Properties of simple substances of metals and non-metals USE. Chemical properties of simple substances of metals and non-metals

METALS, THEIR PROPERTIES, OBTAINING, APPLICATION. ELECTROLYSIS.

1. Does not react with water:

1) magnesium 2) beryllium 3) barium 4) strontium

2. The reaction of dilute nitric acid with copper corresponds to the equation:

1) 3 Cu + 8 HNO 3 \u003d 3 Cu (NO 3) 2 + 2 NO + 4 H 2 O

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

3) Cu + 2 HNO 3 = CuO + NO 2 + H 2 O

4) Cu + HNO 3 = CuO + NH 4 NO 3 + H 2 O

3. Compare the processes occurring on the electrodes during the electrolysis of the melt and sodium chloride solution.

4. During the electrolysis of AgNO solution 3 on the cathode is released:

1) silver 2) hydrogen 3) silver and hydrogen 4) oxygen and hydrogen

5. During the electrolysis of a solution of potassium chloride on the cathode, the following occurs:

1) water reduction 2) water oxidation

3) reduction of potassium ions 4) oxidation of chlorine

6. What process occurs on a copper anode during the electrolysis of a sodium bromide solution?

1) water oxidation 2) bromine ion oxidation

3) copper oxidation 4) copper recovery

7. Reaction is possible between:

1) Ag and K 2 SO 4 (solution) 2) Zn and KCl (solution)

3) Mg and SnCl 2 (solution) 4) Ag and CuSO 4 (solution)

8. In what sequence are these metals reduced during the electrolysis of solutions of their salts?

1) Au, Cu, Ag, Fe 2) Cu, Ag, Fe, Au

3) Fe, Cu, Ag, Au 4) Au, Ag, Cu, Fe

9. With concentrated HNO 3 without heating does not interact:

1) Cu 2) Ag 3) Zn 4) Fe

10. Nitric acid accumulates in the electrolyzer when an electric current is passed through an aqueous solution

1) calcium nitrate 2) silver nitrate 3) aluminum nitrate 4) cesium nitrate

11. Of the metals below, the most active is:

1) beryllium 2) magnesium 3) calcium 4) barium

12. Iron reacts with each of two substances:

1) sodium chloride and nitrogen 2) oxygen and chlorine

3) aluminum oxide and potassium carbonate 4) water and aluminum hydroxide

13. Each of the two metals reacts with water at room temperature:

1) barium and copper 2) aluminum and mercury 3) calcium and lithium 4) silver and sodium

14. When aluminum is fused with sodium hydroxide, the following is formed:

1) NaAlO 2 2) AlH 3 3) Na 4) Al 2 O 3

15. With diluted HNO 3 without heating does not interact:

1) Cu 2) Ag 3) Zn 4) Pt

16. Hydrogen is not displaced from acids:

1) chromium 2) iron 3) copper 4) zinc

17. Copper dissolves in a dilute aqueous acid solution:

1) sulfuric 2) hydrochloric 3) nitrogen 4) hydrofluoric

18. Copper products that are in contact with air are gradually covered with a green coating, the main component

The component of which is:

1) CuO 2) CuCO 3 3) Cu(OH) 2 4) (CuOH) 2 CO 3

19. When heating magnesium in a nitrogen atmosphere:

1) the reaction does not go on 2) magnesium nitride is formed

3) magnesium nitrite is formed 4) magnesium nitrate is formed

20. At ordinary temperature, magnesiumdoes not interact With:

A) water

B) alkali solutions

C) dilute H 2 SO 4 and HNO 3

D) concentrated H 2 SO 4 and HNO 3

D) gray

Answer:

21. At room temperature, chromium interacts with:

A) HCl (diff.) B) H 2 O C) H 2 SO 4 (diff.) D) N 2 E) H 2

Answer: ____________________ . (Write down the corresponding letters in alphabetical order.)

22. During the electrolysis of an aqueous solution of KI not generated:

1) K 2) KOH 3) H 2 4) I 2

23. The substance from which the same products are formed during the electrolysis of an aqueous solution and a melt has

Formula:

1) CuCl 2 2) KBr 3) NaOH 4) NaCl

24. Gaseous substances will be released on the cathode and anode during the electrolysis of an aqueous solution:

1) AgNO 3 2) KNO 3 3) CuCl 2 4) HgCl 2

25. During the electrolysis of a solution of Cr 2 (SO 4 ) 3 on the cathode is released:

1) oxygen 2) hydrogen and chromium 3) chromium 4) oxygen and chromium 26. Two inert electrodes were lowered into a glass containing a mixture of aqueous solutions of salts with the same concentration

Cium AgNO 3 , Cu(NO 3 ) 2 , Hg(NO 3 ) 2 , NaNO 3 . The first particles to be reduced during electrolysis are:

1) Hg +2 2) Ag + 3) Cu +2 4) H 2 O

27. During the electrolysis of a dilute aqueous solution of Ni (NO 3 ) 2 on the cathode is released:

1) Ni 2) O 2 3) Ni and H 2 4) H 2 and O 2

28. Nitric acid accumulates in an electrolytic cell when an electric current is passed through an aqueous solution.

1) potassium nitrate 2) aluminum nitrate 3) magnesium nitrate 4) copper nitrate

29. The release of oxygen occurs during the electrolysis of an aqueous solution of salt:

30. During the electrolysis of an aqueous solution of silver nitrate on the cathode, the following is formed:

1) Ag 2) NO 2 3) NO 4) H 2

31. Calcium in industry is obtained by:

1) electrolysis of CaCl solution 2 2) electrolysis of CaCl melt 2

3) electrolysis of Ca(OH) solution 2 4) the action of a more active metal on aqueous solutions of salts

32. During the electrolysis of a solution of sodium iodide at the cathode, the color of litmus in solution:

1) red 2) blue 3) purple 4) yellow

33. During the electrolysis of an aqueous solution of potassium nitrate, the following is released at the anode:

1) O 2 2) NO 2 3) N 2 4) H 2

34. Hydrogen is formed during the electrolysis of an aqueous solution:

1) CaCl 2 2) CuSO 4 3) Hg(NO 3 ) 2 4) AgNO 3

35. When lithium interacts with water, hydrogen is formed and:

1) oxide 2) peroxide 3) hydride 4) hydroxide

36. Metallic properties are weakest expressed in:

1) sodium 2) magnesium 3) calcium 4) aluminum

37. Are the following judgments about alkali metals correct?

A. In all compounds, they have an oxidation state of +1.

B. With non-metals, they form compounds with ionic bonds.

1) only A is true 2) only B is true

3) both judgments are true 4) both judgments are wrong

38. At room temperature, chromium interacts with:

1) H 2 SO 4 (solution) 2) H 2 O 3) N 2 4) O 2

39. When chromium interacts with hydrochloric acid, the following are formed:

1) CrCl 2 and H 2 2) CrCl 3 and H 2 O 3) CrCl 2 and H 2 O 4) CrCl 3 and H 2

40. Copper does not interact With:

1) dilute HNO 3 2) concentrated HNO 3

3) dilute HCl 4) concentrated H 2 SO 4

41. Which of the metals does not displace hydrogen from dilute sulfuric acid?

1) iron 2) chromium 3) copper 4) zinc

42. Reacts most vigorously with water:

1) Al 2) Mg 3) Ca 4) K

43. Under normal conditions, it reacts with water:

1) Mg 2) Ca 3) Pb 4) Zn

44. As a result of the reaction of calcium with water, the following are formed:

1) CaO and H 2 2) Ca (OH) 2 and H 2 3) CaH 2 and O 2 4) Ca (OH) 2 and O 2

45. Chemical reaction does not happen between:

1) Zn and HCl 2) Al and HCl 3) Mg and H 2 SO 4 (diff.) 4) Ag and H 2 SO 4 (diff.)

46. ​​Hydrochloric acid reacts with:

1) Cu 2) Zn 3) Ag 4) Hg

47. For aluminum, under normal conditions, interaction with:

A) HgCl 2 B) CaO C) CuSO 4 D) HNO 3 (conc.) E) Na 2 SO 4 E) Fe 3 O 4

Answer: ____________________ . (Write down the corresponding letters in alphabetical order.)

48. Establish a correspondence between the starting materials and products of redox reactions.

STARTING SUBSTANCES REACTION PRODUCTS

1) Fe + Cl 2 → A) FeSO 4 + H 2

2) Fe + HCl → B) Fe 2 (SO 4) 3 + H 2

3) Fe + H 2 SO 4 (diff.) → B) Fe 2 (SO 4) 3 + SO 2 + H 2 O

4) Fe + H 2 SO 4 (conc.) → D) FeCl 2 + H 2

E) FeCl 3 + H 2

E) FeCl 3

49. Write the equations for the reactions taking place on the cathode and anode, and the general equation for the electrolysis of water

A solution of copper(II) sulfate on inert electrodes.

50. Write the equations for the reactions taking place on the cathode and anode, and the general equation for the electrolysis of an aqueous solution

Barium chloride on inert electrodes.

51. Write the equations for the reactions taking place on the cathode and anode, and the general equation for the electrolysis of an aqueous solution

Potassium iodide on inert electrodes.

52. Write the equations for the reactions taking place on the cathode and anode, and the general equation for the electrolysis of an aqueous solution

Sulfuric acid on inert electrodes.

53. Write the equations for the reactions taking place on the cathode and anode, and the general equation for the electrolysis of an aqueous solution

Lithium bromide on inert electrodes.

54. Under normal conditions, calcium reacts with:

1) oxygen 2) carbon 3) sulfur 4) nitrogen

55. Write the equations for the reactions taking place on the cathode and anode, and the general equation for the electrolysis of an aqueous solution

Potassium nitrate on inert electrodes.

56. Write the equations for the reactions taking place on the cathode and anode, and the general equation for the electrolysis of an aqueous solution

Sodium sulfate on inert electrodes.

57. At ordinary temperature, copper reacts with:

1) water 2) oxygen 3) hydrochloric acid 4) nitric acid

58. Write the equations for the reactions taking place on the cathode and anode, and the general equation for the electrolysis of an aqueous solution

Potassium hydroxide on inert electrodes.

59. Dissolves in dilute sulfuric acid:

1) Cu 2) Zn 3) Ag 4) Au

60. Write the equations for the reactions taking place on the cathode and anode, and the general equation for the electrolysis of an aqueous solution

Nitric acid on inert electrodes.

61. When heated, copper reacts with:

1) hydrogen 2) hydrosulphuric acid

Video lesson 1: Inorganic chemistry. Metals: alkali, alkaline earth, aluminum

Video lesson 2: transition metals

Lecture: Characteristic chemical properties and production of simple substances - metals: alkali, alkaline earth, aluminum; transition elements (copper, zinc, chromium, iron)

Chemical properties of metals

All metals in chemical reactions manifest themselves as reducing agents. They easily part with valence electrons, being oxidized at the same time. Recall that the further to the left a metal is located in the electrochemical series of tension, the stronger the reducing agent it is. Therefore, the strongest is lithium, the weakest is gold and vice versa, gold is the strongest oxidizing agent, and lithium is the weakest.

Li→Rb→K→Ba→Sr→Ca→Na→Mg→Al→Mn→Cr→Zn→Fe→Cd→Co→Ni→Sn→Pb→H→Sb→Bi→Cu→Hg→Ag→Pd→ Pt→Au

All metals displace other metals from the salt solution, i.e. restore them. All except alkaline and alkaline earth as they interact with water. Metals located before H displace it from solutions of dilute acids, and they themselves dissolve in them.

Consider some general chemical properties of metals:

• The interaction of metals with oxygen forms basic (CaO, Na 2 O, 2Li 2 O, etc.) or amphoteric (ZnO, Cr 2 O 3, Fe 2 O 3, etc.) oxides.
• The interaction of metals with halogens (the main subgroup of group VII) forms hydrohalic acids (HF - hydrogen fluoride, HCl - hydrogen chloride, etc.).
• The interaction of metals with non-metals forms salts (chlorides, sulfides, nitrides, etc.).
• The interaction of metals with metals forms intermetallic compounds (MgB 2 , NaSn, Fe 3 Ni, etc.).
• The interaction of active metals with hydrogen forms hydrides (NaH, CaH 2, KH, etc.).
• The interaction of alkali and alkaline earth metals with water forms alkalis (NaOH, Ca (OH) 2, Cu (OH) 2, etc.).
• The interaction of metals (only those standing in the electrochemical series up to H) with acids forms salts (sulfates, nitrites, phosphates, etc.). It should be borne in mind that metals react with acids quite reluctantly, while they almost always interact with bases and salts. In order for the reaction of the metal with the acid to take place, the metal must be active and the acid strong.

Chemical properties of alkali metals

The group of alkali metals includes the following chemical elements: lithium (Li), sodium (Na), potassium (K), rubidium (Rb), cesium (Cs), francium (Fr). As they move from top to bottom in Group I of the Periodic Table, their atomic radii increase, which means that their metallic and reducing properties increase.

Consider the chemical properties of alkali metals:

• They do not have signs of amphotericity, as they have negative values ​​of electrode potentials.
• The strongest reducing agents among all metals.
• In compounds, they exhibit only the +1 oxidation state.
• Giving a single valence electron, the atoms of these chemical elements are converted into cations.
• They form numerous ionic compounds.
• Almost all are soluble in water.

Interaction of alkali metals with other elements:

1. With oxygen, forming individual compounds, so the oxide forms only lithium (Li 2 O), sodium forms peroxide (Na 2 O 2), and potassium, rubidium and cesium form superoxides (KO 2, RbO 2, CsO 2).

2. With water, forming alkalis and hydrogen. Remember, these reactions are explosive. Without an explosion, only lithium reacts with water:

2Li + 2H 2 O → 2LiO H + H 2.

3. With halogens, forming halides (NaCl - sodium chloride, NaBr - sodium bromide, NaI - sodium iodide, etc.).

4. With hydrogen when heated, forming hydrides (LiH, NaH, etc.)

5. With sulfur when heated, forming sulfides (Na 2 S, K 2 S, etc.). They are colorless and highly soluble in water.

6. With phosphorus when heated, forming phosphides (Na 3 P, Li 3 P, etc.), they are very sensitive to moisture and air.

7. With carbon, when heated, carbides form only lithium and sodium (Li 2 CO 3, Na 2 CO 3), while potassium, rubidium and cesium do not form carbides, they form binary compounds with graphite (C 8 Rb, C 8 Cs, etc.) .

8. Under normal conditions, only lithium reacts with nitrogen, forming Li 3 N nitride, with other alkali metals, the reaction is possible only when heated.

9. They react explosively with acids, so carrying out such reactions is very dangerous. These reactions are ambiguous, because the alkali metal actively reacts with water, forming an alkali, which is then neutralized by an acid. This creates competition between alkali and acid.

10. With ammonia, forming amides - analogues of hydroxides, but stronger bases (NaNH 2 - sodium amide, KNH 2 - potassium amide, etc.).

11. With alcohols, forming alcoholates.

Francium is a radioactive alkali metal, one of the rarest and least stable of all radioactive elements. Its chemical properties are not well understood.

To obtain alkali metals, they mainly use the electrolysis of melts of their halides, most often chlorides, which form natural minerals:

• NaCl → 2Na + Cl 2 .

• Na 2 CO 3 + 2C → 2Na + 3CO.

• 2Li 2 O + Si + 2CaO → 4Li + Ca 2 SiO 4 .

• KCl + Na → K + NaCl.

Chemical properties of alkaline earth metals

Alkaline earth metals include elements of the main subgroup of group II: calcium (Ca), strontium (Sr), barium (Ba), radium (Ra). The chemical activity of these elements grows in the same way as that of alkali metals, i.e. increasing down the subgroup.

Chemical properties of alkaline earth metals:

The structure of the valence shells of atoms of these elements ns 2 .

• Giving two valence electrons, the atoms of these chemical elements are converted into cations.
• The compounds exhibit an oxidation state of +2.
• The charges of atomic nuclei are greater by one than those of alkaline elements of the same periods, which leads to a decrease in the radius of atoms and an increase in ionization potentials.

Interaction of alkaline earth metals with other elements:

1. With oxygen, all alkaline earth metals, except for barium, form oxides, barium forms peroxide BaO 2. Of these metals, beryllium and magnesium, coated with a thin protective oxide film, interact with oxygen only at very high t. Basic oxides of alkaline earth metals react with water, with the exception of beryllium oxide BeO, which has amphoteric properties. The reaction of calcium oxide and water is called the lime slaking reaction. If the reagent is CaO, quicklime is formed, if Ca(OH) 2, slaked. Also, basic oxides react with acidic oxides and acids. Eg:

• 3CaO + P 2 O 5 → Ca 3 (PO 4) 2 .

2. With water, alkaline earth metals and their oxides form hydroxides - white crystalline substances, which, in comparison with alkali metal hydroxides, are less soluble in water. Hydroxides of alkaline earth metals are alkalis, except for the amphoteric Be(OH ) 2 and weak base Mg(OH)2. Since beryllium does not react with water, Be (OH ) 2 can be obtained in other ways, for example, by hydrolysis of nitride:

• Be 3 N 2+ 6H 2 O → 3 Be (OH)2+ 2N N 3.

3. Under normal conditions, everything reacts with halogens, except for beryllium. The latter reacts only at high t. Halides are formed (MgI 2 - magnesium iodide, CaI 2 - calcium iodide, CaBr 2 - calcium bromide, etc.).

4. All alkaline earth metals, except beryllium, react with hydrogen when heated. Hydrides are formed (BaH 2 , CaH 2 , etc.). For the reaction of magnesium with hydrogen, in addition to high t, an increased hydrogen pressure is also required.

5. Sulfur forms sulfides. Eg:

• Ca + S → CaS.

Sulfides are used to obtain sulfuric acid and the corresponding metals.

6. They form nitrides with nitrogen. Eg:

• 3Be + N 2 → Be 3 N 2.

7. With acids, forming salts of the corresponding acid and hydrogen. Eg:

• Be + H 2 SO 4 (razb.) → BeSO 4 + H 2.

These reactions proceed in the same way as in the case of alkali metals.

Obtaining alkaline earth metals:

• BeF 2 + Mg –t o → Be + MgF 2

• 3BaO + 2Al –t o → 3Ba + Al 2 O 3

• CaCl 2 → Ca + Cl 2

Chemical properties of aluminum

Aluminum is an active, light metal, number 13 in the table. In nature, the most common of all metals. And of the chemical elements, it occupies the third position in terms of distribution. High heat and electrical conductor. Resistant to corrosion, as it is covered with an oxide film. The melting point is 660 0 С.

Consider the chemical properties and interaction of aluminum with other elements:

1. In all compounds, aluminum is in the +3 oxidation state.

2. It exhibits reducing properties in almost all reactions.

3. Amphoteric metal exhibits both acidic and basic properties.

4. Restores many metals from oxides. This method of obtaining metals is called aluminothermy. Example of getting chromium:

2Al + Cr 2 O 3 → Al 2 O 3 + 2Cr.

5. Reacts with all dilute acids to form salts and release hydrogen. Eg:

2Al + 6HCl → 2AlCl 3 + 3H 2;

2Al + 3H2SO4 → Al 2 (SO 4) 3 + 3H 2.

In concentrated HNO 3 and H 2 SO 4 aluminum is passivated. Thanks to this, it is possible to store and transport these acids in containers made of aluminum.

6. Interacts with alkalis, as they dissolve the oxide film.

7. Reacts with all non-metals except hydrogen. To carry out the reaction with oxygen, finely divided aluminum is needed. The reaction is only possible at high t:

• 4Al + 3O 2 → 2Al 2 O 3 .

According to its thermal effect, this reaction is exothermic. Interaction with sulfur forms aluminum sulfide Al 2 S 3 , with phosphorus phosphide AlP, with nitrogen nitride AlN, with carbon carbide Al 4 C 3 .

8. It interacts with other metals, forming aluminides (FeAl 3 CuAl 2, CrAl 7, etc.).

Metallic aluminum is obtained by electrolysis of a solution of alumina Al 2 O 3 in molten cryolite Na 2 AlF 6 at 960–970°C.

• 2Al2O3 → 4Al + 3O 2 .
  

Chemical properties of transition elements

Transitional elements include elements of secondary subgroups of the Periodic Table. Consider the chemical properties of copper, zinc, chromium and iron.

Chemical properties of copper

1. In the electrochemical series, it is located to the right of H, so this metal is inactive.

2. Weak reducer.

3. In compounds, it exhibits oxidation states +1 and +2.

4. Reacts with oxygen when heated to form:

• copper oxide (I) 2Cu + O 2 → 2CuO(at t 400 0 C)
• or copper(II) oxide: 4Cu + O2 → 2Cu2O(at t 200 0 C).

Oxides have basic properties. When heated in an inert atmosphere, Cu 2 O disproportionates: Cu2O → CuO + Cu. Copper (II) oxide CuO forms cuprates in reactions with alkalis, for example: CuO + 2NaOH → Na 2 CuO 2 + H 2 O.

5. Copper hydroxide Cu (OH) 2 is amphoteric, the main properties prevail in it. It dissolves easily in acids:

• Cu (OH) 2 + 2HNO 3 → Cu(NO 3) 2 + 2H 2 O,

and in concentrated solutions of alkalis with difficulty:

• Сu(OH) 2 + 2NaOH → Na 2.

6. The interaction of copper with sulfur under various temperature conditions also forms two sulfides. When heated to 300-400 0 C in a vacuum, copper (I) sulfide is formed:

• 2Cu+S → Cu2S.

At room temperature, by dissolving sulfur in hydrogen sulfide, copper (II) sulfide can be obtained:

• Cu+S → CuS.

7. Of the halogens, it interacts with fluorine, chlorine and bromine, forming halides (CuF 2 , CuCl 2 , CuBr 2), iodine, forming copper (I) iodide CuI; does not interact with hydrogen, nitrogen, carbon, silicon.

8. It does not react with acids - non-oxidizing agents, because they oxidize only metals located to hydrogen in the electrochemical series. This chemical element reacts with oxidizing acids: dilute and concentrated nitric and concentrated sulfuric:

3Cu + 8HNO 3 (diff) → 3Cu(NO 3) 2 + 2NO + 4H 2 O;

Cu + 4HNO 3 (conc) → Cu(NO 3) 2 + 2NO 2 + 2H 2 O;

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

9. Interacting with salts, copper displaces from their composition the metals located to the right of it in the electrochemical series. Eg,

2FeCl 3 + Cu → CuCl 2 + 2FeCl 2 .

Here we see that copper went into solution, and iron (III) was reduced to iron (II). This reaction is of great practical importance and is used to remove copper deposited on plastic.

Chemical properties of zinc

2. It has pronounced reducing properties and amphoteric properties.

3. In compounds, it exhibits an oxidation state of +2.

4. In air, it is covered with an oxide film of ZnO.

5. Interaction with water is possible at a temperature of red heat. As a result, zinc oxide and hydrogen are formed:

• Zn + H 2 O → ZnO + H 2.

6. Interacts with halogens, forming halides (ZnF 2 - zinc fluoride, ZnBr 2 - zinc bromide, ZnI 2 - zinc iodide, ZnCl 2 - zinc chloride).

7. With phosphorus it forms the phosphides Zn 3 P 2 and ZnP 2 .

8. With sulfur chalcogenide ZnS.

9. Does not directly react with hydrogen, nitrogen, carbon, silicon and boron.

10. It interacts with non-oxidizing acids, forming salts and displacing hydrogen. Eg:

• H 2 SO 4 + Zn → ZnSO 4 + H 2
• Zn + 2HCl → ZnCl 2 + H 2 .

It also reacts with acids - oxidizing agents: with conc. sulfuric acid forms zinc sulfate and sulfur dioxide:

• Zn + 2H 2 SO 4 → ZnSO 4 + SO 2 + 2H 2 O.

11. It actively reacts with alkalis, since zinc is an amphoteric metal. With alkali solutions, it forms tetrahydroxozincates and releases hydrogen:

• Zn + 2NaOH + 2H 2 O → Na 2 + H 2 .

Gas bubbles appear on the zinc granules after the reaction. With anhydrous alkalis, when fused, it forms zincates and releases hydrogen:

• Zn+ 2NaOH → Na 2 ZnO 2 + H 2.

Chemical properties of chromium

2.

3. Forms colored compounds.

4. In compounds, it exhibits oxidation states +2 (basic oxide CrO black), +3 (amphoteric oxide Cr 2 O 3 and hydroxide Cr (OH) 3 green) and +6 (acid oxide chromium (VI) CrO 3 and acids: chromic H 2 CrO 4 and two-chrome H 2 Cr 2 O 7, etc.).

5. It interacts with fluorine at t 350-400 0 C, forming chromium (IV) fluoride:

• Cr+2F 2 → CrF 4 .

6. With oxygen, nitrogen, boron, silicon, sulfur, phosphorus and halogens at t 600 0 C:

• connection with oxygen forms chromium oxide (VI) CrO 3 (dark red crystals),
• nitrogen compound - chromium nitride CrN (black crystals),
• compound with boron - chromium boride CrB (yellow crystals),
• compound with silicon - chromium silicide CrSi,
• connection with carbon - chromium carbide Cr 3 C 2 .

7. It reacts with water vapor, being in a hot state, forming chromium (III) oxide and hydrogen:

• 2Cr + 3H 2 O → Cr 2 O 3 + 3H 2 .

8. It does not react with alkali solutions, but slowly reacts with their melts, forming chromates:

• 2Cr + 6KOH → 2KCrO 2 + 2K 2 O + 3H 2 .

9. It dissolves in dilute strong acids to form salts. If the reaction takes place in air, Cr 3+ salts are formed, for example:

• 2Cr + 6HCl + O 2 → 2CrCl 3 + 2H 2 O + H 2 .

• Cr + 2HCl → CrCl 2 + H 2 .

10. With concentrated sulfuric and nitric acids, as well as with aqua regia, it reacts only when heated, because. at low temperatures, these acids passivate chromium. Reactions with acids when heated look like this:

2Cr + 6H 2 SO 4 (conc) → Cr 2 (SO 4) 3 + 3SO 2 + 6H 2 O

Cr + 6HNO 3 (conc) → Cr(NO 3) 3 + 3NO 2 + 3H 2 O

Chromium(II) oxide CrO- solid black or red, insoluble in water.

Chemical properties:

• It has basic and restorative properties.
• When heated to 100 0 C in air, it oxidizes to Cr 2 O 3 - chromium (III) oxide.
• It is possible to restore chromium with hydrogen from this oxide: CrO + H 2 → Cr + H 2 O or coke: CrO + C → Cr + CO.
• Reacts with hydrochloric acid, while releasing hydrogen: 2CrO + 6HCl → 2CrCl 3 + H 2 + 2H 2 O.
• Does not react with alkalis, dilute sulfuric and nitric acids.

Chromium oxide (III) Cr 2 O 3- a refractory substance, dark green in color, insoluble in water.

Chemical properties:

• It has amphoteric properties.
• How basic oxide interacts with acids: Cr 2 O 3 + 6HCl → CrCl 3 + 3H 2 O.
• How acidic oxide interacts with alkalis: Cr 2 O 3 + 2KOH → 2KCrO 3 + H 2 O.
• Strong oxidizing agents oxidize Cr 2 O 3 to chromate H 2 CrO 4 .
• Strong reducing agents restoreCr out Cr2O3.

Chromium(II) hydroxide Cr(OH) 2 - solid yellow or brown color, poorly soluble in water.

Chemical properties:

• Weak base, exhibits basic properties.
• In the presence of moisture in air, it oxidizes to Cr(OH) 3 - chromium (III) hydroxide.
• Reacts with concentrated acids to form blue chromium (II) salts: Cr(OH) 2 + H 2 SO 4 → CrSO 4 + 2H 2 O.
• Does not react with alkalis and dilute acids.

Chromium (III) hydroxide Cr(OH) 3 - a gray-green substance, insoluble in water.

Chemical properties:

• It has amphoteric properties.
• How basic hydroxide interacts with acids: Cr(OH) 3 + 3HCl → CrCl 3 + 3H 2 O.
• How acid hydroxide interacts with alkalis: Cr(OH) 3 + 3NaOH → Na 3 [Cr(OH)6].

Chemical properties of iron

1. Active metal with high reactivity.

2. It has restorative properties, as well as pronounced magnetic properties.

3. In compounds, it exhibits the main oxidation states +2 (with weak oxidizing agents: S, I, HCl, salt solutions), +3 (with strong oxidizing agents: Br and Cl) and less characteristic +6 (with O and H 2 O). In weak oxidizing agents, iron takes the oxidation state +2, in stronger ones +3. +2 oxidation states correspond to black oxide FeO and green hydroxide Fe (OH) 2, which have basic properties. +3 oxidation states correspond to red-brown oxide Fe 2 O 3 and brown hydroxide Fe (OH) 3, which have weakly pronounced amphoteric properties. Fe (+2) is a weak reducing agent, and Fe (+3) is often a weak oxidizing agent. When the redox conditions change, the oxidation states of iron can change with each other.

• 3Fe + 2O 2 → Fe 3 O 4.

5. Reacts with halogens when heated:

• connection with chlorine forms iron (III) chloride FeCl 3,
• compound with bromine - iron (III) bromide FeBr 3,
• compound with iodine - iron (II,III) iodide Fe 3 I 8,
• compound with fluorine - iron (II) fluoride FeF 2, iron (III) fluoride FeF 3.

• connection with sulfur forms iron(II) sulfide FeS,
• connection with nitrogen - iron nitride Fe 3 N,
• compound with phosphorus - phosphides FeP, Fe 2 P and Fe 3 P,
• compound with silicon - iron silicide FeSi,
• compound with carbon - iron carbide Fe 3 C.

9. It does not react with alkali solutions, but slowly reacts with alkali melts, which are strong oxidizing agents:

• Fe + KClO 3 + 2KOH → K 2 FeO 4 + KCl + H 2 O.

10. Restores metals located in the electrochemical row to the right:

• Fe + SnCl 2 → FeCl 2 + Sn.

• 3Fe2O3 + CO → CO 2 + 2Fe 3 O 4,
• Fe 3 O 4 + CO → CO 2 + 3FeO,
• FeO + CO → CO 2 + Fe.

Iron(II) oxide FeO - a black crystalline substance (wustite) that does not dissolve in water.

Chemical properties:

• Has basic properties.
• Reacts with dilute hydrochloric acid: FeO + 2HCl → FeCl 2 + H 2 O.
• Reacts with concentrated nitric acid:FeO + 4HNO 3 → Fe(NO 3) 3 + NO 2 + 2H 2 O.
• Does not react with water and salts.
• With hydrogen at t 350 0 C it is reduced to pure metal: FeO + H 2 → Fe + H 2 O.
• It is also reduced to pure metal when combined with coke: FeO + C → Fe + CO.
• This oxide can be obtained in various ways, one of them is heating Fe at low pressure O: 2Fe + O 2 → 2FeO.

Iron(III) oxideFe2O3- brown powder (hematite), a substance insoluble in water. Other names: iron oxide, iron minium, food coloring E172, etc.

Chemical properties:

• Fe 2 O 3 + 6HCl → 2 FeCl 3 + 3H 2 O.
• It does not react with alkali solutions, it reacts with their melts, forming ferrites: Fe 2 O 3 + 2NaOH → 2NaFeO 2 + H 2 O.
• When heated with hydrogen, it exhibits oxidizing properties:Fe 2 O 3 + H 2 → 2FeO + H 2 O.
• Fe 2 O 3 + 3KNO 3 + 4KOH → 2K 2 FeO 4 + 3KNO 2 + 2H 2 O.

Iron oxide (II, III) Fe 3 O 4 or FeO Fe 2 O 3 - a grayish-black solid (magnetite, magnetic iron ore), a substance insoluble in water.

Chemical properties:

• Decomposes when heated above 1500 0 С: 2Fe 3 O 4 → 6FeO + O 2.
• Reacts with dilute acids: Fe 3 O 4 + 8HCl → FeCl 2 + 2FeCl 3 + 4H 2 O.
• Does not react with alkali solutions, reacts with their melts: Fe 3 O 4 + 14NaOH → Na 3 FeO 3 + 2Na 5 FeO 4 + 7H 2 O.
• When reacting with oxygen, it oxidizes: 4Fe 3 O 4 + O 2 → 6Fe 2 O 3.
• With hydrogen, when heated, it is restored:Fe 3 O 4 + 4H 2 → 3Fe + 4H 2 O.
• It is also reduced when combined with carbon monoxide: Fe 3 O 4 + 4CO → 3Fe + 4CO 2.

Iron(II) hydroxide Fe(OH) 2 - white, rarely greenish crystalline substance, insoluble in water.

Chemical properties:

• It has amphoteric properties with a predominance of basic ones.
• It enters into the neutralization reaction of the non-oxidizing acid, showing the main properties: Fe(OH) 2 + 2HCl → FeCl 2 + 2H 2 O.
• When interacting with nitric or concentrated sulfuric acids, it exhibits reducing properties, forming iron (III) salts: 2Fe(OH) 2 + 4H 2 SO 4 → Fe 2 (SO 4) 3 + SO 2 + 6H 2 O.
• When heated, it reacts with concentrated alkali solutions: Fe (OH) 2 + 2NaOH → Na 2.

Iron hydroxide (I I I) Fe (OH) 3- brown crystalline or amorphous substance, insoluble in water.

Chemical properties:

• It has mild amphoteric properties with a predominance of basic ones.
• Easily interacts with acids: Fe(OH) 3 + 3HCl → FeCl 3 + 3H 2 O.
• With concentrated alkali solutions it forms hexahydroxoferrates (III): Fe (OH) 3 + 3NaOH → Na 3.
• It forms ferrates with alkali melts:2Fe(OH) 3 + Na 2 CO 3 → 2NaFeO 2 + CO 2 + 3H 2 O.
• In an alkaline environment with strong oxidizing agents, it exhibits reducing properties: 2Fe(OH) 3 + 3Br 2 + 10KOH → 2K 2 FeO 4 + 6NaBr + 8H 2 O.

All chemical elements are divided into metals And nonmetals depending on the structure and properties of their atoms. Also, simple substances formed by elements are classified into metals and non-metals, based on their physical and chemical properties.

In the Periodic system of chemical elements D.I. Mendeleev, non-metals are located diagonally: boron - astatine and above it in the main subgroups.

Metal atoms are characterized by relatively large radii and a small number of electrons at the outer level from 1 to 3 (exceptions: germanium, tin, lead - 4; antimony and bismuth - 5; polonium - 6 electrons).

Non-metal atoms, on the contrary, are characterized by small atomic radii and the number of electrons at the outer level from 4 to 8 (the exception is boron, it has three such electrons).

Hence the tendency of metal atoms to give up external electrons, i.e. reducing properties, and for non-metal atoms - the desire to receive missing electrons to a stable eight-electron level, i.e. oxidizing properties.

Metals

In metals, there is a metallic bond and a metallic crystal lattice. At the lattice sites there are positively charged metal ions bound by socialized external electrons belonging to the entire crystal.

This determines all the most important physical properties of metals: metallic luster, electrical and thermal conductivity, plasticity (the ability to change shape under external influence) and some others characteristic of this class of simple substances.

Group I metals of the main subgroup are called alkali metals.

Group II metals: calcium, strontium, barium - alkaline earth.

Chemical properties of metals

In chemical reactions, metals exhibit only reducing properties, i.e. their atoms donate electrons, forming positive ions as a result.

1. Interact with non-metals:

a) oxygen (with the formation of oxides)

Alkali and alkaline earth metals oxidize easily under normal conditions, so they are stored under a layer of vaseline oil or kerosene.

4Li + O 2 = 2Li 2 O

2Ca + O 2 \u003d 2CaO

Please note: when sodium interacts, peroxide is formed, potassium - superoxide

2Na + O 2 \u003d Na 2 O 2, K + O2 \u003d KO2

and oxides are obtained by calcining peroxide with the corresponding metal:

2Na + Na 2 O 2 \u003d 2Na 2 O

Iron, zinc, copper and other less active metals slowly oxidize in air and actively when heated.

3Fe + 2O 2 = Fe 3 O 4 (a mixture of two oxides: FeO and Fe 2 O 3)

2Zn + O 2 = 2ZnO

2Cu + O 2 \u003d 2CuO

Gold and platinum metals are not oxidized by atmospheric oxygen under any circumstances.

b) hydrogen (with the formation of hydrides)

2Na + H2 = 2NaH

Ca + H 2 \u003d CaH 2

c) chlorine (with the formation of chlorides)

2K + Cl 2 \u003d 2KCl

Mg + Cl 2 \u003d MgCl 2

2Al + 3Cl 2 \u003d 2AlCl 3

Please note: when iron reacts, iron (III) chloride is formed:

2Fe + 3Cl 2 = 2FeCl 3

d) sulfur (with the formation of sulfides)

2Na + S = Na 2 S

Hg + S = HgS

2Al + 3S = Al 2 S 3

Please note: when iron reacts, iron (II) sulfide is formed:

Fe + S = FeS

e) nitrogen (with the formation of nitrides)

6K + N 2 = 2K 3 N

3Mg + N 2 \u003d Mg 3 N 2

2Al + N 2 = 2AlN

2. Interact with complex substances:

It must be remembered that, according to the restorative ability, the metals are arranged in a row, which is called the electrochemical series of voltages or activity of metals (Beketov N.N. displacement series):

Li, K, Ba, Ca, Na, Mg, Al, Mn, Zn, Cr, Fe, Co, Ni, Sn, Pb, (H 2), Cu, Hg, Ag, Au, Pt

a) water

Metals located in a row up to magnesium, under normal conditions, displace hydrogen from water, forming soluble bases - alkalis.

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

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

Magnesium interacts with water when boiled.

Mg + 2H 2 O \u003d Mg (OH) 2 + H 2

Aluminum reacts violently with water when the oxide film is removed.

2Al + 6H 2 O \u003d 2Al (OH) 3 + 3H 2

The rest of the metals, standing in a row up to hydrogen, under certain conditions, can also react with water with the release of hydrogen and the formation of oxides.

3Fe + 4H 2 O \u003d Fe 3 O 4 + 4H 2

b) acid solutions

(Except concentrated sulfuric acid and nitric acid of any concentration. See redox reactions.)

Please note: do not use insoluble silicic acid for reactions

Metals ranging from magnesium to hydrogen displace hydrogen from acids.

Mg + 2HCl \u003d MgCl 2 + H 2

Please note: ferrous salts are formed.

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

The formation of an insoluble salt prevents the reaction from proceeding. For example, lead practically does not react with a solution of sulfuric acid due to the formation of insoluble lead sulfate on the surface.

Metals in the row after hydrogen do NOT displace hydrogen.

c) salt solutions

Metals that are in the row up to magnesium and actively react with water are not used to carry out such reactions.

For other metals, the rule is fulfilled:

Each metal displaces from salt solutions other metals located in the row to the right of it, and can itself be displaced by metals located to the left of it.

Cu + HgCl 2 \u003d Hg + CuCl 2

Fe + CuSO 4 \u003d FeSO 4 + Cu

As with acid solutions, the formation of an insoluble salt prevents the reaction from proceeding.

d) alkali solutions

Metals interact, the hydroxides of which are amphoteric.

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

2Al + 2KOH + 6H 2 O = 2K + 3H 2

e) with organic substances

Alkali metals with alcohols and phenol.

2C 2 H 5 OH + 2Na \u003d 2C 2 H 5 ONa + H 2

2C 6 H 5 OH + 2Na \u003d 2C 6 H 5 ONa + H 2

Metals participate in reactions with haloalkanes, which are used to obtain lower cycloalkanes and for syntheses, during which the carbon skeleton of the molecule becomes more complex (A. Wurtz reaction):

CH 2 Cl-CH 2 -CH 2 Cl + Zn = C 3 H 6 (cyclopropane) + ZnCl 2

2CH 2 Cl + 2Na \u003d C 2 H 6 (ethane) + 2NaCl

non-metals

In simple substances, the atoms of non-metals are connected by a covalent non-polar bond. In this case, single (in H 2, F 2, Cl 2, Br 2, I 2), double (in O 2 molecules), triple (in N 2 molecules) covalent bonds are formed.

The structure of simple substances - non-metals:

1. molecular

Under normal conditions, most of these substances are gases (H 2, N 2, O 2, O 3, F 2, Cl 2) or solids (I 2, P 4, S 8) and only a single bromine (Br 2) is liquid. All these substances have a molecular structure, therefore they are volatile. In the solid state, they are fusible due to the weak intermolecular interaction that keeps their molecules in the crystal, and are capable of sublimation.

2. atomic

These substances are formed by crystals, in the nodes of which there are atoms: (B n, C n, Si n, Gen, Se n, Te n). Due to the high strength of covalent bonds, they, as a rule, have high hardness, and any changes associated with the destruction of the covalent bond in their crystals (melting, evaporation) are performed with a large expenditure of energy. Many of these substances have high melting and boiling points, and their volatility is very low.

Many elements - non-metals form several simple substances - allotropic modifications. Allotropy can be associated with a different composition of molecules: oxygen O 2 and ozone O 3 and with different crystal structures: allotropic modifications of carbon are graphite, diamond, carbine, fullerene. Elements - non-metals with allotropic modifications: carbon, silicon, phosphorus, arsenic, oxygen, sulfur, selenium, tellurium.

Chemical properties of non-metals

The atoms of non-metals are dominated by oxidizing properties, that is, the ability to attach electrons. This ability is characterized by the value of electronegativity. Among the non-metals

At, B, Te, H, As, I, Si, P, Se, C, S, Br, Cl, N, O, F

electronegativity increases and oxidizing properties are enhanced.

It follows that for simple substances - non-metals, both oxidizing and reducing properties will be characteristic, with the exception of fluorine, the strongest oxidizing agent.

1. Oxidizing properties

a) in reactions with metals (metals are always reducing agents)

2Na + S = Na 2 S (sodium sulfide)

3Mg + N 2 = Mg 3 N 2 (magnesium nitride)

b) in reactions with non-metals located to the left of this one, that is, with a lower value of electronegativity. For example, when phosphorus and sulfur interact, sulfur will be the oxidizing agent, since phosphorus has a lower electronegativity value:

2P + 5S = P 2 S 5 (phosphorus V sulfide)

Most non-metals will be oxidizing agents in reactions with hydrogen:

H 2 + S = H 2 S

H 2 + Cl 2 \u003d 2HCl

3H 2 + N 2 \u003d 2NH 3

c) in reactions with some complex substances

Oxidizing agent - oxygen, combustion reactions

CH 4 + 2O 2 \u003d CO 2 + 2H 2 O

2SO 2 + O 2 \u003d 2SO 3

Oxidizing agent - chlorine

2FeCl 2 + Cl 2 = 2FeCl 3

2KI + Cl 2 \u003d 2KCl + I 2

CH 4 + Cl 2 \u003d CH 3 Cl + HCl

Ch 2 \u003d CH 2 + Br 2 \u003d CH 2 Br-CH 2 Br

2. Restorative properties

a) in reactions with fluorine

S + 3F 2 = SF 6

H 2 + F 2 \u003d 2HF

Si + 2F 2 = SiF 4

b) in reactions with oxygen (except fluorine)

S + O 2 \u003d SO 2

N 2 + O 2 \u003d 2NO

4P + 5O 2 \u003d 2P 2 O 5

C + O 2 = CO 2

c) in reactions with complex substances - oxidizing agents

H 2 + CuO \u003d Cu + H 2 O

6P + 5KClO 3 \u003d 5KCl + 3P 2 O 5

C + 4HNO 3 \u003d CO 2 + 4NO 2 + 2H 2 O

H 2 C \u003d O + H 2 \u003d CH 3 OH

3. Disproportionation reactions: the same non-metal is both an oxidizing agent and a reducing agent

Cl 2 + H 2 O \u003d HCl + HClO

3Cl 2 + 6KOH \u003d 5KCl + KClO 3 + 3H 2 O

General properties of metals.

The presence of valence electrons weakly bound to the nucleus determines the general chemical properties of metals. In chemical reactions, they always act as a reducing agent; simple substances, metals, never exhibit oxidizing properties.

Getting metals:
- recovery from oxides with carbon (C), carbon monoxide (CO), hydrogen (H2) or more active metal (Al, Ca, Mg);
- recovery from salt solutions with a more active metal;
- electrolysis of solutions or melts of metal compounds - recovery of the most active metals (alkali, alkaline earth metals and aluminum) using electric current.

In nature, metals are found mainly in the form of compounds, only low-active metals are found in the form of simple substances (native metals).

Chemical properties of metals.
1. Interaction with simple substances non-metals:
Most metals can be oxidized with non-metals such as halogens, oxygen, sulfur, nitrogen. But most of these reactions require preheating to start. In the future, the reaction can proceed with the release of a large amount of heat, which leads to the ignition of the metal.
At room temperature, reactions are possible only between the most active metals (alkali and alkaline earth) and the most active non-metals (halogens, oxygen). Alkali metals (Na, K) react with oxygen to form peroxides and superoxides (Na2O2, KO2).

a) interaction of metals with water.
At room temperature, alkali and alkaline earth metals interact with water. As a result of the substitution reaction, an alkali (soluble base) and hydrogen are formed: Metal + H2O \u003d Me (OH) + H2
When heated, other metals interact with water, standing in the activity series to the left of hydrogen. Magnesium reacts with boiling water, aluminum - after a special surface treatment, as a result, insoluble bases are formed - magnesium hydroxide or aluminum hydroxide - and hydrogen is released. Metals in the activity range from zinc (inclusive) to lead (inclusive) interact with water vapor (i.e. above 100 C), while oxides of the corresponding metals and hydrogen are formed.
Metals to the right of hydrogen in the activity series do not interact with water.
b) interaction with oxides:
active metals interact in a substitution reaction with oxides of other metals or non-metals, reducing them to simple substances.
c) interaction with acids:
Metals located to the left of hydrogen in the activity series react with acids to release hydrogen and form the corresponding salt. Metals to the right of hydrogen in the activity series do not interact with acid solutions.
A special place is occupied by the reactions of metals with nitric and concentrated sulfuric acids. All metals except noble ones (gold, platinum) can be oxidized by these oxidizing acids. As a result of these reactions, the corresponding salts will always be formed, water and the product of nitrogen or sulfur reduction, respectively.
d) with alkalis
Metals that form amphoteric compounds (aluminum, beryllium, zinc) are capable of reacting with melts (with the formation of medium salts of aluminates, beryllates or zincates) or alkali solutions (with the formation of the corresponding complex salts). All reactions will release hydrogen.
e) In accordance with the position of the metal in the activity series, reactions of reduction (displacement) of a less active metal from a solution of its salt by another more active metal are possible. As a result of the reaction, a salt of a more active and simple substance is formed - a less active metal.

General properties of nonmetals.

There are much fewer non-metals than metals (22 elements). However, the chemistry of non-metals is much more complicated due to the greater filling of the external energy level of their atoms.
The physical properties of non-metals are more diverse: among them are gaseous (fluorine, chlorine, oxygen, nitrogen, hydrogen), liquids (bromine) and solids, which differ greatly from each other in melting point. Most non-metals do not conduct electricity, but silicon, graphite, germanium have semiconductor properties.
Gaseous, liquid and some solid non-metals (iodine) have a molecular structure of the crystal lattice, the rest of the non-metals have an atomic crystal lattice.
Fluorine, chlorine, bromine, iodine, oxygen, nitrogen and hydrogen under normal conditions exist in the form of diatomic molecules.
Many non-metal elements form several allotropic modifications of simple substances. So oxygen has two allotropic modifications - oxygen O2 and ozone O3, sulfur has three allotropic modifications - rhombic, plastic and monoclinic sulfur, phosphorus has three allotropic modifications - red, white and black phosphorus, carbon - six allotropic modifications - soot, graphite, diamond , carbine, fullerene, graphene.

Unlike metals, which exhibit only reducing properties, non-metals in reactions with simple and complex substances can act both as a reducing agent and as an oxidizing agent. According to their activity, non-metals occupy a certain place in the series of electronegativity. Fluorine is considered the most active non-metal. It exhibits only oxidizing properties. Oxygen is in second place in terms of activity, nitrogen is in third, then halogens and other non-metals. Hydrogen has the lowest electronegativity among non-metals.

Chemical properties of non-metals.

1. Interaction with simple substances:
Nonmetals interact with metals. In such a reaction, metals act as a reducing agent, non-metals as an oxidizing agent. As a result of the reaction of the compound, binary compounds are formed - oxides, peroxides, nitrides, hydrides, salts of oxygen-free acids.
In the reactions of non-metals with each other, a more electronegative non-metal exhibits the properties of an oxidizing agent, a less electronegative one - the properties of a reducing agent. As a result of the compound reaction, binary compounds are formed. It must be remembered that non-metals can exhibit variable oxidation states in their compounds.
2. Interaction with complex substances:
a) with water:
Under normal conditions, only halogens interact with water.
b) with oxides of metals and non-metals:
Many non-metals can react at high temperatures with oxides of other non-metals, reducing them to simple substances. Non-metals to the left of sulfur in the electronegativity series can also interact with metal oxides, reducing metals to simple substances.
c) with acids:
Some non-metals can be oxidized with concentrated sulfuric or nitric acids.
d) with alkalis:
Under the action of alkalis, some non-metals can undergo dismutation, being both an oxidizing agent and a reducing agent.
For example, in the reaction of halogens with alkali solutions without heating: Cl2 + 2NaOH = NaCl + NaClO + H2O or when heated: 3Cl2 + 6NaOH = 5NaCl + NaClO3 + 3H2O.
e) with salts:
When interacting, being strong oxidizing agents, they exhibit reducing properties.
Halogens (except fluorine) enter into substitution reactions with solutions of salts of hydrohalic acids: a more active halogen displaces a less active halogen from a salt solution.

Chemical properties of simple substances - non-metals

Chemical properties of hydrogen

From the point of view of the properties of hydrogen as a simple substance, it nevertheless has more in common with halogens. Hydrogen, like halogens, is a non-metal and forms diatomic molecules similarly to them (H 2 ).

Under normal conditions, hydrogen is a gaseous, inactive substance. The low activity of hydrogen is explained by the high strength of the bond between hydrogen atoms in the molecule, which requires either strong heating or the use of catalysts, or both at the same time, to break it.

Interaction of hydrogen with simple substances

with metals

Among metals, hydrogen reacts only withalkaline and alkaline earth! Alkali metals include metals of the main subgroup of the 1st group (Li, Na, K, Rb, Cs, Fr), and alkaline earth metals - the metals of the main subgroup of the 2nd group, except for beryllium and magnesium (Ca, Sr, Ba, Ra)

When interacting with active metals, hydrogen exhibits oxidizing properties, i.e. lowers its oxidation state. In this case, hydrides of alkali and alkaline earth metals are formed, which have an ionic structure. The reaction proceeds when heated:

2Na+H 2 = 2NaH

Ca + H 2 = CaH 2

It should be noted that the interaction with active metals is the only case when molecular hydrogen H 2 is an oxidizing agent.

with non-metals

Of non-metals, hydrogen reacts only with carbon, nitrogen, oxygen, sulfur, selenium and halogens!

Carbon should be understood as graphite or amorphous carbon, since diamond is an extremely inert allotropic modification of carbon.

When interacting with non-metals, hydrogen can only perform the function of a reducing agent, that is, it can only increase its oxidation state:

Interaction of hydrogen with complex substances

with metal oxides

Hydrogen does not react with metal oxides that are in the activity series of metals up to aluminum (inclusive), however, it is able to reduce many metal oxides to the right of aluminum when heated:

with non-metal oxides

Of the non-metal oxides, hydrogen reacts when heated with oxides of nitrogen, halogens, and carbon. Of all the interactions of hydrogen with non-metal oxides, its reaction with carbon monoxide CO should be especially noted.

Mixture of CO and H 2 even has its own name - “synthesis gas”, since, depending on the conditions, such demanded industrial products as methanol, formaldehyde and even synthetic hydrocarbons can be obtained from it:

with acids

Hydrogen does not react with inorganic acids!

Of the organic acids, hydrogen reacts only with unsaturated acids, as well as with acids containing functional groups capable of being reduced by hydrogen, in particularaldehyde, keto or nitro groups .

with salts

In the case of aqueous solutions of salts, their interaction with hydrogen does not occur. However, when hydrogen is passed over solid salts of some metals of medium and low activity, their partial or complete reduction is possible, for example:

Chemical properties of halogens

Halogens are the chemical elements of group VIIA (F, Cl, Br, I, At), as well as the simple substances they form. Hereinafter, unless otherwise stated, halogens will be understood as simple substances.

All halogens have a molecular structure, which leads to low melting and boiling points of these substances. Halogen molecules are diatomic, i.e. their formula can be written in general form as Hal 2 .

Halogen

Physical Properties

F 2 Light yellow gas with a pungent, irritating odor

Cl 2 Yellow-green gas with a pungent, suffocating odor

Br 2 Red-brown liquid with a pungent odor

I 2 Solid substance with a pungent odor, forming black-violet crystals

It should be noted such a specific physical property of iodine as its ability to sublimate or, in other words, sublimation. Sublimation is a phenomenon in which a substance in the solid state does not melt when heated, but, bypassing the liquid phase, immediately passes into the gaseous state.

As you know, the electronegativity of non-metals decreases when moving down the subgroup, and therefore the activity of halogens decreases in the series: F 2 >Cl 2 >Br 2 > I 2

Interaction of halogens with simple substances

All halogens are highly reactive and react with most simple substances. However, it should be noted that fluorine, due to its extremely high reactivity, can react even with those simple substances with which other halogens cannot react. Such simple substances include oxygen, carbon (diamond), nitrogen, platinum, gold, and some noble gases (xenon and krypton). Those. in fact, fluorine does not react only with certain noble gases.

The remaining halogens, i.e. chlorine, bromine and iodine are also active substances, but less active than fluorine. They react with almost all simple substances except oxygen, nitrogen, carbon in the form of diamond, platinum, gold and noble gases.

Interaction of halogens with non-metals

hydrogen

When all halogens react with hydrogen, hydrogen halides are formed with the general formula HHal. At the same time, the reaction of fluorine with hydrogen begins spontaneously even in the dark and proceeds with an explosion in accordance with the equation: H 2 + F 2 = 2HF

The reaction of chlorine with hydrogen can be initiated by intense ultraviolet irradiation or heating. Also leaks with an explosion: H 2 +Cl 2 = 2HCl

Bromine and iodine react with hydrogen only when heated, and at the same time, the reactionwith iodine is reversible: H 2 + Br 2 = 2 HBr

phosphorus

The interaction of fluorine with phosphorus leads to the oxidation of phosphorus to the highest oxidation state (+5). In this case, the formation of phosphorus pentafluoride occurs: 2P + 5F 2 = 2PF 5

When chlorine and bromine interact with phosphorus, it is possible to obtain phosphorus halides both in the + 3 oxidation state and in the + 5 oxidation state, which depends on the proportions of the reactants:

In the case of white phosphorus in an atmosphere of fluorine, chlorine or liquid bromine, the reaction begins spontaneously.

The interaction of phosphorus with iodine can lead to the formation of only phosphorus triiodide due to the significantly lower oxidizing ability than other halogens:

gray

Fluorine oxidizes sulfur to the highest oxidation state +6, forming sulfur hexafluoride:

Chlorine and bromine react with sulfur, forming compounds containing sulfur in oxidation states that are extremely unusual for it +1 and +2. These interactions are very specific, and to pass the exam in chemistry, the ability to write down the equations of these interactions is not necessary. Therefore, the following three equations are given rather for guidance:

interaction of sulfur with chlorine and bromine

Interaction of halogens with metals

As mentioned above, fluorine is able to react with all metals, even such inactive ones as platinum and gold:

The remaining halogens react with all metals except platinum and gold:

Reactions of halogens with complex substances

Substitution reactions with halogens

More active halogens, i.e. the chemical elements of which are located higher in the periodic table, are able to displace less active halogens from the hydrohalic acids and metal halides they form:

Similarly, bromine and iodine displace sulfur from solutions of sulfides and or hydrogen sulfide:

Chlorine is a stronger oxidizing agent and oxidizes hydrogen sulfide in its aqueous solution not to sulfur, but to sulfuric acid:

Interaction of halogens with water

Water burns in fluorine with a blue flame in accordance with the reaction equation:

Bromine and chlorine react differently with water than fluorine. If fluorine acted as an oxidizing agent, then chlorine and bromine disproportionate in water, forming a mixture of acids. In this case, the reactions are reversible:

The interaction of iodine with water proceeds to such an insignificant degree that it can be neglected and considered that the reaction does not proceed at all.

Interaction of halogens with alkali solutions

Fluorine, when interacting with an aqueous solution of alkali, again acts as an oxidizing agent:

The ability to write this equation is not required to pass the exam. It is enough to know the fact about the possibility of such an interaction and the oxidizing role of fluorine in this reaction.

Unlike fluorine, the remaining halogens disproportionate in alkali solutions, that is, they simultaneously increase and decrease their oxidation state. At the same time, in the case of chlorine and bromine, depending on the temperature, flow in two different directions is possible. In particular, in the cold, the reactions proceed as follows:

Iodine reacts with alkalis exclusively according to the second option, i.e. with the formation of iodate, because hypoiodite is not stable not only when heated, but also at normal temperature and even in the cold:

Chemical properties of oxygen

The chemical element oxygen can exist in the form of two allotropic modifications, i.e. forms two simple substances. Both of these substances have a molecular structure. One of them has the formula O 2 and has the name oxygen, i.e. the same as the name of the chemical element with which it is formed.

Another simple substance formed by oxygen is called ozone. Ozone, unlike oxygen, consists of triatomic molecules, i.e. has the formula O 3 .

Since the main and most common form of oxygen is molecular oxygen O 2 First of all, we will consider its chemical properties.

The chemical element oxygen is in second place in terms of electronegativity among all elements and is second only to fluorine. In this regard, it is logical to assume the high activity of oxygen and the presence of almost exclusively oxidizing properties in it. Indeed, the list of simple and complex substances with which oxygen can react is huge. However, it should be noted that since there is a strong double bond in the oxygen molecule, most reactions with oxygen require the use of heat. Most often, strong heating is required at the very beginning of the reaction (ignition), after which many reactions proceed independently without heat supply from the outside.

Among simple substances, only noble metals (Ag, Pt, Au), halogens and inert gases are not oxidized by oxygen.

Sulfur burns in oxygen to form sulfur dioxide:

Characteristic chemical properties of oxygen and sulfur

Phosphorus depending on the excess or lack of oxygen, it can form both phosphorus (V) oxide and phosphorus (III) oxide:

Interaction of oxygenwith nitrogen proceeds under extremely harsh conditions, since the binding energies in oxygen and especially nitrogen molecules are very high. The high electronegativity of both elements also contributes to the complexity of the reaction. The reaction begins only at temperatures above 2000 o C and is reversible:

Not all simple substances react with oxygen to form oxides. So, for example, sodium, burning in oxygen, forms a peroxide:

Most often, when complex substances are burned in oxygen, a mixture of oxides of the elements that formed the original substance is formed. For example:

However, when nitrogen-containing organic substances are burned in oxygen, molecular nitrogen N is formed instead of nitric oxide. 2 . For example:

When chlorine derivatives are burned in oxygen, instead of chlorine oxides, hydrogen chloride is formed:

Chemical properties of ozone:

Ozone is a stronger oxidizing agent than oxygen. This is due to the fact that one of the oxygen-oxygen bonds in the ozone molecule breaks easily and as a result, extremely active atomic oxygen is formed. Ozone, unlike oxygen, does not require heating to manifest its high oxidizing properties. It shows its activity at ordinary and even low temperatures: PbS + 4O 3 = PbSO 4 + 4O 2

As stated above,silver does not react with oxygen, however, it reacts with ozone:

2Ag+O 3 = Ag 2 O+O 2

A qualitative reaction to the presence of ozone is that when the test gas is passed through a solution of potassium iodide, the formation of iodine is observed:

2KI+O 3 + H 2 O=I 2 ↓ +O 2 + 2KOH

Chemical properties of sulfur

Sulfur as a chemical element can exist in several allotropic modifications. Distinguish rhombic, monoclinic and plastic sulfur. Monoclinic sulfur can be obtained by slow cooling of a rhombic sulfur melt, while plastic, on the contrary, is obtained by sharp cooling of a sulfur melt that has been previously brought to a boil. Plastic sulfur has a rare property of elasticity for inorganic substances - it is able to reversibly stretch under the action of an external force, returning to its original form when this effect is terminated. The rhombic sulfur is the most stable under normal conditions, and all other allotropic modifications pass into it over time.

The rhombic sulfur molecules consist of eight atoms, i.e. its formula can be written as S 8 . However, since the chemical properties of all modifications are quite similar, so as not to make it difficult to write the reaction equations, any sulfur is simply denoted by the symbol S.

Sulfur can interact with both simple and complex substances. In chemical reactions, it exhibits both oxidizing and reducing properties.

Oxidizing properties of sulfur appear when it interacts with metals, as well as non-metals formed by atoms of a less electronegative element (hydrogen, carbon, phosphorus):

As a reducing agent, sulfur acts when interacting with non-metals formed by more electronegative elements (oxygen, halogens), as well as complex substances with a pronounced oxidizing function, for example, concentrated sulfuric and nitric acids:

Sulfur also interacts during boiling with concentrated aqueous solutions of alkalis. The interaction proceeds according to the type of disproportionation, i.e. sulfur simultaneously both lowers and increases its oxidation state:

Chemical properties of nitrogen

The chemical element nitrogen forms only one simple substance. This substance is gaseous and is formed by diatomic molecules, i.e. has the formula N 2 . Despite the fact that the chemical element nitrogen has a high electronegativity, molecular nitrogen N 2 is an extremely inert substance. This fact is due to the fact that an extremely strong triple bond (N≡N) takes place in the nitrogen molecule. For this reason, almost all reactions with nitrogen proceed only at elevated temperatures.

Interaction of nitrogen with metals

The only substance that reacts with nitrogen under normal conditions is lithium:

Interesting is the fact that with other active metals, i.e. alkaline and alkaline earth, nitrogen reacts only when heated:

The interaction of nitrogen with metals of medium and low activity (except for Pt and Au) is also possible, but requires incomparably higher temperatures.

Interaction of nitrogen with non-metals

Nitrogen reacts with hydrogen when heated in the presence of catalysts. The reaction is reversible, therefore, to increase the ammonia yield in industry, the process is carried out at high pressure:

As a reducing agent, nitrogen reacts with fluorine and oxygen. With fluorine, the reaction proceeds under the action of an electric discharge:

With oxygen, the reaction proceeds under the influence of an electric discharge or at a temperature of more than 2000 O C and is reversible:

Of the non-metals, nitrogen does not react with halogens and sulfur.

The interaction of nitrogen with complex substances

As part of the USE school course, we can assume that nitrogen does not react with any complex substances other than active metal hydrides:

Chemical properties of phosphorus

There are several allotropic modifications of phosphorus, in particular white phosphorus, red phosphorus and black phosphorus.

White phosphorus is formed by four-atomic molecules P 4 , is not a stable modification of phosphorus. Poisonous. At room temperature, it is soft and, like wax, can be easily cut with a knife. In air, it slowly oxidizes, and due to the peculiarities of the mechanism of such oxidation, it glows in the dark (the phenomenon of chemiluminescence). Even with low heating, spontaneous ignition of white phosphorus is possible.

Of all the allotropic modifications, white phosphorus is the most active.

Red phosphorus consists of long molecules of variable composition Pn. Some sources indicate that it has an atomic structure, but it is still more correct to consider its structure as molecular. Due to structural features, it is a less active substance compared to white phosphorus, in particular, unlike white phosphorus, it oxidizes much more slowly in air and requires ignition to ignite it.

Black phosphorus consists of continuous Pn chains and has a layered structure similar to that of graphite, which is why it looks like it. This allotropic modification has an atomic structure. The most stable of all allotropic modifications of phosphorus, the most chemically passive. For this reason, the chemical properties of phosphorus discussed below should be attributed primarily to white and red phosphorus.

The interaction of phosphorus with non-metals

The reactivity of phosphorus is higher than that of nitrogen. So, phosphorus is able to burn after ignition under normal conditions, forming an acid oxide P 2 O 5 :

and with a lack of oxygen, phosphorus (III) oxide:

The reaction with halogens also proceeds intensively. So, during chlorination and bromination of phosphorus, depending on the proportions of the reagents, phosphorus trihalides or pentahalides are formed:

Due to the significantly weaker oxidizing properties of iodine compared to other halogens, it is possible to oxidize phosphorus with iodine only to an oxidation state of +3:

Unlike nitrogen, phosphorus does not react with hydrogen.

The interaction of phosphorus with metals

Phosphorus reacts when heated with active metals and metals of medium activity to form phosphides:

The interaction of phosphorus with complex substances

Phosphorus is oxidized by oxidizing acids, in particular, concentrated nitric and sulfuric acids:

interaction of phosphorus with oxidizing acids

You should know that white phosphorus reacts with aqueous solutions of alkalis. However, due to the specificity, the ability to write down the equations of such interactions for the Unified State Examination in Chemistry has not yet been required.

Nevertheless, for those who claim 100 points, for their own peace of mind, you can remember the following features of the interaction of phosphorus with alkali solutions in the cold and when heated.

In the cold, the interaction of white phosphorus with alkali solutions proceeds slowly. The reaction is accompanied by the formationgas with the smell of rotten fish - phosphine and compounds with a rare oxidation state of phosphorus +1:

When white phosphorus interacts with a concentrated alkali solution, hydrogen is released during boiling and phosphite is formed:

Chemical properties of carbon

Carbon is capable of forming several allotropic modifications. These are diamond (the most inert allotropic modification), graphite, fullerene and carbine.

Charcoal and soot are amorphous carbon. Carbon in this state does not have an ordered structure and actually consists of the smallest fragments of graphite layers. Amorphous carbon treated with hot water vapor is called activated carbon. 1 gram of activated carbon, due to the presence of many pores in it, has a total surface of more than three hundred square meters! Due to its ability to absorb various substances, activated carbon is widely used as a filter filler, as well as an enterosorbent for various types of poisoning.

From a chemical point of view, amorphous carbon is its most active form, graphite exhibits medium activity, and diamond is an extremely inert substance. For this reason, the chemical properties of carbon considered below should primarily be attributed to amorphous carbon.

Reducing properties of carbon

As a reducing agent, carbon reacts with non-metals such as oxygen, halogens, and sulfur.

Depending on the excess or lack of oxygen, coal combustion may produce carbon monoxide CO or carbon dioxide CO 2 :

When carbon interacts with fluorine carbon tetrafluoride is formed:

When carbon is heated with sulfur carbon disulfide CS is formed 2 :

Carbon can reduce metals after aluminum in the activity series of their oxides. For example:

Alsocarbon also reacts with oxides of active metals , however, in this case, as a rule, it is not the reduction of the metal that is observed, but the formation of its carbide:

Interaction of carbon with non-metal oxides

Carbon enters into a co-proportionation reaction with carbon dioxide CO 2 :

One of the most important processes from an industrial point of view is the so-called steam reforming of coal. The process is carried out by passing water vapor through hot coal. In this case, the following reaction takes place:

At high temperatures, carbon is able to reduce even such an inert compound as silicon dioxide. In this case, depending on the conditions, the formation of silicon or silicon carbide (carborundum) is possible:

Also, carbon as a reducing agent reacts with oxidizing acids, in particular, concentrated sulfuric and nitric acids:

Oxidizing properties of carbon

The chemical element carbon is not highly electronegative, so the simple substances it forms rarely exhibit oxidizing properties with respect to other non-metals.

An example of such reactions is the interaction of amorphous carbon with hydrogen when heated in the presence of a catalyst:

as well as with silicon at a temperature of 1200-1300 O WITH:

Carbon exhibits oxidizing properties in relation to metals . Carbon is able to react with active metals and some metals of intermediate activity. Reactions proceed when heated:

Active metal carbides are hydrolyzed by water:

as well as solutions of non-oxidizing acids:

In this case, hydrocarbons are formed containing carbon in the same oxidation state as in the original carbide.

Chemical properties of silicon

Silicon can exist, as well as carbon in the crystalline and amorphous state, and, just as in the case of carbon, amorphous silicon is significantly more chemically active than crystalline silicon.

Sometimes amorphous and crystalline silicon is called its allotropic modifications, which, strictly speaking, is not entirely true. Amorphous silicon is essentially a conglomerate of the smallest particles of crystalline silicon randomly arranged relative to each other.

Interaction of silicon with simple substances

non-metals

Under normal conditions, silicon, due to its inertness, reacts only with fluorine:

Si+2F 2 = SiF 4

Silicon reacts with chlorine, bromine and iodine only when heated. It is characteristic that, depending on the activity of the halogen, a correspondingly different temperature is required:

All silicon halides are easily hydrolyzed by water:

as well as alkali solutions:

The reaction of silicon with oxygen proceeds, but requires very strong heating (1200-1300 O C) due to the fact that a strong oxide film makes it difficult to interact:

At a temperature of 1200-1500 O With silicon slowly interacts with carbon in the form of graphite with the formation of carborundum SiC - a substance with an atomic crystal lattice similar to diamond and almost not inferior to it in strength:

Silicon does not react with hydrogen.

metals

Due to its low electronegativity, hydrogen can exhibit oxidizing properties only with respect to metals. Of the metals, silicon reacts with active (alkaline and alkaline earth), as well as many metals of medium activity. As a result of this interaction, silicides are formed: 2Mg + Si = Mg 2 Si

Silicides of active metals are easily hydrolyzed with water or dilute solutions of non-oxidizing acids:

This produces a gas silane SiH 4 – analogue of methane CH 4 .

Interaction of silicon with complex substances

Silicon does not react with water even when boiling, but amorphous silicon interacts with superheated water vapor at a temperature of about 400-500 O C. This produces hydrogen and silicon dioxide:

Of all acids, silicon (in its amorphous state) reacts only with concentrated hydrofluoric acid:

Silicon dissolves in concentrated alkali solutions. The reaction is accompanied by the evolution of hydrogen: