There are a lot of representatives of each of them, but oxides undoubtedly occupy the leading position. One chemical element can have several different binary compounds with oxygen. Copper also has this property. She has three oxides. Let's look at them in more detail.

Copper(I) oxide

Its formula is Cu 2 O. In some sources, this compound may be called copper hemioxide, dicopper oxide, or cuprous oxide.

Properties

Is crystalline substance having a brown-red color. This oxide is insoluble in water and ethanol. It can melt without decomposing at a temperature of just over 1240 ° C. This substance does not interact with water, but can be transferred into solution if the participants in the reaction with it are concentrated hydrochloric acid, alkali, nitric acid, ammonia hydrate, ammonium salts, sulfuric acid .

Obtaining copper oxide (I)

It can be obtained by heating metallic copper, or in an environment where oxygen has a low concentration, as well as in a stream of certain nitrogen oxides and together with copper (II) oxide. In addition, it can become a reaction product of the thermal decomposition of the latter. Copper (I) oxide will also be obtained if copper (I) sulfide is heated in a stream of oxygen. There are others more difficult ways its production (for example, the reduction of one of the copper hydroxides, the ion exchange of any salt of monovalent copper with alkali, etc.), but they are practiced only in laboratories.

Application

Needed as a pigment when painting ceramics, glass; component of paints that protect the underwater part of the vessel from fouling. Also used as a fungicide. Copper oxide valves cannot do without it.

Copper(II) oxide

Its formula is CuO. In many sources it can be found under the name of copper oxide.

Properties

It is the highest copper oxide. The substance has the appearance of black crystals, which are almost insoluble in water. It reacts with acid and during this reaction forms the corresponding salt of divalent copper, as well as water. When it is fused with alkali, the reaction products are represented by cuprates. The decomposition of copper oxide (II) occurs at a temperature of about 1100 o C. Ammonia, carbon monoxide, hydrogen and coal are able to extract metallic copper from this compound.

Receipt

It can be obtained by heating metallic copper in air under one condition - the heating temperature must be below 1100 o C. Also, copper (II) oxide can be obtained by heating carbonate, nitrate, divalent copper hydroxide.

Application

With the help of this oxide, enamel and glass are colored green or blue, and a copper-ruby variety of the latter is also produced. In the laboratory, this oxide is used to discover the reducing properties of substances.

Copper(III) oxide

Its formula is Cu 2 O 3. It has a traditional name, which probably sounds a little unusual - copper oxide.

Properties

It has the appearance of red crystals that do not dissolve in water. The decomposition of this substance occurs at a temperature of 400 ° C, the products of this reaction are copper (II) oxide and oxygen.

Receipt

It can be obtained by oxidizing divalent copper hydroxide with potassium peroxydisulphate. A necessary condition for the reaction is an alkaline environment in which it must occur.

Application

This substance is not used by itself. In science and industry, more wide use find the products of its decomposition - copper (II) oxide and oxygen.

Conclusion

That's all copper oxides. There are several of them due to the fact that copper has a variable valence. There are other elements that have several oxides, but we'll talk about them another time.

Chemical properties copper(II) oxide

Brief description of copper oxide (II):

copper oxide(II) – black inorganic substance.

2. reaction of copper (II) oxide with carbon:

CuO + C → Cu + CO (t = 1200 o C).

carbon.

3.copper oxide reaction(II) with gray:

CuO + 2S → Cu + S 2 O (t = 150-200 o C).

The reaction takes place in a vacuum. As a result of the reaction, copper and oxide are formed sulfur.

4. copper oxide reaction(II) with aluminum:

3CuO + 2Al → 3Cu + Al 2 O 3 (t = 1000-1100 o C).

As a result of the reaction, copper and oxide are formed aluminum.

5.copper oxide reaction(II) with copper:

CuO + Cu → Cu 2 O (t = 1000-1200 o C).

As a result of the reaction, copper (I) oxide is formed.

6. copper oxide reaction(II) With lithium oxide:

CuO + Li 2 O → Li 2 CuO 2 (t = 800-1000 o C, O 2).

The reaction takes place in a flow of oxygen. As a result of the reaction, lithium cuprate is formed.

7. copper oxide reaction(II) with sodium oxide:

CuO + Na 2 O → Na 2 CuO 2 (t = 800-1000 o C, O 2).

The reaction takes place in a flow of oxygen. As a result of the reaction, sodium cuprate is formed.

8.copper oxide reaction(II) with carbon monoxide:

CuO + CO → Cu + CO 2.

As a result of the reaction, copper and carbon monoxide (carbon dioxide) are formed.

9. copper oxide reaction(II) with oxide gland:

CuO + Fe 2 O 3 → CuFe 2 O 4 (t o).

As a result of the reaction, a salt is formed - copper ferrite. The reaction proceeds when the reaction mixture is calcined.

10. copper oxide reaction(II) with hydrofluoric acid:

CuO + 2HF → CuF 2 + H 2 O.

As a result of a chemical reaction, a salt is obtained - copper fluoride and water.

11.copper oxide reaction(II) with nitric acid:

CuO + 2HNO 3 → 2Cu(NO 3) 2 + H 2 O.

As a result of a chemical reaction, a salt is obtained - copper nitrate and water .

Copper oxide reacts similarly(II) and with other acids.

12. copper oxide reaction(II) with hydrogen bromide (hydrogen bromide):

CuO + 2HBr → CuBr 2 + H 2 O.

As a result of a chemical reaction, a salt is obtained - copper bromide and water .

13. copper oxide reaction(II) with hydrogen iodine:

CuO + 2HI → CuI 2 + H 2 O.

As a result of a chemical reaction, a salt is obtained - copper iodide and water .

14. copper oxide reaction(II) With sodium hydroxide :

CuO + 2NaOH → Na 2 CuO 2 + H 2 O.

As a result of a chemical reaction, a salt is obtained - sodium cuprate and water .

15.copper oxide reaction(II) With potassium hydroxide :

CuO + 2KOH → K 2 CuO 2 + H 2 O.

As a result of a chemical reaction, a salt is obtained - potassium cuprate and water .

16.copper oxide reaction(II) with sodium hydroxide and water:

CuO + 2NaOH + H 2 O → Na 2 2 (t = 100 o C).

Sodium hydroxide is dissolved in water. A solution of sodium hydroxide in water 20-30%. The reaction proceeds at boiling. As a result of a chemical reaction, sodium tetrahydroxocuprate is obtained.

17.copper oxide reaction(II) with potassium superoxide:

2CuO + 2KO 2 → 2KCuO 2 + O 2 (t = 400-500 o C).

As a result of a chemical reaction, a salt is obtained - potassium cuprate (III) and

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

Request

Copper (Cu) belongs to the d-elements and is located in the IB group of the periodic table of D.I. Mendeleev. Electronic configuration 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 oxidation state +1 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 runs at room temperature.

with halogens

Copper reacts with fluorine, chlorine and bromine to form halides with 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 is passivated under normal conditions with 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, chemical element, which is in group VIIIB and has 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 environment, 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 USE assignments, as a question on the topic to what degree of oxidation iron will be oxidized under the action of dilute and concentrated hydrochloric acid on it. 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

Please note that diluted sulphuric 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..