Non-metals are chemical elements that have typical non-metallic properties and are located in the upper right corner of the Periodic Table. What properties are inherent in these elements, and with what do nonmetals react?

Non-metals: general characteristics

Nonmetals differ from metals in that they have more electrons in their outer energy level. Therefore, their oxidizing properties are more pronounced than those of metals. Non-metals are characterized by high electronegativity values ​​and high reduction potential.

Non-metals include chemical elements that are in a gaseous, liquid or solid state of aggregation. So, for example, nitrogen, oxygen, fluorine, chlorine, hydrogen are gases; iodine, sulfur, phosphorus - solid; bromine is a liquid (at room temperature). There are 22 non-metals in total.

Rice. 1. Non-metals - gases, solids, liquids.

With an increase in the charge of the atomic nucleus, a pattern of changes in the properties of chemical elements from metallic to non-metallic is observed.

Chemical properties of non-metals

Hydrogen properties of non-metals are mainly volatile compounds, which in aqueous solutions are acidic. They have molecular structures as well as a covalent polar bond. Some, such as water, ammonia, or hydrogen fluoride, form hydrogen bonds. Compounds are formed by the direct interaction of non-metals with hydrogen. Example:

S + H 2 \u003d H 2 S (up to 350 degrees, the balance is shifted to the right)

All hydrogen compounds have reducing properties, with their reducing power increasing from right to left in a period and from top to bottom in a group. So, hydrogen sulfide burns with a large amount of oxygen:

2H 2 S + 3O 3 \u003d 2SO 2 + 2H 2 O + 1158 kJ.

Oxidation can go in a different way. So, already in air, an aqueous solution of hydrogen sulfide becomes cloudy as a result of the formation of sulfur:

H 2 S + 3O 2 \u003d 2S + 2H 2 O

Compounds of non-metals with oxygen, as a rule, are acid oxides, which correspond to oxygen-containing acids (oxo acids). The structure of oxides of typical non-metals is molecular.

The higher the oxidation state of the non-metal, the stronger the corresponding oxygen-containing acid. So, chlorine does not directly interact with oxygen, but forms a number of oxo acids, which correspond to oxides, anhydrides of these acids.

The best known are such salts of these acids as bleach CaOCl 2 (mixed salt of hypochlorous and hydrochloric acids), bertolet salt KClO 3 (potassium chlorate).

Nitrogen in oxides exhibits positive oxidation states +1, +2, +3, +4, +5. The first two oxides N 2 O and NO are non-salt-forming and are gases. N 2 O 3 (nitric oxide III) - is an anhydride of nitrous acid HNO 2. Nitric oxide IV - brown gas NO 2 - a gas that dissolves well in water, forming two acids. This process can be expressed by the equation:

2NO 2 + H 2 O \u003d HNO 3 (nitric acid) + HNO 2 (nitrous acid) - redox disproportionation reaction

Rice. 2. Nitrous acid.

Nitric acid anhydride N 2 O 5 is a white crystalline substance that is easily soluble in water. Example:

N 2 O 5 + H 2 O \u003d 2HNO 3

Salts of nitric acid are called saltpeters, they are soluble in water. Salts of potassium, calcium, sodium are used to produce nitrogen fertilizers.

Phosphorus forms oxides, showing oxidation states +3 and +5. The most stable oxide is phosphoric anhydride P 2 O 5 , which forms a molecular lattice with P 4 O 10 dimers at its nodes. Salts of phosphoric acid are used as phosphate fertilizers, for example, ammophos NH 4 H 2 PO 4 (ammonium dihydrogen phosphate).

Table of arrangement of non-metals

Group	I	III	IV	V	VI	VII	VIII
First period	H						He
Second period		B	C	N	O	F	Ne
Third period			Si	P	S	Cl	Ar
The fourth period				As	Se	Br	kr
Fifth period					Te	I	Xe
Sixth period						At	Rn

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, resulting in the formation of insoluble bases - 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.

Non-metals are elements that differ significantly in physical and chemical properties from metals. The reason for their differences could be explained in detail only at the end of the 19th century, after the discovery of the electronic structure of the atom. What is the peculiarity of non-metals? What qualities are characteristic of their day? Let's figure it out.

Non-metals - what is it?

The approach to separating elements into metals and non-metals has long existed in the scientific community. The first elements in the periodic table of Mendeleev usually include 94 elements. Mendeleev's non-metals include 22 elements. In they occupy the upper right corner.

In their free form, non-metals are simple substances, the main feature of which is the absence of characteristic metallic properties. They can be in all states of aggregation. So, iodine, phosphorus, sulfur, carbon are found in the form of solid substances. The gaseous state is characteristic of oxygen, nitrogen, fluorine, etc. Only bromine is a liquid.

In nature, non-metal elements can exist both in the form of simple substances and in the form of compounds. Sulfur, nitrogen, oxygen are found in unbound form. In compounds, they form borates, phosphates, etc. In this form, they are present in minerals, water, rocks.

Difference from metals

Non-metals are elements that differ from metals in appearance, structure and chemical properties. They have a large number of unpaired electrons at the external level, which means they are more active in oxidative reactions and more easily attach additional electrons to themselves.

A characteristic difference between the elements is observed in the structure of the crystal lattice. In metals, it is metallic. In non-metals, it can be of two types: atomic and molecular. The atomic lattice gives substances hardness and increases the melting point; it is characteristic of silicon, boron, and germanium. Chlorine, sulfur, oxygen have a molecular lattice. It gives them volatility and a little hardness.

The internal structure of elements determines their physical properties. Metals have a characteristic luster, good conductivity of current and heat. They are hard, ductile, malleable, have a small range of colors (black, shades of gray, sometimes yellowish).

Non-metals are liquid, gaseous or non-lustrous and malleable. Their colors vary greatly and can be red, black, gray, yellow, etc. Almost all non-metals are poor conductors of current (except carbon) and heat (except black phosphorus and carbon).

Chemical properties of non-metals

In chemical reactions, non-metals can play the role of both oxidizing and reducing agents. When interacting with metals, they take on electrons, thus exhibiting oxidizing properties.

Interacting with other non-metals, they behave differently. In such reactions, the less electronegative element acts as a reducing agent, while the more electronegative element acts as an oxidizing agent.

With oxygen, almost all (except fluorine) non-metals act as reducing agents. When interacting with hydrogen, many are oxidizing agents, subsequently forming volatile compounds.

Some non-metal elements have the ability to form several simple substances or modifications. This phenomenon is called allotropy. For example, carbon exists in the form of graphite, diamond, carbine, and other modifications. Oxygen has two of them - ozone and oxygen itself. Phosphorus comes in red, black, white, and metallic.

Nonmetals in nature

Non-metals are found everywhere in varying amounts. They are part of the earth's crust, are part of the atmosphere, hydrosphere, are present in the universe and in living organisms. In outer space, the most common are hydrogen and helium.

On Earth, the situation is quite different. The most important constituents of the earth's crust are oxygen and silicon. They make up more than 75% of its mass. But the smallest amount falls on iodine and bromine.

In the composition of sea water, oxygen accounts for 85.80%, and hydrogen - 10.67%. Its composition also includes chlorine, sulfur, boron, bromine, carbon, fluorine and silicon. Nitrogen (78%) and oxygen (21%) dominate in the composition of the atmosphere.

Non-metals such as carbon, hydrogen, phosphorus, sulfur, oxygen and nitrogen are important organic substances. They support the vital activity of all living beings on our planet, including humans.

If we draw a diagonal from beryllium to astatine in the periodic table of elements of D.I. Mendeleev, then there will be metal elements on the diagonal at the bottom left (they also include elements of secondary subgroups, highlighted in blue), and on the top right - non-metal elements (highlighted in yellow). Elements located near the diagonal - semimetals or metalloids (B, Si, Ge, Sb, etc.) have a dual character (highlighted in pink).

As can be seen from the figure, the vast majority of elements are metals.

By their chemical nature, metals are chemical elements whose atoms donate electrons from the outer or pre-outer energy levels, thus forming positively charged ions.

Almost all metals have relatively large radii and a small number of electrons (from 1 to 3) at the external energy level. Metals are characterized by low electronegativity values ​​and reducing properties.

The most typical metals are located at the beginning of periods (starting from the second), further from left to right, the metallic properties weaken. In a group from top to bottom, metallic properties are enhanced, because the radius of atoms increases (due to an increase in the number of energy levels). This leads to a decrease in the electronegativity (the ability to attract electrons) of the elements and an increase in the reducing properties (the ability to donate electrons to other atoms in chemical reactions).

typical metals are s-elements (elements of the IA group from Li to Fr. elements of the PA group from Mg to Ra). The general electronic formula of their atoms is ns 1-2. They are characterized by oxidation states + I and + II, respectively.

The small number of electrons (1-2) in the outer energy level of typical metal atoms suggests easy loss of these electrons and the manifestation of strong reducing properties, which reflect low electronegativity values. This implies the limited chemical properties and methods for obtaining typical metals.

A characteristic feature of typical metals is the tendency of their atoms to form cations and ionic chemical bonds with non-metal atoms. Compounds of typical metals with non-metals are ionic crystals "metal cation anion of non-metal", for example, K + Br -, Ca 2+ O 2-. Typical metal cations are also included in compounds with complex anions - hydroxides and salts, for example, Mg 2+ (OH -) 2, (Li +) 2CO 3 2-.

The A-group metals forming the amphoteric diagonal in the Be-Al-Ge-Sb-Po Periodic System, as well as the metals adjacent to them (Ga, In, Tl, Sn, Pb, Bi) do not exhibit typically metallic properties. The general electronic formula of their atoms ns 2 np 0-4 implies a greater variety of oxidation states, a greater ability to retain their own electrons, a gradual decrease in their reducing ability and the appearance of an oxidizing ability, especially in high oxidation states (typical examples are compounds Tl III, Pb IV, Bi v). A similar chemical behavior is also characteristic of most (d-elements, i.e., elements of the B-groups of the Periodic Table (typical examples are the amphoteric elements Cr and Zn).

This manifestation of duality (amphoteric) properties, both metallic (basic) and non-metallic, is due to the nature of the chemical bond. In the solid state, compounds of atypical metals with non-metals contain predominantly covalent bonds (but less strong than bonds between non-metals). In solution, these bonds are easily broken, and the compounds dissociate into ions (completely or partially). For example, gallium metal consists of Ga 2 molecules, in the solid state aluminum and mercury (II) chlorides AlCl 3 and HgCl 2 contain strongly covalent bonds, but in a solution AlCl 3 dissociates almost completely, and HgCl 2 - to a very small extent (and even then into HgCl + and Cl - ions).

General physical properties of metals

Due to the presence of free electrons ("electron gas") in the crystal lattice, all metals exhibit the following characteristic general properties:

1) Plastic- the ability to easily change shape, stretch into a wire, roll into thin sheets.

2) metallic luster and opacity. This is due to the interaction of free electrons with light incident on the metal.

3) Electrical conductivity. It is explained by the directed movement of free electrons from the negative to the positive pole under the influence of a small potential difference. When heated, the electrical conductivity decreases, because. as the temperature rises, vibrations of atoms and ions in the nodes of the crystal lattice increase, which makes it difficult for the directed movement of the "electron gas".

4) Thermal conductivity. It is due to the high mobility of free electrons, due to which the temperature is quickly equalized by the mass of the metal. The highest thermal conductivity is in bismuth and mercury.

5) Hardness. The hardest is chrome (cuts glass); the softest - alkali metals - potassium, sodium, rubidium and cesium - are cut with a knife.

6) Density. It is the smaller, the smaller the atomic mass of the metal and the larger the radius of the atom. The lightest is lithium (ρ=0.53 g/cm3); the heaviest is osmium (ρ=22.6 g/cm3). Metals having a density less than 5 g/cm3 are considered "light metals".

7) Melting and boiling points. The most fusible metal is mercury (m.p. = -39°C), the most refractory metal is tungsten (t°m. = 3390°C). Metals with t°pl. above 1000°C are considered refractory, below - low melting point.

General chemical properties of metals

Strong reducing agents: Me 0 – nē → Me n +

A number of stresses characterize the comparative activity of metals in redox reactions in aqueous solutions.

I. Reactions of metals with non-metals

1) With oxygen:
2Mg + O 2 → 2MgO

2) With sulfur:
Hg + S → HgS

3) With halogens:
Ni + Cl 2 – t° → NiCl 2

4) With nitrogen:
3Ca + N 2 – t° → Ca 3 N 2

5) With phosphorus:
3Ca + 2P – t° → Ca 3 P 2

6) With hydrogen (only alkali and alkaline earth metals react):
2Li + H 2 → 2LiH

Ca + H 2 → CaH 2

II. Reactions of metals with acids

1) Metals standing in the electrochemical series of voltages up to H reduce non-oxidizing acids to hydrogen:

Mg + 2HCl → MgCl 2 + H 2

2Al+ 6HCl → 2AlCl 3 + 3H 2

6Na + 2H 3 PO 4 → 2Na 3 PO 4 + 3H 2

2) With oxidizing acids:

In the interaction of nitric acid of any concentration and concentrated sulfuric acid with metals hydrogen is never released!

Zn + 2H 2 SO 4 (K) → ZnSO 4 + SO 2 + 2H 2 O

4Zn + 5H 2 SO 4(K) → 4ZnSO 4 + H 2 S + 4H 2 O

3Zn + 4H 2 SO 4(K) → 3ZnSO 4 + S + 4H 2 O

2H 2 SO 4 (c) + Cu → Cu SO 4 + SO 2 + 2H 2 O

10HNO 3 + 4Mg → 4Mg(NO 3) 2 + NH 4 NO 3 + 3H 2 O

4HNO 3 (c) + Сu → Сu (NO 3) 2 + 2NO 2 + 2H 2 O

III. Interaction of metals with water

1) Active (alkali and alkaline earth metals) form a soluble base (alkali) and hydrogen:

2Na + 2H 2 O → 2NaOH + H 2

Ca+ 2H 2 O → Ca(OH) 2 + H 2

2) Metals of medium activity are oxidized by water when heated to oxide:

Zn + H 2 O – t° → ZnO + H 2

3) Inactive (Au, Ag, Pt) - do not react.

IV. Displacement by more active metals of less active metals from solutions of their salts:

Cu + HgCl 2 → Hg + CuCl 2

Fe+ CuSO 4 → Cu+ FeSO 4

In industry, not pure metals are often used, but their mixtures - alloys in which the beneficial properties of one metal are complemented by the beneficial properties of another. So, copper has a low hardness and is of little use for the manufacture of machine parts, while alloys of copper with zinc ( brass) are already quite hard and are widely used in mechanical engineering. Aluminum has high ductility and sufficient lightness (low density), but is too soft. On its basis, an alloy with magnesium, copper and manganese is prepared - duralumin (duralumin), which, without losing the useful properties of aluminum, acquires high hardness and becomes suitable in the aircraft industry. Alloys of iron with carbon (and additions of other metals) are widely known cast iron and steel.

Metals in free form are reducing agents. However, the reactivity of some metals is low due to the fact that they are covered with surface oxide film, to varying degrees resistant to the action of such chemical reagents as water, solutions of acids and alkalis.

For example, lead is always covered with an oxide film; its transition into solution requires not only exposure to a reagent (for example, dilute nitric acid), but also heating. The oxide film on aluminum prevents its reaction with water, but is destroyed under the action of acids and alkalis. Loose oxide film (rust), formed on the surface of iron in moist air, does not interfere with the further oxidation of iron.

Under the influence concentrated acids are formed on metals sustainable oxide film. This phenomenon is called passivation. So, in concentrated sulfuric acid passivated (and then do not react with acid) such metals as Be, Bi, Co, Fe, Mg and Nb, and in concentrated nitric acid - metals A1, Be, Bi, Co, Cr, Fe, Nb, Ni, Pb , Th and U.

When interacting with oxidizing agents in acidic solutions, most metals turn into cations, the charge of which is determined by the stable oxidation state of a given element in compounds (Na +, Ca 2+, A1 3+, Fe 2+ and Fe 3+)

The reducing activity of metals in an acidic solution is transmitted by a series of stresses. Most metals are converted into a solution with hydrochloric and dilute sulfuric acids, but Cu, Ag and Hg - only with sulfuric (concentrated) and nitric acids, and Pt and Au - with "aqua regia".

Corrosion of metals

An undesirable chemical property of metals is their, i.e., active destruction (oxidation) upon contact with water and under the influence of oxygen dissolved in it (oxygen corrosion). For example, the corrosion of iron products in water is widely known, as a result of which rust is formed, and the products crumble into powder.

Corrosion of metals proceeds in water also due to the presence of dissolved CO 2 and SO 2 gases; an acidic environment is created, and H + cations are displaced by active metals in the form of hydrogen H 2 ( hydrogen corrosion).

The point of contact between two dissimilar metals can be especially corrosive ( contact corrosion). Between one metal, such as Fe, and another metal, such as Sn or Cu, placed in water, a galvanic couple appears. The flow of electrons goes from the more active metal, which is to the left in the series of voltages (Re), to the less active metal (Sn, Cu), and the more active metal is destroyed (corrodes).

It is because of this that the tinned surface of cans (tin-plated iron) rusts when stored in a humid atmosphere and carelessly handled (iron quickly collapses after even a small scratch appears, allowing contact of iron with moisture). On the contrary, the galvanized surface of an iron bucket does not rust for a long time, because even if there are scratches, it is not iron that corrodes, but zinc (a more active metal than iron).

Corrosion resistance for a given metal is enhanced when it is coated with a more active metal or when they are fused; for example, coating iron with chromium or making an alloy of iron with chromium eliminates the corrosion of iron. Chrome-plated iron and steel containing chromium ( stainless steel) have high corrosion resistance.

electrometallurgy, i.e., obtaining metals by electrolysis of melts (for the most active metals) or salt solutions;

pyrometallurgy, i.e., the recovery of metals from ores at high temperature (for example, the production of iron in the blast furnace process);

hydrometallurgy, i.e., the isolation of metals from solutions of their salts by more active metals (for example, the production of copper from a CuSO 4 solution by the action of zinc, iron or aluminum).

Native metals are sometimes found in nature (typical examples are Ag, Au, Pt, Hg), but more often metals are in the form of compounds ( metal ores). By prevalence in the earth's crust, metals are different: from the most common - Al, Na, Ca, Fe, Mg, K, Ti) to the rarest - Bi, In, Ag, Au, Pt, Re.

The position of non-metal elements in the Periodic system of chemical elements D.I. Mendeleev

Elements-non-metals:

s-element - hydrogen;

p-elements of the 3rd group - 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.

Elements of the 8th group - 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 able to donate 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 (give away). Intermediate redox properties have nitrogen, sulfur, iodine.

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 \u003d 2NH 3, H 2 + Br 2 \u003d 2HBr;

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

S + O 2 \u003d SO 2, 4P + 5O 2 \u003d 2P 2 O 5;

when interacting with fluorine fluorine is an oxidizing agent, and oxygen is a reducing agent: 2F 2 + O 2 \u003d 2OF 2;

non-metals interact between themselves , a more electronegative metal plays the role of an oxidizing agent, a less electronegative one plays the role of a reducing agent: S + 3F 2 \u003d SF 6, C + 2Cl 2 \u003d CCl 4.

Halogens (Group 7)

Chemical properties of halogens.

OXYGEN-CONTAINING CHLORINE ACID

· Hypochlorous acid HCl +1 O salt - hypo chlorites

It 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 the light, releasing atomic oxygen HClO = HCl + O

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

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

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

Chemical properties

1.HClO2 + KOH = KClO2 + H2O

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

Perchloric acid HCl O3 (HClO3 - Strong acid and strong oxidizing agent; chloric acid salts - chlorates)

KClO 3 - Berthollet's salt; it is obtained by passing chlorine through a heated (40 ° C) KOH solution:

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

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

4KClO 3 = KCl + 3KClO 4 (no catalyst)

2KClO 3 \u003d 2KCl + 3O 2 (MnO 2 catalyst)

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

Obtaining KClO4 + H2SO4 = KHSO4 + HClO4

Chemical properties

1) Interacts 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.

At the outer energy level, chalcogens have 6 electrons. Before the completion of the external energy level, atoms lack 2 electrons, so they add electrons and show an oxidation state of -2 in their compounds.

Sulfur, selenium and tellurium atoms in their compounds with more electronegative elements exhibit positive oxidation states +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%.

Getting 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 takes place in the presence of a catalyst - manganese oxide IV

In the laboratory, the decomposition reaction of potassium permanganate - KMnO 4 - "potassium permanganate" is also used.

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

2KClO 3 \u003d 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 is formed.

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

· Oxygen reacts with non-metals.

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

Almost all reactions with oxygen are exothermic (that is, accompanied by the release of heat). An 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 \u003d CO 2 + 2H 2 O 2H 2 S + 3O 2 \u003d 2SO 2 + 2H 2 O

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