Non-metals. Chemical properties

There are only 16 non-metal chemical elements, but two of them, oxygen and silicon make up 76% of the mass of the earth's crust. Non-metals make up 98.5% of the mass of plants and 97.6% of the mass of man. From carbon, hydrogen, oxygen, sulfur, phosphorus and nitrogen, all the most important organic substances are composed, they are elements of life. Hydrogen and helium - the main elements of the Universe; all cosmic objects, including our Sun, are composed of them.

Non-metals are chemical elements whose atoms take electrons to complete the external energy level, while forming negatively charged ions. Almost all non-metals have relatively small radii and a large number of electrons at the external energy level from 4 to 7, they are characterized by high values \u200b\u200bof electronegativity and oxidative properties.

If we draw a diagonal from beryllium to astatine in the Periodic System, then non-metals will be on the right up diagonal, and metals on the bottom left, they also include elements of all secondary subgroups, lanthanides and actinides. Elements located near the diagonal, for example, beryllium, aluminum, titanium, germanium, antimony, have a dual character and belong to metalloids. The elements of group 18 are inert gases, have a completely complete outer electronic layer, they are sometimes referred to as non-metals, but formally, by physical signs.

Electronic configurations of valence electrons of non-metal elements are given in the table:

Patterns in changing the properties of non-metal elements

In the period with increasing nuclear charge (from left to right):

  • the radius of the atom decreases,
  • the number of electrons at the external energy level increases,
  • electronegativity increases
  • oxidizing properties are enhanced
  • non-metallic properties are enhanced.

In Group with increasing nuclear charge (top to bottom):

  • the radius of the atom increases,
  • the number of electrons at the external energy level does not change,
  • electronegativity decreases,
  • oxidizing properties weaken,
  • non-metallic properties weaken.

Thus, the more to the right and higher the element is in the Periodic system, the more pronounced are its non-metallic properties.

Non-metals in the main subgroup of the IV group of the Periodic system Mendeleev are carbon and silicon. At the external energy level of these elements are 4 electrons (ns 2 np 2). In its inorganic compounds, carbon has an oxidation state of +2 (in an unexcited state) and +4 (in an excited state). In organic compounds, the degree of oxidation of carbon can be any from –4 to +4.

For silicon, the most stable oxidation state is +4. Carbon and silicon form acid oxides of the general formula EO 2, as well as volatile hydrogen compounds of the general formula EN 4.

Non-metals in the V group of the main subgroup of the Periodic system Mendeleev are nitrogen, phosphorus, arsenic. At the external energy level of these elements are five electrons: ns 2 np 3. Nitrogen in its compounds can exhibit oxidation states of –3, –2, +1, +2, +3, +4, +5.
  Phosphorus is characterized by oxidation states of –3, +3, +5. Since the nitrogen atom does not have a d-sublevel, it cannot be pentavalent, but it is able to form a fourth covalent bond by the donor-acceptor mechanism. With an increase in the serial number within the subgroup, the radii of atoms and ions increase, and the ionization energy decreases. Non-metallic properties are weakened and metal is strengthened.
  With oxygen, the elements of the main subgroup of group V form higher oxides of the composition R 2 O 5. All of them are acid oxides. With hydrogen, nitrogen, phosphorus and arsenic form volatile gaseous compounds of the composition EN 3.

Non-metals of the main subgroup of the VI group of the Periodic system Mendeleev's are oxygen, sulfur, selenium and tellurium. The configuration of the external electronic level of these elements is ns 2 np 4. In their compounds, they exhibit the most characteristic oxidation states –2, +4, +6 (except oxygen). With an increase in the serial number within the subgroup, the ionization energy decreases, the sizes of atoms and ions increase, the non-metallic features of the elements weaken, and the metallic ones grow. Sulfur and selenium form higher oxides of the RO 3 type. These compounds are typical acidic oxides to which strong acids such as H 2 RO 4 correspond. Non-metals of the main subgroup of group VI are characterized by volatile hydrogen compounds with the general formula H 2 R. Moreover, the polarity and bond strength weaken from H 2 O to H 2 Te. All hydrogen compounds, except water, are gaseous substances. Aqueous solutions of H 2 S, H 2 Se, H 2 Te are weak acids.

Elements of group VII of the main subgroup - fluorine, chlorine, bromine, iodine are typical non-metals. The group name of these elements is halogens from the Greek halos - salt and genes - giving birth. The configuration of the external electronic level of these halogens is ns 2 np 5. The most characteristic oxidation state of halogens is -1. In addition, chlorine, bromine and iodine can exhibit oxidation states of + 3, + 5, + 7. Within each period, halogens are the most electronegative elements. Within the subgroup, when passing from fluorine to astatine, an increase in the radius of the atom occurs, nonmetallic properties decrease, a decrease in oxidizing properties and an increase in reduction properties. All halogens form simple substances - diatomic molecules Hal 2. Fluorine is the most electronegative of chemical elements. In all its compounds, it has oxidation states of –1. Higher halogen oxides (except fluorine) have the general formula R 2 O 7, are acid oxides. Strong acids of the general formula HRO 4 (R \u003d Cl, Br) correspond to them. Hydrogen compounds of halogens - hydrogen halides have the general formula HHal. Their aqueous solutions are acids, the strength of which increases from HF to HI. For halogens there is a pattern: each previous halogen is able to displace the next of its compounds with metals and hydrogen, for example: Cl 2 + 2KBr \u003d 2KCl + Br 2.

The periodic table of Dmitry Ivanovich Mendeleev is very convenient and universal in its use. It can be used to determine some characteristics of elements, and most surprisingly, to predict some properties of chemical elements that are still undiscovered, not discovered by scientists (for example, we know some properties of the alleged unbighexia, although it has not yet been discovered and synthesized).

These properties depend on the ability of the element.  give or attract electrons to yourself. It is important to remember one rule, metals - give away electrons, and non-metals - take. Accordingly, metallic properties are the ability of a particular chemical element to donate its electrons (from an external electron cloud) to another chemical element. For non-metals, everything is exactly the opposite. The easier non-metal accepts electrons, the higher its non-metallic properties.

Metals will never accept the electrons of another chemical element. This is characteristic of the following elements;

  • sodium;
  • potassium;
  • lithium;
  • france and so on.

With non-metals, things are similar. Fluorine more than any other non-metal exhibits its properties; it can only attract particles of another element to itself, but under no circumstances will it give up its own. It has the greatest non-metallic properties.. Oxygen (according to its characteristics) goes immediately after fluorine. Oxygen can form a compound with fluorine, giving up its electrons, but it takes away negative particles from other elements.

List of non-metals with the most pronounced characteristics:

  1. fluorine;
  2. oxygen;
  3. nitrogen;
  4. chlorine;
  5. bromine.

Non-metallic and metallic properties are explained by the fact that all chemicals seek to complete their energy level. To do this, at the last electronic level there should be 8 electrons. The fluorine atom on the last electron shell has 7 electrons, trying to complete it, it attracts another electron. The sodium atom on the outer shell has one electron, to get 8, it is easier to give 1 to it, and at the last level there will be 8 negatively charged particles.

Noble gases do not interact with other substances precisely because they have completed the energy level, they do not need to attract or donate electrons.

How metal properties change in a periodic system

The periodic table of Mendeleev consists of groups and periods. The periods are arranged horizontally so that the first period includes: lithium, beryllium, boron, carbon, nitrogen, oxygen, and so on. Chemical elements are arranged strictly by increasing the serial number.

The groups are arranged vertically so that the first group includes: lithium, sodium, potassium, copper, rubidium, silver, and so on. The group number indicates the number of negative particles at the external level of a particular chemical element. While the period number indicates the number of electronic clouds.

Metallic properties are enhanced in a row  from right to left or, in another way, weaken in the period. That is, magnesium has greater metallic properties than aluminum, but less than sodium. This is because in the period the number of electrons on the outer shell increases, therefore, it is more difficult for a chemical element to donate its electrons.

In a group, everything is the other way round; metallic properties are enhanced in a row from top to bottom. For example, potassium appears stronger than copper, but weaker than sodium. The explanation for this is very simple, the number of electron shells in the group increases, and the farther the electron is from the nucleus, the easier it is to give it to an element. The force of attraction between the nucleus of an atom and an electron in the first shell is greater than between the nucleus and an electron in 4 shell.

Compare the two elements - calcium and barium. Barium in the periodic system is lower than calcium. And this means that the electrons from the outer shell of calcium are located closer to the nucleus, therefore, they are better attracted than barium.

It is more difficult to compare elements that are in different groups and periods. Take, for example, calcium and rubidium. Rubidium will give better negative particles than calcium. Since it stands lower and to the left. But using only the periodic table, one cannot unequivocally answer this question by comparing magnesium and scandium (since one element is lower and to the right, and the other above and to the left). To compare these elements, special tables will be needed (for example, the electrochemical series of metal stresses).

How nonmetallic properties change in a periodic system

Non-metallic properties in the periodic table change exactly the opposite, rather than metallic. In fact, these two traits are antagonists.

Strengthen in the period (in a row from right to left). For example, sulfur is less able to attract electrons to itself than chlorine, but more than phosphorus. The explanation for this phenomenon is the same. The number of negatively charged particles on the outer layer increases, and therefore it is easier for an element to complete its energy level.

Non-metallic properties decrease in a row from top to bottom (in a group). For example, phosphorus is able to give back negatively charged particles more than nitrogen, but it is better able to attract than arsenic. Particles of phosphorus are attracted to the core better than particles of arsenic, which gives it the advantage of an oxidizing agent in reactions to lower and increase the degree of oxidation (redox reactions).

Compare, for example, sulfur and arsenic. Sulfur is higher and to the right, which means that it is easier for her to complete her energy level. Like metals, non-metals are difficult to compare if they are in different groups and periods. For example, chlorine and oxygen. One of these elements is higher and to the left, and the other is lower and to the right. For an answer, we have to turn to the table of electronegativity of non-metals, from which we see that oxygen attracts negative particles more easily than chlorine.

Periodic table  helps to find out not only the number of protons in an atom, atomic mass and serial number, but also helps to determine the properties of elements.

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In everyday life, a person interacts with many substances. All elements can be classified by physical and chemical qualities. In the article, we will examine how metals differ from non-metals, their properties and concept.

Definition of metal and its properties

Every day we deal with metals and it is no accident. Most elements of the periodic table are them. All of them have their own characteristics and properties.

As a rule, metals are elements that conduct heat and electricity well. Also, metals are very plastic, which allows them to change their shape by forging, and they also have a high coefficient of hardness. A distinctive feature of this element is the shine, which is called metallic. The properties of the metal are divided into two main fractions, such as:

  • Physical properties
  • Chemical properties.

How do metals differ from metals in physical characteristics? Physical properties include:

  • Color. Metals, as a rule, have a dense structure that does not let light through. And their color is determined by the reflection of light from its surface. So, metals in most cases have a color from gray to silver. But there are exceptions, such as copper, which has a red color, and gold, which has a yellow color.
  • Form condition, hardness and density. Metals themselves have a solid state of aggregation, but are capable of passing into liquid at high temperatures. So, metals melt at temperatures from 40 to 3400 degrees Celsius. But there are metals whose main state of aggregation is liquid. These elements include mercury.
  • Electrical conductivity. A feature is its decrease with increasing temperature of the substance.
  • Thermal conductivity and boiling / melting point.

How do metals differ from metals in chemical properties? In this group, there are:

How metals differ from each other

Many do not know how metals differ from metals. Their differences can be classified:

  • Metals differ in color, such as gold and copper.
  • Also, metals melt at different temperatures. Some metals, such as tin and lead, can be melted at home, but others require a higher temperature.
  • Between themselves, metals are divided into two groups: heavy and light. Heavy metals include those whose density is from 5 g / cm 3, light metals have a density of less than 5 g / cm 3. Light metals include lithium, which has a density of 0.2 g / cm 3; the site of the heaviest metal is divided between osmium and iridium. Their density is 22.6 g / cm 3.
  • Metals differ in plasticity and electrical conductivity. Some of them are very plastic. For example, from just 1 gram of gold, you can make a thin wire of 3.5 kilometers. It will be flexible and will not break. Repeat this with less ductile metal does not work.
  • Also, some metals conduct current better than others. The most electrically conductive metals are copper, silver and aluminum. They are most often used as conductive elements.

How do non-metals differ from metals

Non-metals are called elements that have non-metallic properties. What is the difference between metals and non-metals? Let's consider in more detail:


Wood is a plant material. Metal is the result of a natural chemical compound. What is the difference between wood and metal:


Semiconductors are non-metals that have some metallic properties. Metals and semiconductors have similarities in the fact that both are capable of conducting current.

But semiconductors have a distinctive feature, which is that their electrical conductivity can increase several times depending on external factors. Thus, the semiconductor conducts current better with increasing temperature. In metals, electrical conductivity decreases with increasing temperature. Also, the presence of impurities can affect the electrical conductivity. So, in metals, impurities lower electrical conductivity, and in semiconductors increase.

Semiconductors, unlike metals, can have positive and negative electrical conductivity. By themselves, semiconductors, by their ability to pass current through themselves, stand between the metal and elements that do not conduct current at all.

The difference between metal and steel

The fact is that metals are a whole group of elements that have metallic properties. This group also includes iron. Steel is nothing but an alloy of iron with elements included in a group of metals.

Most often, in addition to iron, steel also includes such elements of the periodic table as molybdenum, chromium and vanadium. Also carbon is a part of steel. With it, increase the strength of iron.

Thus, by varying the amount of carbon in the alloy, a very strong material can be obtained. But the stronger the steel, the more it becomes brittle. So, with a long dynamic load, steel easily breaks. Adding other impurities to it helps to achieve resistance to any effects.

So, the article examined how metals differ from metals and non-metals. The characteristics of all elements can be compared by chemical and physical properties. Every day, a person uses these elements and creates new substances to improve the quality of life.

General characteristics of non-metals.

Non-metals  - chemical elements that form simple bodies that do not have properties characteristic of metals. A qualitative characteristic of non-metals is electronegativity.

Electronegativity  - this is the ability to polarize a chemical bond, to attract common electronic pairs to itself.

Non-metals include 22 elements.

The position of non-metallic elements in the periodic system of chemical elements

1st period

2nd period

3rd period

4th period

5th period

6th period

As can be seen from the table, non-metallic elements are mainly located in the upper right part of the periodic system.

The structure of non-metal atoms

A characteristic feature of non-metals is a larger (compared to metals) number of electrons at the external energy level of their atoms. This determines their greater ability to attach additional electrons and manifest a higher oxidative activity than metals. Particularly strong oxidizing properties, i.e., the ability to attach electrons, are manifested by non-metals located in the 2nd and 3rd periods of groups VI-VII. If we compare the arrangement of electrons along the orbitals in the atoms of fluorine, chlorine, and other halogens, then we can also judge about their distinctive properties. The fluorine atom has no free orbitals. Therefore, fluorine atoms can only show valency I and oxidation state of 1. The strongest oxidizing agent is fluorine. In atoms of other halogens, for example, in a chlorine atom, there are free d-orbitals at the same energy level. Due to this, the pairing of electrons can occur in three different ways. In the first case, chlorine can manifest an oxidation state of +3 and form hydrochloric acid HClO 2, which corresponds to salts - chlorites, for example, potassium chlorite KClO 2. In the second case, chlorine can form compounds in which the oxidation state of chlorine is +5. Such compounds include perchloric acid HClO 3 and its salts - chlorates, for example potassium chlorate KClO 3 (Bertoletova salt). In the third case, chlorine exhibits an oxidation state of +7, for example, in perchloric acid HClO 4 and in its salts, perchlorates (in potassium perchlorate KClO 4).

Structures of non-metal molecules. Physical properties of non-metals

In a gaseous state at room temperature are:

    hydrogen is H 2;

    nitrogen - N 2;

    oxygen is O 2;

    fluorine - F 2;

    chlorine - CI 2.

And inert gases:

    helium - He;

    neon - Ne;

    argon - Ar;

    krypton - Kr;

    xenon - Xe;

    radon - Rn).

In the liquid - bromine - Br.

In solid:

  • carbon is C;

    silicon - Si;

    phosphorus - P;

  • arsenic - As;

    selenium - Se;

    tellurium - Te;

  • astatine - At.

The spectrum of colors is much richer in non-metals: red - for phosphorus, brown - for bromine, yellow - for sulfur, yellow-green - for chlorine, violet - for iodine vapors, etc.

The most typical non-metals have a molecular structure, while the less typical non-metals have a non-molecular structure. This explains the difference in their properties.

Composition and properties of simple substances - non-metals

Non-metals form both monatomic and diatomic molecules. TO monatomic non-metals include inert gases that practically do not react even with the most active substances. Inert gases are located in group VIII of the periodic system, and the chemical formulas of the corresponding simple substances are He, Ne, Ar, Kr, Xe and Rn.

Some non-metals form diatomic  molecules. These are H 2, F 2, Cl 2, Br 2, Cl 2 (elements of group VII of the periodic system), as well as oxygen O 2 and nitrogen N 2. Of triatomic molecules composed of gas ozone (O 3). For non-metal substances in the solid state, it is rather difficult to formulate the chemical formula. The carbon atoms in graphite are connected to each other in various ways. It is difficult to isolate a single molecule in the above structures. When writing the chemical formulas of such substances, as in the case of metals, the assumption is made that such substances consist only of atoms. Chemical formulas, in this case, are written without indices: C, Si, S, etc. Such simple substances as ozone and oxygen, having the same qualitative composition (both consist of the same element - oxygen), but differing in number atoms in a molecule have different properties. So, oxygen has no smell, while ozone has a pungent odor, which we perceive during a thunderstorm. The properties of solid non-metals, graphite and diamond, which also have the same qualitative composition, but different structures, differ sharply (graphite is brittle, diamond is hard). Thus, the properties of a substance are determined not only by its qualitative composition, but also by how many atoms are contained in the substance’s molecule and how they are related. Non-metals in the form of simple bodies are in a solid or gaseous state (excluding bromine - liquid). They do not have the physical properties inherent in metals. Solid non-metals do not have a characteristic gloss for metals, they are usually brittle, poorly conduct electric current and heat (with the exception of graphite). Crystalline boron B (like crystalline silicon) has a very high melting point (2075 ° C) and high hardness. The electric conductivity of boron increases significantly with increasing temperature, which makes it possible to widely use it in semiconductor technology. The addition of boron to steel and to alloys of aluminum, copper, nickel, etc. improves their mechanical properties. Borides (compounds of boron with some metals, for example, with titanium: TiB, TiB 2) are necessary in the manufacture of parts for jet engines, gas turbine blades. As can be seen from Scheme 1, carbon — C, silicon — Si, boron — B have a similar structure and have some common properties. As simple substances, they are found in two modifications - crystalline and amorphous. The crystalline modifications of these elements are very solid, with high melting points. Crystalline silicon has semiconductor properties. All these elements form compounds with metals - carbides, silicides and borides (CaC 2, Al 4 C 3, Fe 3 C, Mg 2 Si, TiB, TiB 2). Some of them have greater hardness, for example Fe 3 C, TiB. Calcium carbide is used to produce acetylene.

Chemical properties of non-metals

In accordance with the numerical values \u200b\u200bof the relative electronegativities, the oxidative abilities of nonmetals increase in the following order: Si, B, H, P, C, S, I, N, Cl, O, F.

Non-metals as oxidizing agents

The oxidizing properties of non-metals are manifested during their interaction:

    with metals: 2Na + Cl 2 \u003d 2NaCl;

    with hydrogen: H 2 + F 2 \u003d 2HF;

    with non-metals that have lower electronegativity: 2P + 5S \u003d P 2 S 5;

    with some complex substances: 4NH 3 + 5O 2 \u003d 4NO + 6H 2 O,

2FeCl 2 + Cl 2 \u003d 2 FeCl 3.

Non-metals as reducing agents

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

S + O 2 \u003d SO 2, 2H 2 + O 2 \u003d 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 directly combine with oxygen, but its oxides (Cl 2 O, ClO 2, Cl 2 O 2) can be obtained indirectly, in which chlorine exhibits a positive oxidation state. At high temperature, nitrogen directly combines with oxygen and exhibits reducing properties. Sulfur reacts even easier with oxygen.

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

ZnO + C \u003d Zn + CO, S + 6HNO 3 conc \u003d H 2 SO 4 + 6NO 2 + 2H 2 O.

    There are also such reactions in which the same non-metal is both an oxidizing agent and a reducing agent:

Cl 2 + H 2 O \u003d HCl + HClO.

    Fluorine is the most typical non-metal to which the reducing properties are uncharacteristic, i.e., the ability to donate electrons in chemical reactions.

Non-metal compounds

Non-metals can form compounds with different intramolecular bonds.

Types of non-metal compounds

The general formulas of hydrogen compounds for groups of the periodic system of chemical elements are given in the table:

Non-volatile hydrogen compounds

Volatile hydrogen compounds

With metals, hydrogen forms (with some exceptions) non-volatile compounds, which are solid substances 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, the 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.

Non-metals form acid oxides with oxygen. In some oxides, they exhibit a maximum oxidation state equal to the group number (for example, SO 2, N 2 O 5), and others a lower degree (for example, SO 2, N 2 O 3). Acid oxides correspond to acids, and of the two oxygen acids of one non-metal, the one in which it exhibits a higher degree of oxidation is stronger. For example, nitric acid HNO 3 is stronger than nitrous HNO 2, and sulfuric acid H 2 SO 4 is stronger than sulfur dioxide H 2 SO 3.

Characteristics of non-metal oxygen compounds

    The properties of higher oxides (i.e., oxides, which include an element of this group with a high degree of oxidation) in the periods from left to right gradually change from basic to acidic.

    In groups from top to bottom, the acid properties of higher oxides are gradually weakening. This can be judged by the properties of the acids corresponding to these oxides.

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

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

Halogens.

The structure of halogen atoms

Halogens include elements of group VIII of the periodic system, the atoms of these elements contain seven electrons at the external energy level, and until its completion they lack only one electron, therefore halogens exhibit bright oxidizing properties. In the subgroup, with an increase in the serial number, these properties decrease due to an increase in the radius of atoms: from fluorine to astatine, and, accordingly, their reduction properties increase. The value of the relative electronegativity of halogens decreases similarly. As the most electronegative element, fluorine in compounds with other elements exhibits a constant oxidation state -1 . Other halogens can exhibit both this oxidation state in compounds with metals, hydrogen and less electronegative elements, as well as positive odd oxidation states from +1   before +7   in compounds with more electronegative elements: oxygen, fluorine.

Simple substances halogens and their properties

Chlorine, bromine and iodine in glass vessels

Characterizing simple substances - halogens, it is necessary to recall the basic theoretical information about the types of chemical bonds and the crystalline structure of the substance. In diatomic halogen molecules, atoms are linked by a covalent non-polar bond G · G  or G ― G  and have a molecular crystal lattice.

Under normal conditions F 2   - bright yellow gas with an orange tint, Cl 2   - yellow-green poisonous gas with a characteristic suffocating odor, Br 2   - volatile brown liquid (bromine vapor is highly toxic, bromine burns are very painful and do not heal for a long time), and I 2   - crystalline solid, capable of sublimation. In a row F 2,   Cl 2 , Br 2 , I 2   - the density of simple substances increases, and the color intensity increases. Consequently, the same pattern is manifested in the change in the properties of atoms and simple substances - halogens: with an increase in the serial number, non-metallic properties weaken, and metallic ones increase.

Chemical properties of halogens

    The interaction of halogens with metals with the formation of halides:

2Na + I 2 ―― 2Na +1 I -1 (sodium iodide);

2Al + 3I 2 \u003d 2Al +3 I 3 -1 (aluminum iodide);

2Al + 3Br 2 \u003d 2Al +3 Br 3 -1 (aluminum bromide).

During reactions of metals of sub-groups (transition metals) with halogens, halides with a high degree of metal oxidation are formed, for example:

2Fe + 3Cl 2 \u003d 2FeCl 3,

but 2CHl + Fe \u003d FeCl 2 + H 2.

    The interaction of halogens with hydrogen with the formation of hydrogen halides (the type of bond is covalent polar, the type of lattice is molecular). Comparison of the rate of chemical reactions of different halogens with hydrogen allows us to repeat its dependence on the nature of the reacting substances. So, fluorine has such a high reaction rate that it interacts with hydrogen with an explosion, even in the dark. The reaction of chlorine with hydrogen under normal conditions is slow and only when ignited or illuminated does its speed increase many times (an explosion occurs). Bromine and iodine interact even more slowly with hydrogen, with the latter reaction already acquiring an endothermic character:

Only fluorine interacts with hydrogen irreversibly, other halogens, depending on conditions, can give a reversible reaction.

Aqueous solutions of hydrogen halides are acids: HF - hydrofluoric (hydrofluoric), HCl - hydrochloric (hydrochloric), HBr - hydrobromic, HI - hydroiodic.

    Halogens interact with water:

2F 2 + 2H 2 O \u003d 4HF + O 2

Water in fluorine burns, oxygen is not a cause, but a consequence of combustion, acting in an unusual role for it as a reducing agent.

    To characterize the ability of some halogens (not halogen atoms, but simple substances) to displace others from solutions of their compounds, you can use the “activity series” of halogens, which is written as follows:

F 2\u003e Cl 2\u003e Br 2\u003e I 2,

i.e., oxidizing properties are reduced.

So, chlorine displaces bromine and iodine (but not fluorine), and bromine is able to displace only iodine from solutions of the corresponding salts:

2NaBr + Cl 2 \u003d 2NaCl + Br 2

2KI + Br 2 \u003d 2KBr + I 2.

Biological significance and application of halogens

Fluorineplays a very important role in the life of plants, animals and humans. Without fluoride, the development of the bone skeleton and especially the teeth is impossible. The fluorine content in the bones is 80-100 mg per 100 g of dry matter. In enamel, fluorine is present in the form of a compound Ca 4 F 2 (PO 4) 2 and gives it hardness and whiteness. With a lack of fluoride in the human body, damage to the dental tissue (caries) occurs, and its excess contributes to the disease of the teeth with fluorosis. The daily human need for fluoride is 2-3 mg. Chlorine(chlorine ion) is more important for the life of animals and humans than for plants. It is part of the kidneys, lungs, spleen, blood, saliva, cartilage, hair. Chlorine ions regulate the blood buffer system. Sodium chloride is an integral part of blood plasma and cerebrospinal fluid and is involved in the regulation of water metabolism in the body. Free hydrochloric acid is part of the gastric juice of all mammals and is actively involved in digestion. A healthy person contains 0.2-0.3% hydrochloric acid in the stomach. A lack of chlorine in the body leads to tachycardia, a decrease in blood pressure, and convulsions. A sufficient amount of chlorine is found in vegetables such as celery, radishes, cucumbers, white cabbage, dill, peppers, onions, and artichokes. Bromineit is also among the necessary trace elements and most of all it is found in the pituitary gland, blood. Thyroid, adrenal glands. Bromides in small doses (0.1-0.3 adults) do not positively act on the central nervous system as enhancers of inhibition in the cerebral cortex. In nature, bromides accumulate in plants such as rye, wheat, barley, potatoes, carrots, cherries, apples. A lot of bromine is found in Dutch cheese. Iodine  in the human body begins to accumulate in the womb. The human thyroid hormone - thyroxine - contains 60% of bound iodine. This hormone with blood flow enters the liver, kidneys, mammary glands, and the gastrointestinal tract. A lack of iodine in the human body causes diseases such as endemic goiter and cretinism, in which growth slows down and mental retardation develops. In combination with other elements, iodine promotes the growth and fatness of animals, improves their health and fertility. The main suppliers of iodine for humans are cereals, eggplant, beans, white cabbage and cauliflower, potatoes, onions, carrots, cucumbers, pumpkin, lettuce, sea kale, squid.

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  • The main educational program of primary and secondary general education "secondary school number 10"

    Basic educational program

    Are specified are common  main goals of common  education taking into account the specifics of the subject; 2) overallcharacterization  training ... in electrolyte solutions. Variety of substances Totalcharacteristicnon-metals  based on their position in the periodic ...

    • I. Elements.  Non-metals form p-elements, as well as hydrogen and helium, which are s-elements. In a long-period table p-elements forming non-metals are located to the right and above the conditional boundary B - At.

    II. Atoms  Non-metal atoms are small (orbital radius less than 0.1 nm). Most of them have four to eight valence electrons (they are external), but the hydrogen atom has one, the helium atom has two, and the boron atom has three valence electrons. Non-metal atoms are relatively easy to attach foreign electrons (but not more than three). Non-metal atoms do not have a tendency to give electrons.

    At atoms of non-metal elements in the period with increasing serial number

    • core charge increases;
    • the radii of atoms are reduced;
    • the number of electrons on the outer layer increases;
    • the number of valence electrons increases;
    • electronegativity increases;
    • oxidizing (non-metallic) properties are enhanced (except for elements of group VIIIA).

    At atoms of non-metal elements in a subgroup (in a long-period table - in a group) with increasing serial number

    • core charge increases;
    • the radius of the atom increases;
    • electronegativity decreases;
    • the number of valence electrons does not change;
    • the number of external electrons does not change (with the exception of hydrogen and helium);
    • oxidative (non-metallic) properties are weakened (except for elements of group VIIIA).

    III. Simple substances.  Most non-metals are simple substances in which atoms are linked by covalent bonds; there are no chemical bonds in noble gases. Among non-metals there are both molecular and non-molecular substances. All this leads to the fact that there are no physical properties characteristic of all non-metals.

    Molecular non-metals: H 2, N 2, P 4 (white phosphorus), As 4, O 2, O 3, S 8, F 2, Cl 2, Br 2, I 2. Noble gases (He, Ne, Ar, Kr, Kx, Rn), whose atoms are like “monatomic molecules,” can also be attributed to them.

    At room temperature, hydrogen, nitrogen, oxygen, ozone, fluorine and chlorine are gases; bromine is a liquid; phosphorus, arsenic, sulfur and iodine are solids.

    Non-molecular non-metals: B (several allotropic modifications), C (graphite), C (diamond), Si, Ge, P (red), P (black), As, Se, Te. All of them are solids, silicon, germanium, selenium and some others have semiconductor properties.

    IV. Chemical properties.  Typical for most non-metals are oxidative properties. As oxidizing agents, they react with metals:


      with complex substances:

    With complex substances:

    H 2 + HCHO \u003d CH 3 OH 6P + 5KClO 3 \u003d 5KCl + 3P 2 O 5

    V. Hydrogen compounds.  All non-metals (except elements of noble gases) form molecular hydrogen compounds, and carbon and boron are very numerous. The simplest hydrogen compounds:

    All he gases except water. Substances in bold in aqueous solution are strong acids.

    In the group, with an increase in the serial number, their stability decreases, and the recovery activity increases.

    In the period with an increase in the serial number, the acidic properties of their solutions intensify; in the group, these properties weaken.

    VI. Oxides and hydroxides.  All non-metal oxides are acidic or non-salt forming. Non-salt forming oxides: CO, SiO, N 2 O, NO.

    The following acids correspond to higher non-metal oxides (strong acids are shown in bold)

    In the period with increasing serial number, the strength of higher acids increases. In groups, there is no pronounced dependence.

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