Unified State Exam. CHEMICAL PROPERTIES OF NON-METALS

CHEMICAL PROPERTIES OF HYDROGEN

1. WITH METALS

(Li, Na, K, Rb, Cs, Ca, Sr, Ba) → forms unstable solid hydrides when heated with alkali and alkaline earth metals; other metals do not react.

2K + H₂ = 2KH (potassium hydride)

Ca + H₂ = CaH₂

2. WITH NON-METALS

with oxygen, halogens under normal conditions, when heated it reacts with phosphorus, silicon and carbon, with nitrogen in the presence of pressure and a catalyst.

2Н₂ + O₂ = 2Н₂O Н₂ + Cl₂ = 2HCl

3Н₂ + N₂↔ 2NH₃ H₂ + S = H₂S

3. INTERACTION WITH WATER

Does not react with water

4. INTERACTION WITH OXIDES

Reduces oxides of metals (inactive) and non-metals to simple substances:

CuO + H₂ = Cu + H₂O 2NO + 2H₂ = N₂ + 2H₂O

SiO₂ + H₂ = Si + H₂O

5. INTERACTION WITH ACIDS

Does not react with acids

6. INTERACTION WITH ALKALI

Does not react with alkalis

7. INTERACTION WITH SALT

Recovers low-active metals from salts

CuCl₂ + H₂ = Cu + 2HCl

CHEMICAL PROPERTIES OF OXYGEN

1. INTERACTION WITH METALS

With alkali metals under normal conditions - oxides and peroxides (lithium - oxide, sodium - peroxide, potassium, cesium, rubidium - superoxide

4Li + O2 = 2Li2O (oxide)

2Na + O2 = Na2O2 (peroxide)

K+O2=KO2 (superoxide)

With other metals of the main subgroups under normal conditions it forms oxides with an oxidation state equal to the group number

2 WITHa+O2=2WITHaO

4Al + O2 = 2Al2O3

1. INTERACTION WITH METALS

With metals of secondary subgroups, under normal conditions and when heated, it forms oxides of varying degrees of oxidation, and with iron, iron scaleFe3 O4 ( FeO∙ Fe2 O3)

3Fe + 2O2 = Fe3O4 4Cu + O₂ = 2Cu₂⁺¹O (red);

2Cu + O₂ = 2Cu⁺²O (black); 2Zn + O₂ = ZnO

4Cr + 3O2 = 2Cr2⁺³O3

forms oxides – often of intermediate oxidation state

C + O₂(izb)=CO₂; C+ O₂ (week) =CO

S + O₂ = SO₂N₂ + O₂ = 2NO - Q

3. INTERACTION WITH WATER

Does not react with water

4. INTERACTION WITH OXIDES

Oxidizes lower oxides to oxides with a higher oxidation state

Fe⁺²O + O2 = Fe2⁺³O3; C⁺²O + O2 = C⁺⁴O2

5. INTERACTION WITH ACIDS

Anhydrous oxygen-free acids (binary compounds) burn in an oxygen atmosphere

2H2S + O2 = 2S + 2H2O 2H2S + 3O2 = 2SO2 + 2H2O

In oxygen-containing compounds, it increases the oxidation state of the non-metal.

2HN⁺³O2 + O2 = 2HN⁺⁵O3

6.INTERACTION WITH BASES

Oxidizes unstable hydroxides in aqueous solutions to a higher oxidation state

4Fe(OH)2 + O2 + 2H2O = 4Fe(OH)3

7. INTERACTION WITH SALT AND BINARY COMPOUNDS

Enters into combustion reactions.

4FeS2 +11O2 = 2Fe2O3 + 8SO2

CH4 + 2O2 = CO2 + 2H2O

4NH3 + 3O2 = 2N2 + 6H2O

Catalytic oxidation

NH3 + O2 = NO + H2O

CHEMICAL PROPERTIES OF HALOGENS

1. INTERACTION WITH METALS

With alkaline under normal conditions, withF, Cl, Brignite:

2 Na + Cl2 = 2 NaCl(chloride)

Alkaline earths and aluminum react under normal conditions:

WITHa+Cl2=WITHaCl2 2Al+3Cl2 = 2AlCl3

Metals of secondary subgroups at elevated temperatures

Cu + Cl₂ = Cu⁺²Cl₂

2Cu + I₂ = 2Cu⁺¹I (there is no copper (II) iodide!)

2Fe + 3С12 = 2Fe⁺³Cl3 iron (III) chloride

Fluorine reacts with metals (often explosively), including gold and platinum.

2Au + 3F₂ = 2AuF

2. INTERACTION WITH NON-METALS

They do not interact directly with oxygen (except for F₂), but react with sulfur, phosphorus, and silicon. The chemical activity of bromine and iodine is less pronounced than that of fluorine and chlorine:

H2+F2 = 2HF ; Si + 2 F2 = SiF4.; 2 P + 3 Cl2 = 2 P⁺³ Cl3; 2 P + 5 Cl2 = 2 P⁺⁵ Cl5; S + 3 F2 = S⁺⁶ F6;

S + Cl2 = S⁺²Cl2

F₂

Reacts with oxygen:F2 + O2 = O⁺² F2

Reacts with other halogens:Cl₂ + F₂ = 2 Cl⁺¹ F¯¹

Reacts even with inert gases 2F₂ + Xe= Xe⁺⁸ F₄¯¹.

3. INTERACTION WITH WATER

Fluorine under normal conditions forms hydrofluoric acid + + O₂

2F2 + 2H2O → 4НF + О2

When the temperature increases, chlorine forms hydrochloric acid + O₂,

2Сl₂ + 2H₂O → 4HCl + O₂

at no. - “chlorine water”

Cl2 + H2O ↔ HCl + HClO (hydrochloric and hypochlorous acids)

Bromine under normal conditions forms “bromine water”

Br2 + H2O ↔ HBr + HBrO (hydrobromic and hypobromous acids

Iodine →reaction does not occur

I2 + H₂O ≠

5. INTERACTION WITH OXIDES

Only fluorine F₂ reacts, displacing oxygen from the oxide, forming fluorides

SiO2‾² + 2F2⁰ = SiF4‾¹ + O2⁰

6. INTERACTION WITH ACIDS.

react with oxygen-free acids, displacing less active nonmetals.

H2S‾² + I2⁰ → S⁰↓+ 2HI‾

7. INTERACTION WITH ALKALI

FLUORINE forms fluoride + oxygen and water

2F2 + 4NaOH = 4NaF¯¹ + O2 + 2H2O

CHLORINE when heated forms chloride, chlorate and water

3 Cl₂ + 6 KOH = 5 KCl¯¹ + KCl⁺⁵ O3 + 3 H2 O

In the cold, chloride, hypochlorate and water, with calcium hydroxide, bleach and water

Cl2 + 2KOH-(cold)= KCl¯¹ + KCl⁺¹O + H2O

Cl2 + Ca(OH) 2 = CaOCl2 (bleach – a mixture of chloride, hypochlorite and hydroxide) + H2O

Bromine when heated → bromide, bromate and water

3Br2 + 6KOH =5KBr¯¹ + KBr ⁺⁵O3 + 3H2O

Iodine when heated → iodide, iodate and water

3I2 + 6NaOH = 5NaI¯¹ + NaI ⁺⁵O3 + 3H2O

9. INTERACTION WITH SALT

Displacement of less active halogens from salts

2KBr + Cl2 → 2KCl + Br2
2KCl + Br2 ≠
2KCl + F2→ 2KF + Cl2
2KBr + J2≠

Oxidize nonmetals in salts to a higher oxidation state

2Fe⁺²Cl2 + Cl2⁰ → 2Fe⁺³Cl 3 ‾¹

Na2S⁺⁴O3 + Br2⁰ + 2H2O →Na2S⁺⁶O4 + 2HBr‾

CHEMICAL PROPERTIES OF SULFUR

1. INTERACTION WITH METALS

reacts when heated even with alkali metals, with mercury under normal conditions: with sulfur - sulfides:

2K + S = K2S

2Cr + 3S = Cr2⁺³S3 Fe + S = Fe⁺²S

2. INTERACTION WITH NON-METALS

When heated with hydrogen,coxygen (sulfur dioxide)chalogens (except iodine), with carbon, nitrogen and silicon and does not react

S + Cl₂ = S⁺²Cl₂ ; S + O₂ =S⁺⁴O₂

H₂ + S = H₂S¯² ; 2P + 3S = P₂S₃¯²

WITH+ 3S = CS₂¯²

WITH WATER, OXIDES, SALTS

DOES NOT REACT

3. INTERACTION WITH ACIDS

Oxidized by sulfuric acid when heated to sulfur dioxide and water

2H2SO4 (conc.) = 2H2O + 3S⁺⁴O2

Nitric acid when heated to sulfuric acid, nitric oxide (+4) and water

S+6HNO3(conc.) =H2SO4 + 6N⁺⁴O2 + 2H2O

4. INTERACTION WITH ALKALI

When heated, forms sulfite, sulfide + water

3S + 6KOH = K2SO3 + 2K2S + 3H2O

CHEMICAL PROPERTIES OF NITROGEN

1. INTERACTION WITH METALS

reactions occur when heated (exception: lithium with nitrogen under normal conditions):

With nitrogen - nitrides

6Li + N2 = 3Li2N (lithium nitride) (n.s.) 3Mg + N2 =Mg3N2 (magnesium nitride) 2Cr + N2 = 2CrN

Iron in these compounds has an oxidation state of +2

2. INTERACTION WITH NON-METALS

(due to the triple bond, nitrogen is very inactive). Under normal conditions it does not react with oxygen. Reacts with oxygen only at high temperatures (electric arc), in nature - during a thunderstorm

N2+O2=2NO (el. arc, 3000 0C)

With hydrogen at high pressure, elevated temperature and in the presence of a catalyst:

t,p,kat

3N2+3H2 ↔ 2NH3

WITH WATER, OXIDES, ACIDS, ALKALI AND SALTS

DOES NOT REACT

CHEMICAL PROPERTIES OF PHOSPHORUS

1. INTERACTION WITH METALS

reactions occur when heated with phosphorus - phosphides

3Ca + 2P =K3P2, Iron in these compounds has an oxidation state of +2

2. INTERACTION WITH NON-METALS

Combustion in oxygen

4P + 5O₂ = 2P₂⁺⁵O₅ 4P + 3O₂ = 2P₂⁺³O₃

With halogens and sulfur when heated

2P + 3Cl₂ = 2P⁺³Cl₃ 2P + 5Cl₂ = 2P⁺⁵Cl₅; 2P + 5S = P₂⁺⁵S₅

Does not directly interact with hydrogen, carbon, silicon

WITH WATER AND OXIDES

DOES NOT REACT

3. INTERACTION WITH ACIDS

With concentrated nitric acid, nitric oxide (+4), with dilute nitric oxide (+2) and phosphoric acid

3P + 5HNO₃(conc) =3H₃PO₄ + 5N⁺⁴O₂

3P + 5HNO₃ + 2H₂O =3H₃PO₄ + 5N⁺²O

With concentrated sulfuric acid, phosphoric acid, sulfur oxide (+4) and water are formed

3P + 5H₂SO₄(conc.) =3H₃PO₄ + 5S⁺⁴O₂+ 2H₂O

4. INTERACTION WITH ALKALI

Forms phosphine and hypophosphite with alkali solutions

4P⁰ + 3NaOH + 3H2O = P¯³H 3 + 3NaH 2 P ⁺1O 2

5. INTERACTION WITH SALT

5. INTERACTION WITH SALT

With strong oxidizing agents, exhibiting reducing properties

3P⁰ + 5NaN⁺⁵O₃ = 5NaN⁺³O₂ + P₂⁺⁵O₅

CHEMICAL PROPERTIES OF CARBON

1. INTERACTION WITH METALS

reactions occur when heated

Metals - d-elements form compounds of non-stoichiometric composition with carbon, such as solid solutions: WC, ZnC, TiC - are used to produce superhard steels

with carbon carbides 2Li + 2C = Li2C2,

Ca + 2C = CaC2

2. INTERACTION WITH NON-METALS

Of the halogens, it reacts directly only with fluorine, with the rest when heated.

C + 2F₂ = CF₄.

Interaction with oxygen:

2С + О₂ (insufficient) = 2С⁺²О (carbon monoxide),

C + O₂(g) = C⁺⁴O₂(carbon dioxide).

Interaction with other non-metals at elevated temperatures, does not interact with phosphorus

C + Si = SiC¯⁴ ; C + N₂ = C₂⁺⁴N₂ ;

C + 2H₂ = C¯⁴H₄ ; C + 2S = C⁺⁴S₂;

3. INTERACTION WITH WATER

Passing water vapor through hot coal produces carbon monoxide and hydrogen (synthesis gas)

C + H₂O = CO + H₂

4. INTERACTION WITH OXIDES

CARBON RESTORES METALS AND NON-METALS FROM OXIDES TO SIMPLE SUBSTANCE WHEN HEATED (CARBOHERMY), in carbon dioxide it reduces the degree of oxidation

2ZnO + C = 2Zn + CO; 4WITH+ Fe₃O₄ = 3Fe + 4CO ;

P₂O₅ + C = 2P + 5CO; 2WITH+ SiO₂ = Si + 2CO;

WITH+ C⁺⁴O₂ = 2C⁺²O

5. INTERACTION WITH ACIDS

Oxidized with concentrated nitric and sulfuric acid to carbon dioxide

C +2H2SO4(conc)=C⁺⁴O2+ 2S⁺⁴O2+ 2H2O; C+4HNO3 (conc) =C⁺⁴O2 + 4N⁺⁴O2 + 2H2O.

WITH ALKALI AND SALT

DOES NOT REACT

CHEMICAL PROPERTIES OF SILICON

1. INTERACTION WITH METALS

reactions occur when heated: active metals - silicides - react with silicon

4Cs + Si = Cs4Si,

1. INTERACTION WITH NON-METALS

From halogens directly only with fluorine.

Reacts with chlorine when heated

Si + 2F2 = SiF4; Si + 2Cl2 = SiCl4;

Si + O₂ = SiO₂; Si + C = SiC; 3Si + 2N₂ = Si₃N;

Does not interact with hydrogen

3. INTERACTION WITH ACIDS

reacts only with a mixture of hydrofluoric and nitric acids, forming hexafluorosilicic acid

3Si + 4HNO₃ + 18HF = 3H₂ + 4NO + 8H₂O

Interaction with hydrogen halides (these are not acids) - displaces hydrogen, forming silicon halides and hydrogen

Reacts with hydrogen fluoride under normal conditions.

Si + 4HF = SiF₄ + 2H₂

4. INTERACTION WITH ALKALI

Dissolves when heated in alkalis, forming silicate and hydrogen:

Si +2NaOH +H₂O = Na₂SiO₃ + 2H₂

Position of non-metal elements in the Periodic Table of Chemical Elements D.I. Mendeleev

· Non-metal elements:

· s-element – ​​hydrogen;

· p-elements of group 3 – boron;

· 4 groups – carbon and silicon;

· 5 groups – nitrogen, phosphorus and arsenic,

· 6 groups – oxygen, sulfur, selenium and tellurium

· 7 groups – fluorine, chlorine, bromine, iodine and astatine.

Group 8 elements - inert gases - occupy a special position; they have a completely completed outer electron layer.

Non-metal chemical elements can exhibit both oxidizing and reducing properties, depending on the chemical transformation in which they take part.

The atoms of the most electronegative element - fluorine - are not capable of donating electrons; it always exhibits only oxidizing properties; other elements can also exhibit reducing properties, although to a much lesser extent than metals. The most powerful oxidizing agents (accept electrons) are fluorine, oxygen and chlorine; hydrogen, boron, carbon, silicon, phosphorus, arsenic and tellurium exhibit predominantly reducing properties (donate). Nitrogen, sulfur, and iodine have intermediate redox properties.

1. Interaction with metals:

2Na + Cl 2 = 2NaCl, Fe + S = FeS, 6Li + N 2 = 2Li 3 N, 2Ca + O 2 = 2CaO

in these cases, non-metals exhibit oxidizing properties; they accept electrons, forming negatively charged particles.

2. Interaction with other non-metals:

· interacting with hydrogen , most non-metals exhibit oxidizing properties, forming volatile hydrogen compounds - covalent hydrides:

3H 2 + N 2 = 2NH 3, H 2 + Br 2 = 2HBr;

· interacting with oxygen , all non-metals, except fluorine, exhibit reducing properties:

S + O 2 = SO 2, 4P + 5O 2 = 2P 2 O 5;

· during interaction with fluoride fluorine is an oxidizing agent, and oxygen is a reducing agent: 2F 2 + O 2 = 2OF 2 ;

· non-metals interact between themselves , a more electronegative metal plays the role of an oxidizing agent, a less electronegative metal plays the role of a reducing agent: S + 3F 2 = SF 6, C + 2Cl 2 = CCl 4.

Halogens (group 7)

Chemical properties of halogens.

OXYGEN-CONTAINING CHLORINE ACIDS

· Hypochlorous acid HCl +1 O salts – hypo chlorites

Exists only in the form of dilute aqueous solutions.

Obtaining Cl2 + H2O = HCl + HClO

Chemical properties

HClO is a weak acid and a strong oxidizing agent:

1) Decomposes in light, releasing atomic oxygen HClO = HCl + O

2) With alkalis it gives salts - hypochlorites HClO + KOH = KClO + H2O

3) Reacts with hydrogen halides 2HI + HClO = I2 + HCl + H2O

Chlorous acid HClO2 (HClO2 is a weak acid and a strong oxidizing agent; salts of chlorous acid - chlorites)

Chemical properties

1.HClO2 + KOH = KClO2 + H2O

2. Unstable, decomposes during storage 4HClO2 = HCl + HClO3 + 2ClO2 + H2O

Hypochlorous acid HCl O3 (HClO3 - Strong acid and strong oxidizing agent; salts of perchloric acid - chlorates)

KClO3 - Berthollet's salt; it is obtained by passing chlorine through a heated (40°C) KOH solution:

3Cl 2 + 6KOH = 5KCl + KClO 3 + 3H 2 O

Berthollet's salt is used as an oxidizing agent; When heated, it decomposes:

4KClO 3 = KCl + 3KClO 4 (without catalyst)

2KClO 3 = 2KCl + 3O 2 (MnO 2 catalyst)

Perchloric acid HClO4 (HClO4 is a very strong acid and a very strong oxidizing agent; salts of perchloric acid - perchlorates)

Preparation of KClO4 + H2SO4 = KHSO4 + HClO4

Chemical properties

1) Reacts with alkalis HClO4 + KOH = KClO4 + H2O

2) When heated, perchloric acid and its salts decompose:

4HClO4 = 4ClO2 + 3O2 + 2H2O KClO4 = KCl + 2O2

Chalcogens (group VIA elements)

Oxygen, S, Se, Te, Po. The name chalcogens means “giving birth to ores.” Sulfur compounds: pyrite, or iron pyrite - FeS2, cinnabar - HgS, zinc blende - ZnS.

Chalcogens have 6 electrons in their outer energy level. The atoms lack 2 electrons before completing the outer energy level, so they gain electrons and exhibit a -2 oxidation state in their compounds.

Sulfur, selenium and tellurium atoms in their compounds with more electronegative elements exhibit positive oxidation states of +2, +4 and +6.

Oxygen n=8 1s 2 2s 2 2p 4

Oxygen is part of such ores as corundum - Al2O3, magnetic iron ore - Fe3O4, red iron ore - Fe2O3, brown iron ore - Fe2O3

Oxygen combined with fluorine – OF2 exhibits an oxidation state of +2. Oxygen is part of the atmosphere, where it accounts for 21%.

Obtaining oxygen.

· In industry, oxygen is obtained from liquid air.

· Oxygen can also be obtained by decomposing water in a special device - an electrolyzer.

· Hydrogen peroxide (H2O2) is used in the laboratory. This reaction occurs in the presence of a catalyst - manganese oxide IV

· in the laboratory they also use the decomposition reaction of potassium permanganate - KMnO 4 - “potassium permanganate”.

· In laboratory conditions, oxygen is released when berthollet salt (potassium chlorate) is heated.

2KClO 3 = 2KCl + 3O 2 The catalyst is manganese oxide (MnO 2).

oxygen exists in the form of two allotropic modifications –O 2 and O 3 .

Chemical properties

Oxygen does not interact with halogens, noble gases, gold and platinum.

· Oxygen reacts vigorously with metals. For example, in a reaction with lithium, lithium oxide is formed, in a reaction with copper - copper (II) oxide.

4Li + O 2 = 2Li 2 O 2Cu + O 2 = 2CuO

· Oxygen reacts with non-metals.

S + O 2 = SO 2 4P + 5O 2 = 2P 2 O 5

Almost all reactions with oxygen are exothermic (that is, accompanied by the release of heat). The exception is the reaction of nitrogen with oxygen, which is endothermic.

N 2 + O 2 ↔ 2NO – Q

· Oxygen is a complex substance.

CH 4 + 2O 2 = CO 2 + 2H 2 O 2H 2 S + 3O 2 = 2SO 2 + 2H 2 O

SULFUR n=16 1s 2 2s 2 2p 6 3s 2 3p 4

If most metal elements are not colored, the only exceptions being copper and gold, then almost all non-metals have their own color: fluorine - orange-yellow, chlorine - greenish-yellow, bromine - brick-red, iodine - violet, sulfur - yellow, phosphorus can be white, red and black, and liquid oxygen is blue.

All nonmetals do not conduct heat or electricity because they do not have free charge carriers - electrons; all of them are used to form chemical bonds. Crystals of non-metals are non-plastic and brittle, since any deformation leads to the destruction of chemical bonds. Most non-metals do not have a metallic luster.

The physical properties of nonmetals are varied and are determined by different types of crystal lattices.

1.4.1 Allotropy

ALLOTROPY - the existence of chemical elements in two or more molecular or crystalline forms. For example, allotropes are ordinary oxygen O 2 and ozone O 3 ; in this case, allotropy is due to the formation of molecules with different numbers of atoms. Most often, allotropy is associated with the formation of crystals of various modifications. Carbon exists in two distinct crystalline allotropes: diamond and graphite. Previously it was believed that the so-called. amorphous forms of carbon, charcoal and soot are also its allotropic modifications, but it turned out that they have the same crystalline structure as graphite. Sulfur occurs in two crystalline modifications: orthorhombic (a-S) and monoclinic (b-S); at least three of its non-crystalline forms are known: l-S, m-S and violet. For phosphorus, white and red modifications have been well studied, black phosphorus has also been described; at temperatures below –77°C there is another type of white phosphorus. Allotropic modifications of As, Sn, Sb, Se, and, at high temperatures, of iron and many other elements have been discovered.

1.5. Chemical properties of non-metals

Non-metal chemical elements can exhibit both oxidizing and reducing properties, depending on the chemical transformation in which they take part.

The atoms of the most electronegative element - fluorine - are not capable of donating electrons; it always exhibits only oxidizing properties; other elements can also exhibit reducing properties, although to a much lesser extent than metals. The most powerful oxidizing agents are fluorine, oxygen and chlorine; hydrogen, boron, carbon, silicon, phosphorus, arsenic and tellurium exhibit predominantly reducing properties. Nitrogen, sulfur, and iodine have intermediate redox properties.

Interaction with simple substances

Interaction with metals:

2Na + Cl 2 = 2NaCl,

6Li + N 2 = 2Li 3 N,

2Ca + O2 = 2CaO

in these cases, non-metals exhibit oxidizing properties; they accept electrons, forming negatively charged particles.

Interaction with other non-metals:

When interacting with hydrogen, most non-metals exhibit oxidizing properties, forming volatile hydrogen compounds - covalent hydrides:

3H 2 + N 2 = 2NH 3,

H 2 + Br 2 = 2HBr;

When interacting with oxygen, all nonmetals, except fluorine, exhibit reducing properties:

S + O 2 = SO 2,

4P + 5O 2 = 2P 2 O 5 ;

When interacting with fluorine, fluorine is an oxidizing agent, and oxygen is a reducing agent:

2F 2 + O 2 = 2OF 2;

Nonmetals interact with each other, the more electronegative metal plays the role of an oxidizing agent, the less electronegative one plays the role of a reducing agent:

S + 3F 2 = SF 6,

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 metal substances never exhibit oxidizing properties.

Obtaining metals:
- reduction from oxides with carbon (C), carbon monoxide (CO), hydrogen (H2) or a more active metal (Al, Ca, Mg);
- reduction from salt solutions with a more active metal;
- electrolysis of solutions or melts of metal compounds - reduction 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 by non-metals such as halogens, oxygen, sulfur, and nitrogen. But most of these reactions require preheating to begin. Subsequently, the reaction can proceed with the release of a large amount of heat, which leads to 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, alkali (soluble base) and hydrogen are formed: Metal + H2O = Me(OH) + H2
When heated, other metals that are to the left of hydrogen in the activity series interact with water. Magnesium reacts with boiling water, aluminum - after special surface treatment, resulting in the formation of insoluble bases - magnesium hydroxide or aluminum hydroxide - and hydrogen is released. Metals in the activity series from zinc (inclusive) to lead (inclusive) interact with water vapor (i.e. above 100 C), and oxides of the corresponding metals and hydrogen are formed.
Metals located in the activity series to the right of hydrogen do not interact with water.
b) interaction with oxides:
active metals react by substitution reaction with oxides of other metals or non-metals, reducing them to simple substances.
c) interaction with acids:
Metals located in the activity series to the left of hydrogen react with acids to release hydrogen and form the corresponding salt. Metals located in the activity series to the right of hydrogen 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. These reactions will always produce the corresponding salts, water and the reduction product of nitrogen or sulfur, respectively.
d) with alkalis
Metals that form amphoteric compounds (aluminum, beryllium, zinc) are capable of reacting with melts (in this case, medium salts aluminates, beryllates or zincates are formed) or alkali solutions (in this case the corresponding complex salts are formed). All reactions will produce 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 metal and a simple substance - a less active metal - are formed.

General properties of non-metals.

There are much fewer nonmetals than metals (22 elements). However, the chemistry of nonmetals is much more complex due to the greater occupancy of the outer energy level of their atoms.
The physical properties of non-metals are more diverse: among them there are gaseous (fluorine, chlorine, oxygen, nitrogen, hydrogen), liquid (bromine) and solid substances that differ greatly from each other in melting point. Most nonmetals do not conduct electricity, but silicon, graphite, and germanium have semiconducting properties.
Gaseous, liquid and some solid non-metals (iodine) have a molecular structure of a crystal lattice, other 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 nonmetallic elements form several allotropic modifications of simple substances. So oxygen has two allotropic modifications - oxygen O2 and ozone O3, sulfur has three allotropic modifications - orthorhombic, plastic and monoclinic sulfur, phosphorus has three allotropic modifications - red, white and black phosphorus, carbon - six allotropic modifications - soot, graphite, diamond , carbyne, fullerene, graphene.

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

Chemical properties of nonmetals.

1. Interaction with simple substances:
Nonmetals interact with metals. In such reactions, metals act as a reducing agent, and non-metals act as an oxidizing agent. As a result of the compound reaction, binary compounds are formed - oxides, peroxides, nitrides, hydrides, salts of oxygen-free acids.
In the reactions of nonmetals with each other, the more electronegative nonmetal exhibits the properties of an oxidizing agent, and the less electronegative one exhibits the properties of a reducing agent. The compound reaction produces binary compounds. It must be remembered that non-metals can exhibit varying 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 nonmetals can react at high temperatures with oxides of other nonmetals, reducing them to simple substances. Nonmetals that are to the left of sulfur in the electronegativity series can also interact with metal oxides, reducing metals to simple substances.
c) with acids:
Some nonmetals can be oxidized with concentrated sulfuric or nitric acids.
d) with alkalis:
Under the influence of alkalis, some nonmetals 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 with heating: 3Cl2 + 6NaOH = 5NaCl + NaClO3 + 3H2O.
d) with salts:
When interacting, they are strong oxidizing agents and 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 the salt solution.

Chemical properties of non-metals
In accordance with the numerical values ​​of relative electronegativities the oxidizing capacity of nonmetals increases in the following order: Si, B, H, P, C, S, I, N, Cl, O, F.
Nonmetals as oxidizing agents
The oxidizing properties of nonmetals manifest themselves during their interaction:

· with metals: 2Na + Cl 2 = 2NaCl;

· with hydrogen: H 2 + F 2 = 2HF;

· with nonmetals that have lower electronegativity: 2P + 5S = P 2 S 5 ;

· with some complex substances: 4NH 3 + 5O 2 = 4NO + 6H 2 O,

2FeCl 2 + Cl 2 = 2 FeCl 3.

Nonmetals as reducing agents

1. All non-metals (except fluorine) exhibit reducing properties when interacting with oxygen:

S + O 2 = SO 2, 2H 2 + O 2 = 2H 2 O.

Oxygen in combination with fluorine can also exhibit a positive oxidation state, i.e., be a reducing agent. All other non-metals exhibit reducing properties. For example, chlorine does not combine directly with oxygen, but indirectly it is possible to obtain its oxides (Cl 2 O, ClO 2, Cl 2 O 2), in which chlorine exhibits a positive oxidation state. At high temperatures, nitrogen directly combines with oxygen and exhibits reducing properties. Sulfur reacts even more easily with oxygen.

2. Many non-metals exhibit reducing properties when interacting with complex substances:

ZnO + C = Zn + CO, S + 6HNO 3 conc = H 2 SO 4 + 6NO 2 + 2H 2 O.

3. There are also reactions in which the same nonmetal is both an oxidizing agent and a reducing agent:

Cl 2 + H 2 O = HCl + HClO.

4. Fluorine is the most typical non-metal, which is not characterized by reducing properties, i.e. the ability to donate electrons in chemical reactions.

Non-metal compounds
Nonmetals can form compounds with different intramolecular bonds.
Types of non-metal compounds
General formulas of hydrogen compounds according to groups of the periodic system of chemical elements are given in the table:

	RH 2	RH 3	RH 4	RH 3	H2R
Non-volatile hydrogen compounds				Volatile hydrogen compounds

With metals, hydrogen forms (with some exceptions) non-volatile compounds, which are solids of non-molecular structure. Therefore, their melting points are relatively high. With non-metals, hydrogen forms volatile compounds of molecular structure (for example, hydrogen fluoride HF, hydrogen sulfide H 2 S, ammonia NH 3, methane CH 4). Under normal conditions, these are gases or volatile liquids. When dissolved in water, hydrogen compounds of halogens, sulfur, selenium and tellurium form acids of the same formula as the hydrogen compounds themselves: HF, HCl, HBr, HI, H 2 S, H 2 Se, H 2 Te. When ammonia is dissolved in water, ammonia water is formed, usually denoted by the formula NH 4 OH and called ammonium hydroxide. It is also denoted by the formula NH 3 ∙H 2 O and is called ammonia hydrate.
With oxygen, nonmetals form acidic oxides. In some oxides they exhibit a maximum oxidation state equal to the group number (for example, SO 2, N 2 O 5), while in others it is lower (for example, SO 2, N 2 O 3). Acid oxides correspond to acids, and of the two oxygen acids of one nonmetal, the one in which it exhibits a higher oxidation state is stronger. For example, nitric acid HNO 3 is stronger than nitrous acid HNO 2, and sulfuric acid H 2 SO 4 is stronger than sulfurous acid H 2 SO 3.
Characteristics of oxygen compounds of nonmetals

1. The properties of higher oxides (i.e., oxides that contain an element of a given group with the highest oxidation state) gradually change from basic to acidic in periods from left to right.

2. In groups from top to bottom, the acidic properties of higher oxides gradually weaken. This can be judged by the properties of the acids corresponding to these oxides.

3. The increase in the acidic properties of higher oxides of the corresponding elements in periods from left to right is explained by a gradual increase in the positive charge of the ions of these elements.

4. In the main subgroups of the periodic system of chemical elements, the acidic properties of higher non-metal oxides decrease from top to bottom.