chromium has no other oxidation states.

Under normal conditions, chromium is resistant to corrosion both in air and in water.

Interaction with non-metals

with oxygen

Heated to a temperature of more than 600 o C, powdered metallic chromium burns in pure oxygen to form chromium (III) oxide:

4Cr + 3O 2 = o t=> 2Cr 2 O 3

with halogens

Chromium reacts with chlorine and fluorine at lower temperatures than with oxygen (250 and 300 o C, respectively):

2Cr + 3F 2 = o t=> 2CrF 3

2Cr + 3Cl 2 = o t=> 2CrCl 3

Chromium reacts with bromine at a red heat temperature (850-900 o C):

2Cr + 3Br 2 = o t=> 2CrBr 3

with nitrogen

Metallic chromium interacts with nitrogen at temperatures above 1000 o C:

2Cr + N 2 = ot=> 2CrN

with sulfur

With sulfur, chromium can form both chromium (II) sulfide and chromium (III) sulfide, depending on the proportions of sulfur and chromium:

Cr+S= o t=> CRS

2Cr+3S= o t=> Cr 2 S 3

Chromium does not react with hydrogen.

Interaction with complex substances

Interaction with water

Chromium belongs to the metals of medium activity (located in the activity series of metals between aluminum and hydrogen). This means that the reaction proceeds between red-hot chromium and superheated water vapor:

2Cr + 3H 2 O = o t=> Cr 2 O 3 + 3H 2

Interaction with acids

Chromium, under normal conditions, is passivated by concentrated sulfuric and nitric acids, however, it dissolves in them during boiling, while being oxidized to an oxidation state of +3:

Cr + 6HNO 3 (conc.) = t o=> Cr(NO 3) 3 + 3NO 2 + 3H 2 O

2Cr + 6H 2 SO 4 (conc) = t o=> Cr 2 (SO 4) 3 + 3SO 2 + 6H 2 O

In the case of dilute nitric acid, the main product of nitrogen reduction is a simple substance N 2:

10Cr + 36HNO 3 (razb) \u003d 10Cr (NO 3) 3 + 3N 2 + 18H 2 O

Chromium is located in the activity series to the left of hydrogen, which means that it is able to release H 2 from solutions of non-oxidizing acids. In the course of such reactions, in the absence of access to atmospheric oxygen, chromium (II) salts are formed:

Cr + 2HCl \u003d CrCl 2 + H 2

Cr + H 2 SO 4 (razb.) \u003d CrSO 4 + H 2

When carrying out the reaction in the open air, divalent chromium is instantly oxidized by oxygen contained in the air to an oxidation state of +3. In this case, for example, the equation with hydrochloric acid will take the form:

4Cr + 12HCl + 3O 2 = 4CrCl 3 + 6H 2 O

When chromium metal is fused with strong oxidizing agents in the presence of alkalis, chromium is oxidized to an oxidation state of +6, forming chromates:

Chemical properties of iron

Iron Fe, a chemical element in group VIIIB and having serial number 26 in the periodic table. The distribution of electrons in an iron atom is as follows 26 Fe1s 2 2s 2 2p 6 3s 2 3p 6 3d 6 4s 2 , that is, iron belongs to d-elements, since the d-sublevel is filled in its case. It is most characteristic of two oxidation states +2 and +3. FeO oxide and Fe(OH) 2 hydroxide are dominated by basic properties, Fe 2 O 3 oxide and Fe(OH) 3 hydroxide are markedly amphoteric. So the oxide and hydroxide of iron (lll) dissolve to some extent when boiled in concentrated solutions of alkalis, and also react with anhydrous alkalis during fusion. It should be noted that the oxidation state of iron +2 is very unstable, and easily passes into the oxidation state +3. Iron compounds are also known in a rare oxidation state of +6 - ferrates, salts of the non-existent "iron acid" H 2 FeO 4. These compounds are relatively stable only in the solid state or in strongly alkaline solutions. With insufficient alkalinity of the medium, ferrates quickly oxidize even water, releasing oxygen from it.

Interaction with simple substances

With oxygen

When burned in pure oxygen, iron forms the so-called iron scale, having the formula Fe 3 O 4 and actually representing a mixed oxide, the composition of which can be conditionally represented by the formula FeO∙Fe 2 O 3 . The combustion reaction of iron has the form:

3Fe + 2O 2 = t o=> Fe 3 O 4

With sulfur

When heated, iron reacts with sulfur to form ferrous sulfide:

Fe+S= t o=> FeS

Or with an excess of sulfur iron disulfide:

Fe + 2S = t o=> FeS2

With halogens

With all halogens except iodine, metallic iron is oxidized to an oxidation state of +3, forming iron halides (lll):

2Fe + 3F 2 = t o=> 2FeF 3 - iron fluoride (lll)

2Fe + 3Cl 2 = t o=> 2FeCl 3 - iron chloride (lll)

Iodine, as the weakest oxidizing agent among halogens, oxidizes iron only to the +2 oxidation state:

Fe + I 2 = t o=> FeI 2 - iron iodide (ll)

It should be noted that ferric iron compounds easily oxidize iodide ions in an aqueous solution to free iodine I 2 while recovering to the +2 oxidation state. Examples of similar reactions from the FIPI bank:

2FeCl 3 + 2KI = 2FeCl 2 + I 2 + 2KCl

2Fe(OH) 3 + 6HI = 2FeI 2 + I 2 + 6H 2 O

Fe 2 O 3 + 6HI \u003d 2FeI 2 + I 2 + 3H 2 O

With hydrogen

Iron does not react with hydrogen (only alkali metals and alkaline earth metals react with hydrogen from metals):

Interaction with complex substances

Interaction with acids

With non-oxidizing acids

Since iron is located in the activity series to the left of hydrogen, this means that it is able to displace hydrogen from non-oxidizing acids (almost all acids except H 2 SO 4 (conc.) and HNO 3 of any concentration):

Fe + H 2 SO 4 (diff.) \u003d FeSO 4 + H 2

Fe + 2HCl \u003d FeCl 2 + H 2

It is necessary to pay attention to such a trick in the tasks of the exam, as a question on the topic to what degree of oxidation iron will be oxidized when it is exposed to dilute and concentrated hydrochloric acid. The correct answer is up to +2 in both cases.

The trap here lies in the intuitive expectation of a deeper oxidation of iron (up to s.o. +3) in the case of its interaction with concentrated hydrochloric acid.

Interaction with oxidizing acids

Under normal conditions, iron does not react with concentrated sulfuric and nitric acids due to passivation. However, it reacts with them when boiled:

2Fe + 6H 2 SO 4 = o t=> Fe 2 (SO 4) 3 + 3SO 2 + 6H 2 O

Fe + 6HNO 3 = o t=> Fe(NO 3) 3 + 3NO 2 + 3H 2 O

Note that dilute sulfuric acid oxidizes iron to an oxidation state of +2, and concentrated to +3.

Corrosion (rusting) of iron

In moist air, iron rusts very quickly:

4Fe + 6H 2 O + 3O 2 \u003d 4Fe (OH) 3

Iron does not react with water in the absence of oxygen either under normal conditions or when boiled. The reaction with water proceeds only at a temperature above the red heat temperature (> 800 ° C). those..

Like all d-elements, brightly colored.

Just like with copper, it is observed electron dip- from s-orbital to d-orbital

The electronic structure of the atom:

Accordingly, there are 2 characteristic oxidation states of copper: +2 and +1.

Simple substance: gold-pink metal.

Copper oxides:Сu2O copper oxide (I) \ copper oxide 1 - red-orange color

CuO copper (II) oxide \ copper oxide 2 - black.

Other copper compounds Cu(I), except for the oxide, are unstable.

Copper compounds Cu (II) - firstly, they are stable, and secondly, they are blue or greenish in color.

Why do copper coins turn green? Copper reacts with carbon dioxide in the presence of water to form CuCO3, a green substance.

Another colored copper compound, copper (II) sulfide, is a black precipitate.

Copper, unlike other elements, stands after hydrogen, so it does not release it from acids:

• With hot sulfuric acid: Сu + 2H2SO4 = CuSO4 + SO2 + 2H2O
• With cold sulfuric acid: Cu + H2SO4 = CuO + SO2 + H2O
• with concentrated:
  Cu + 4HNO3 = Cu(NO3)2 + 4NO2 + 4H2O
• with dilute nitric acid:
  3Cu + 8HNO3 = 3 Cu(NO3)2 + 2NO +4 H2O

An example of the task of the exam C2 option 1:

Copper nitrate was calcined, the resulting solid precipitate was dissolved in sulfuric acid. Hydrogen sulfide was passed through the solution, the resulting black precipitate was calcined, and the solid residue was dissolved by heating in nitric acid.

2Сu(NO3)2 → 2CuO↓ +4 NO2 + O2

The solid precipitate is copper(II) oxide.

CuO + H2S → CuS↓ + H2O

Copper(II) sulfide is a black precipitate.

“Fired” means that there was an interaction with oxygen. Do not confuse with "calcination". Ignite - heat, naturally, at a high temperature.

2СuS + 3O2 = 2CuO + 2SO2

The solid residue is CuO if the copper sulfide reacted completely, CuO + CuS if partially.

СuO + 2HNO3 = Cu(NO3)2 + H2O

CuS + 2HNO3 = Cu(NO3)2 + H2S

another reaction is also possible:

СuS + 8HNO3 = Cu(NO3)2 + SO2 + 6NO2 + 4H2O

An example of the task of the exam C2 option 2:

Copper was dissolved in concentrated nitric acid, the resulting gas was mixed with oxygen and dissolved in water. Zinc oxide was dissolved in the resulting solution, then a large excess of sodium hydroxide solution was added to the solution.

As a result of the reaction with nitric acid, Cu(NO3)2, NO2 and O2 are formed.

NO2 mixed with oxygen means oxidized: 2NO2 + 5O2 = 2N2O5. Mixed with water: N2O5 + H2O = 2HNO3.

ZnO + 2HNO3 = Zn(NO3)2 + 2H2O

Zn(NO 3) 2 + 4NaOH \u003d Na 2 + 2NaNO 3

Cuprum (Cu) is one of the low-active metals. It is characterized by the formation of chemical compounds with oxidation states +1 and +2. So, for example, two oxides, which are a compound of two elements Cu and oxygen O: with an oxidation state of +1 - copper oxide Cu2O and an oxidation state of +2 - copper oxide CuO. Despite the fact that they consist of the same chemical elements, but each of them has its own special characteristics. In the cold, the metal interacts very weakly with atmospheric oxygen, becoming covered with a film, which is copper oxide, which prevents further oxidation of cuprum. When heated, this simple substance with serial number 29 in the periodic table is completely oxidized. In this case, copper (II) oxide is also formed: 2Cu + O2 → 2CuO.

The nitrous oxide is a brownish red solid with a molar mass of 143.1 g/mol. The compound has a melting point of 1235°C, a boiling point of 1800°C. It is insoluble in water, but soluble in acids. Copper (I) oxide is diluted in (concentrated), and a colorless complex + is formed, which is easily oxidized in air to a blue-violet ammonium complex 2+, which dissolves in hydrochloric acid to form CuCl2. In the history of semiconductor physics, Cu2O is one of the most studied materials.

Copper(I) oxide, also known as hemioxide, has basic properties. It can be obtained by metal oxidation: 4Cu + O2 → 2 Cu2O. Impurities such as water and acids affect the rate of this process as well as further oxidation to the divalent oxide. Copper oxide can dissolve in this form pure metal and salt: H2SO4 + Cu2O → Cu + CuSO4 + H2O. According to a similar scheme, an oxide with a degree of +1 interacts with other oxygen-containing acids. In the interaction of hemioxide with halogen-containing acids, monovalent metal salts are formed: 2HCl + Cu2O → 2CuCl + H2O.

Oxide of copper (I) occurs in nature in the form of red ore (this is an outdated name, along with such as ruby ​​​​Cu), called the mineral "Cuprite". It takes a long time to educate. It can be produced artificially at high temperatures or under high oxygen pressure. Hemioxide is commonly used as a fungicide, as a pigment, as an antifouling agent in underwater or marine paint, and as a catalyst.

However, the effect of this substance with the chemical formula Cu2O on the body can be dangerous. If inhaled, it causes dyspnoea, coughing, and ulceration and perforation of the respiratory tract. If ingested, it irritates the gastrointestinal tract, which is accompanied by vomiting, pain and diarrhea.

H2 + CuO → Cu + H2O;

CO + CuO → Cu + CO2.

Copper(II) oxide is used in ceramics (as a pigment) to produce glazes (blue, green, and red, and sometimes pink, gray, or black). It is also used as a dietary supplement in animals to reduce cuprum deficiency in the body. It is an abrasive material that is necessary for polishing optical equipment. It is used for the production of dry cells, for the production of other Cu salts. The CuO compound is also used in the welding of copper alloys.

Exposure to the chemical compound CuO can also be dangerous to the human body. Causes lung irritation if inhaled. Copper(II) oxide can cause metal vapor fever (MFF). Cu oxide provokes a change in skin color, vision problems may appear. When ingested, like hemioxide, it leads to poisoning, which is accompanied by symptoms in the form of vomiting and pain.

COPPER AND ITS COMPOUNDS

LESSON IN THE 11th NATURAL SCIENCE CLASS

To increase the cognitive activity and independence of students, we use the lessons of the collective study of the material. At such lessons, each student (or a pair of students) receives a task, the completion of which he must report on in the same lesson, and his report is recorded by the rest of the class in notebooks and is an element of the content of the lesson's educational material. Each student contributes to the study of the topic by the class.
During the lesson, the mode of work of students changes from intraactive (a mode in which information flows are closed within the students, typical for independent work) to interactive (a mode in which information flows are two-way, i.e. information goes both from the student and to the student, there is an exchange of information). At the same time, the teacher acts as the organizer of the process, corrects and supplements the information provided by the students.
The lessons of collective study of the material consist of the following stages:
1st stage - installation, in which the teacher explains the goals and program of work in the lesson (up to 7 minutes);
Stage 2 - independent work of students according to the instructions (up to 15 minutes);
Stage 3 - exchange of information and summing up the lesson (takes all the remaining time).
The lesson "Copper and its compounds" is designed for classes with in-depth study of chemistry (4 hours of chemistry per week), is held for two academic hours, the lesson updates students' knowledge on the following topics: "General properties of metals", "Relationship to metals of concentrated sulfuric acid, nitric acid", "Qualitative reactions to aldehydes and polyhydric alcohols", "Oxidation of saturated monohydric alcohols with copper (II) oxide", "Complex compounds".
Before the lesson, students receive homework: to review the topics listed. The preliminary preparation of the teacher for the lesson consists in compiling instructional cards for students and preparing sets for laboratory experiments.

DURING THE CLASSES

Installation stage

The teacher puts in front of the students the purpose of the lesson: based on existing knowledge about the properties of substances, predict, confirm in practice, generalize information about copper and its compounds.
Students make up the electronic formula of the copper atom, find out what oxidation states copper can exhibit in compounds, what properties (redox, acid-base) copper compounds will have.
A table appears in the students' notebooks.

Properties of copper and its compounds

Metal	Cu 2 O - basic oxide	CuO - basic oxide
Reducing agent	CuOH is an unstable base	Cu (OH) 2 - insoluble base
	CuCl - insoluble salt	CuSO 4 - soluble salt
	Possess redox duality	Oxidizers

Stage of independent work

To confirm and supplement the assumptions, students perform laboratory experiments according to the instructions and write down the equations of the reactions performed.

Instructions for independent work in pairs

1. Ignite the copper wire in a flame. Note how its color has changed. Place the hot calcined copper wire in ethyl alcohol. Note the change in its color. Repeat these manipulations 2-3 times. Check if the smell of ethanol has changed.
Write down two reaction equations corresponding to the transformations carried out. What properties of copper and its oxide are confirmed by these reactions?

2. Add hydrochloric acid to copper(I) oxide.
What are you watching? Write down the reaction equations, given that copper (I) chloride is an insoluble compound. What properties of copper(I) are confirmed by these reactions?

3. a) Place a zinc granule into the copper(II) sulfate solution. If no reaction occurs, heat the solution. b) Add 1 ml of sulfuric acid to copper (II) oxide and heat.
What are you watching? Write down the reaction equations. What properties of copper compounds are confirmed by these reactions?

4. Place a universal indicator strip into the copper(II) sulfate solution.
Explain the result. Write down the ionic equation of hydrolysis for the first stage.
Add a solution of honey(II) sulfate to a solution of sodium carbonate.
What are you watching? Write the equation for the reaction of joint hydrolysis in molecular and ionic forms.

5.
What are you watching?
Add ammonia solution to the resulting precipitate.
What changes have taken place? Write down the reaction equations. What properties of copper compounds are proved by the reactions carried out?

6. Add a solution of potassium iodide to copper(II) sulfate.
What are you watching? Write an equation for the reaction. What property of copper(II) does this reaction prove?

7. Place a small piece of copper wire into a test tube with 1 ml of concentrated nitric acid. Close the tube with a stopper.
What are you watching? (Take the test tube under draft.) Write down the reaction equation.
Pour hydrochloric acid into another test tube, place a small piece of copper wire in it.
What are you watching? Explain your observations. What properties of copper are confirmed by these reactions?

8. Add an excess of sodium hydroxide to copper(II) sulfate.
What are you watching? Heat up the precipitate. What happened? Write down the reaction equations. What properties of copper compounds are confirmed by these reactions?

9. Add an excess of sodium hydroxide to copper(II) sulfate.
What are you watching?
Add a solution of glycerin to the resulting precipitate.
What changes have taken place? Write down the reaction equations. What properties of copper compounds prove these reactions?

10. Add an excess of sodium hydroxide to copper(II) sulfate.
What are you watching?
Pour the glucose solution to the resulting precipitate and heat.
What happened? Write the reaction equation using the general formula for aldehydes to denote glucose

What property of the copper compound does this reaction prove?

11. Add to copper(II) sulfate: a) ammonia solution; b) sodium phosphate solution.
What are you watching? Write down the reaction equations. What properties of copper compounds are proved by the reactions carried out?

Phase of communication and debriefing

The teacher asks a question concerning the properties of a particular substance. The students who performed the corresponding experiments report on the experiment and write down the reaction equations on the blackboard. Then the teacher and students complete the information about the chemical properties of the substance, which could not be confirmed by reactions in the conditions of the school laboratory.

The order of discussion of the chemical properties of copper compounds

1. How does copper react with acids, what other substances can copper react with?

The reactions of copper are written with:

Concentrated and dilute nitric acid:

Cu + 4HNO 3 (conc.) = Cu(NO 3) 2 + 2NO 2 + 2H 2 O,
3Cu + 8HNO 3 (diff.) = 3Cu(NO 3) 2 + 2NO + 4H 2 O;

Concentrated sulfuric acid:

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

Oxygen:

2Cu + O 2 \u003d 2CuO;

Cu + Cl 2 \u003d CuCl 2;

Hydrochloric acid in the presence of oxygen:

2Cu + 4HCl + O 2 = 2CuCl 2 + 2H 2 O;

Iron(III) chloride:

2FeCl 3 + Cu \u003d CuCl 2 + 2FeCl 2.

2. What are the properties of copper(I) oxide and chloride?

Attention is drawn to the main properties, the ability to complex formation, redox duality. The equations of reactions of copper (I) oxide with:

Hydrochloric acid to form CuCl:

Cu 2 O + 2HCl = 2CuCl + H 2 O;

Excess HCl:

CuCl + HCl = H;

Reactions of reduction and oxidation of Cu 2 O:

Cu 2 O + H 2 \u003d 2Cu + H 2 O,

2Cu 2 O + O 2 \u003d 4CuO;

Disproportionation when heated:

Cu 2 O \u003d Cu + CuO,
2CuCl \u003d Cu + CuCl 2.

3. What are the properties of copper(II) oxide?

Attention is drawn to the basic and oxidizing properties. Equations for the reactions of copper(II) oxide with:

Acid:

CuO + 2H + = Cu 2+ + H 2 O;

Ethanol:

C 2 H 5 OH + CuO = CH 3 CHO + Cu + H 2 O;

Hydrogen:

CuO + H 2 \u003d Cu + H 2 O;

Aluminum:

3CuO + 2Al \u003d 3Cu + Al 2 O 3.

4. What are the properties of copper(II) hydroxide?

Attention is drawn to the oxidizing, basic properties, the ability to complex with organic and inorganic compounds. The reaction equations are written with:

Aldehyde:

RCHO + 2Cu(OH) 2 = RCOOH + Cu 2 O + 2H 2 O;

Acid:

Cu(OH) 2 + 2H + = Cu 2+ + 2H 2 O;

Ammonia:

Cu (OH) 2 + 4NH 3 \u003d (OH) 2;

Glycerin:

Decomposition reaction equation:

Cu (OH) 2 \u003d CuO + H 2 O.

5. What are the properties of copper(II) salts?

Attention is drawn to the reactions of ion exchange, hydrolysis, oxidizing properties, complexation. The equations for the reactions of copper sulfate are written with:

Sodium hydroxide:

Cu 2+ + 2OH - \u003d Cu (OH) 2;

Sodium Phosphate:

3Cu 2+ + 2= Cu 3 (PO 4) 2;

Cu 2+ + Zn \u003d Cu + Zn 2+;

Potassium iodide:

2CuSO 4 + 4KI = 2CuI + I 2 + 2K 2 SO 4 ;

Ammonia:

Cu 2+ + 4NH 3 \u003d 2+;

and reaction equations:

Hydrolysis:

Cu 2+ + HOH = CuOH + + H + ;

Co-hydrolysis with sodium carbonate to form malachite:

2Cu 2+ + 2 + H 2 O \u003d (CuOH) 2 CO 3 + CO 2.

In addition, you can tell students about the interaction of copper(II) oxide and hydroxide with alkalis, which proves their amphotericity:

Cu (OH) 2 + 2NaOH (conc.) \u003d Na 2,

Cu + Cl 2 \u003d CuCl 2,

Cu + HgCl 2 \u003d CuCl 2 + Hg,

2Cu + 4HCl + O 2 = 2CuCl 2 + 2H 2 O,

CuO + 2HCl \u003d CuCl 2 + H 2 O,

Cu(OH) 2 + 2HCl = CuCl 2 + 2H 2 O,

CuBr 2 + Cl 2 \u003d CuCl 2 + Br 2,

(CuOH) 2 CO 3 + 4HCl \u003d 2CuCl 2 + 3H 2 O + CO 2,

2CuCl + Cl 2 \u003d 2CuCl 2,

2CuCl \u003d CuCl 2 + Cu,

CuSO 4 + BaCl 2 \u003d CuCl 2 + BaSO 4.)

Exercise 3 Make chains of transformations corresponding to the following schemes and carry them out:

Task 1. An alloy of copper and aluminum was treated first with an excess of alkali and then with an excess of dilute nitric acid. Calculate the mass fractions of metals in the alloy, if it is known that the volumes of gases released in both reactions (under the same conditions) are equal to each other
.

(Answer . Mass fraction of copper - 84%.)

Task 2. On calcination of 6.05 g of hydrated copper(II) nitrate, 2 g of residue was obtained. Determine the formula of the original salt.

(Answer. Cu(NO 3) 2 3H 2 O.)

Task 3. A copper plate weighing 13.2 g was lowered into 300 g of an iron (III) nitrate solution with a mass fraction of salt of 0.112. When it was taken out, it turned out that the mass fraction of iron(III) nitrate became equal to the mass fraction of the formed copper(II) salt. Determine the mass of the plate after it has been removed from the solution.

(Answer. 10 y.)

Homework. Learn the material written in the notebook. Compose a chain of transformations for copper compounds, containing at least ten reactions, and carry it out.

LITERATURE

1. Puzakov S.A., Popkov V.A. A manual on chemistry for university students. Programs. Questions, exercises, tasks. Samples of exam papers. Moscow: Higher school, 1999, 575 p.
2. Kuzmenko N.E., Eremin V.V. 2000 tasks and exercises in chemistry. For schoolchildren and entrants. M.: 1st Federal Book Trade Company, 1998, 512 p.