General characteristics of non-metals.

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

Electronegativity- This is the ability to polarize a chemical bond, to pull off common electron pairs.

Non-metals include 22 elements.

Position of non-metallic elements in the periodic table of chemical elements

1st period
2nd period
3rd period
4th period
5th period
6th period

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

Atomic structure of non-metals

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, that is, the ability to attach electrons, are shown by non-metals located in the 2nd and 3rd periods of VI-VII groups. If we compare the arrangement of electrons in orbitals in the atoms of fluorine, chlorine and other halogens, then we can judge about their distinctive properties. The fluorine atom has no free orbitals. Therefore, fluorine atoms can show only valence I and oxidation state - 1. The strongest oxidizing agent is fluorine... In the atoms of other halogens, for example, in the chlorine atom, there are free d-orbitals at the same energy level. Due to this, the steaming of electrons can occur in three different ways. In the first case, chlorine can exhibit an oxidation state of +3 and form chlorous 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 chloric acid HClO 3 and its salts - chlorates, for example potassium chlorate KClO 3 (Berthollet's 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).

Molecular structures of non-metals. Physical properties of non-metals

In a gaseous state at room temperature are:

hydrogen - H 2;

nitrogen - N 2;

oxygen - O 2;

fluorine - F 2;

chlorine - CI 2.

And inert gases:

helium - He;

neon - Ne;

argon - Ar;

krypton - Kr;

xenon - Xe;

radon - Rn).

In liquid - bromine - Br.

In solid:

• carbon - C;

  silicon - Si;

  phosphorus - P;

• arsenic - As;

  selenium - Se;

  tellurium - Te;

• astatine - At.

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

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

Composition and properties of simple substances - non-metals

Non-metals form both monoatomic 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 as follows: 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. From triatomic molecules consist of ozone gas (O 3). For substances of non-metals in a solid state, it is rather difficult to compose a chemical formula. The carbon atoms in graphite are connected to each other in different ways. It is difficult to isolate a single molecule in the given 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, they are written without indices: C, Si, S, etc. Such simple substances as ozone and oxygen, which have the same qualitative composition (both consist of the same element - oxygen), but differ in the number of atoms in molecule have different properties. So, oxygen has no smell, while ozone has a pungent smell that we feel during a thunderstorm. The properties of hard 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 a molecule of a substance and how they are related to each other. 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 of metals. Solid non-metals do not have the luster characteristic of metals, they are usually fragile, poorly conducting 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 electrical conductivity of boron increases greatly 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 them mechanical properties... Borides (compounds of boron with some metals, for example with titanium: TiB, TiB 2) are necessary in the manufacture of parts 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. Crystalline modifications of these elements are very hard, with high melting points. Crystalline silicon has semiconducting 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 higher hardness, for example Fe 3 C, TiB. Calcium carbide is used to produce acetylene.

Chemical properties non-metals

In accordance with the numerical values ​​of the relative electronegativities, the oxidizing ability of non-metals increases in the following order: Si, B, H, P, C, S, I, N, Cl, O, F.

Non-metals as oxidants

The oxidizing properties of non-metals are manifested when they interact:

with metals: 2Na + Cl 2 = 2NaCl;

with hydrogen: H 2 + F 2 = 2HF;

with non-metals that have a 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.

Non-metals as reducing agents

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, that is, it can be a reducing agent. All other non-metals exhibit reducing properties. So, for example, chlorine does not combine directly with oxygen, but its oxides (Cl 2 O, ClO 2, Cl 2 O 2), in which chlorine exhibits a positive oxidation state, can be obtained indirectly. At high temperatures, nitrogen directly combines with oxygen and exhibits reducing properties. Sulfur reacts even more easily with oxygen.

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.

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

Cl 2 + H 2 O = HCl + HClO.

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

Nonmetal compounds

Non-metals can form compounds with different intramolecular bonds.

Types of compounds of non-metals

General formulas of hydrogen compounds by 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 non-molecular solids. 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). V 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 designated by the formula NH 4 OH and called ammonium hydroxide. It is also designated by the formula NH 3 ∙ H 2 O and is called ammonia hydrate.

With oxygen, non-metals 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), and in others, a lower one (for example, SO 2, N 2 O 3). Acidic oxides correspond to acids, and of the two oxygenic acids of one non-metal, the one in which it exhibits a higher oxidation state is stronger. For example, nitric acid HNO 3 is stronger than nitrous HNO 2, and sulfuric acid H 2 SO 4 is stronger than sulfuric H 2 SO 3.

Characteristics of oxygen compounds of non-metals

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

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

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

In the main subgroups of the periodic table of chemical elements, the acidic properties of higher oxides of non-metals decrease in the direction from top to bottom.

Halogens.

The structure of halogen atoms

Halogens include the elements of group VIII of the periodic table, 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 reducing properties increase. The value of the relative electronegativity of halogens decreases similarly. As the most electronegative element, fluorine, when combined with other elements, exhibits a constant oxidation state. -1 ... The remaining halogens can exhibit both this oxidation state in compounds with metals, hydrogen and less electronegative elements, and odd positive 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

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

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

Chemical properties of halogens

Interaction of halogens with metals with the formation of halides:

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

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

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

During the reactions of metals of side subgroups (transition metals) with halogens, halides with a high oxidation state of the metal are formed, for example:

2Fe + 3Cl 2 = 2FeCl 3,

but 2HCl + Fe = FeCl 2 + H 2.

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

Only fluorine interacts with hydrogen irreversibly, the rest of the halogens, depending on the conditions, can also 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 = 4HF + O 2

The 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, one can use the "activity series" of halogens, which is written as follows:

F 2> Сl 2> Br 2> I 2,

that is, the 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 = 2NаСl + Br 2

2KI + Br 2 = 2KBr + I 2.

Biological significance and application of halogens

Fluorine plays a very important role in the life of plants, animals and humans. Without fluoride, the development of the skeleton and especially teeth is impossible. The fluoride content in 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 an excess of it contributes to dental 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 buffering system of the blood. Sodium chloride is 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 seizures. A sufficient amount of chlorine is found in vegetables such as celery, radishes, cucumbers, white cabbage, dill, peppers, onions, artichokes. Bromine is also one of the essential trace elements and most of all it is found in the pituitary gland, blood. Thyroid gland, adrenal glands. Bromides in small doses (0.1-0.3 adults) have a positive effect on non-central nervous system as amplifiers of inhibition processes in the cerebral cortex. In nature, bromides accumulate in plants such as rye, wheat, barley, potatoes, carrots, cherries, and apples. Dutch cheese contains a lot of bromine. Iodine in the human body begins to accumulate in the womb. The human thyroid hormone - thyroxine - contains 60% of bound iodine. This hormone enters the liver, kidneys, mammary glands, and gastrointestinal tract through the blood stream. 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 contributes to the growth and nutritional status of animals, improves their health and fertility. The main suppliers of iodine for humans are cereals, eggplants, beans, white and cauliflower cabbage, potatoes, onions, carrots, cucumbers, pumpkin, lettuce, seaweed, squid.

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Are concretized general objectives of the main common education, taking into account the specifics of the academic subject; 2) generalcharacterization educational ... in electrolyte solutions. Variety of substances Generalcharacteristicnon-metals based on their position in the periodic ...

The periodic table of Dmitry Ivanovich Mendeleev is very convenient and universal in its use. According to it, it is possible to determine some characteristics of the elements, and what is most surprising, to predict some properties of the still undiscovered, not discovered by scientists, chemical elements (for example, we know some of the properties of the alleged unbigexy, although it has not yet been discovered and synthesized).

These properties depend on the ability of the element donate or attract electrons to oneself. It is important to remember one rule, metals - give up electrons, and non-metals - take. Accordingly, metallic properties are the ability of a certain chemical element to donate its electrons (from the external electron cloud) to another chemical element. For non-metals, the opposite is true. The more easily a non-metal accepts electrons, the higher its non-metallic properties.

Metals will never accept electrons from another chemical element. This is typical for the following elements;

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

With non-metals, things are similar. Fluorine shows its properties more than all other non-metals, 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) comes immediately after fluorine. Oxygen can form a compound with fluorine, donating its electrons, but it takes 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 tend to complete their energy level. For this, there must be 8 electrons at the last electronic level. The fluorine atom has 7 electrons on the last electron shell, trying to complete it, it attracts one more electron. The sodium atom on the outer shell has one electron, to get 8, it is easier for it to give 1, and at the last level there will be 8 negatively charged particles.

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

How metallic properties change in the periodic table

Periodic table of Mendeleev consists of groups and periods. The periods are arranged horizontally in such a way that the first period includes: lithium, beryllium, boron, carbon, nitrogen, oxygen, and so on. Chemical elements are arranged strictly in order of increasing 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 outer level of a particular chemical element. While the period number indicates the amount of electron 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 during 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 the group, the opposite is true, the metallic properties increase in the 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 to element give it away. 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 the fourth shell.

Let's compare two elements - calcium and barium. Barium in the periodic table is lower than calcium. This means that electrons from the outer shell of calcium are located closer to the nucleus, therefore, they are better attracted than that of barium.

Harder to compare items that are in different groups and periods. Take calcium and rubidium, for example. Rubidium will give off negative particles better than calcium. Since it stands below and to the left. But using only the periodic table, it is impossible to unequivocally answer this question when comparing magnesium and scandium (since one element is lower and to the right, and the other is higher and to the left). To compare these elements, special tables will be needed (for example, the electrochemical series of metal voltages).

How non-metallic properties change in the periodic table

Non-metallic properties in Mendeleev's periodic system change exactly the opposite than metallic ones. In fact, these two traits are antagonists.

Strengthened in the period (in a row from right to left). For example, sulfur is able to attract less 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 the 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 off negatively charged particles more than nitrogen, but at the same time it is able to attract better than arsenic. Phosphorus particles are attracted to the core better than arsenic particles, which gives it the advantage of an oxidizing agent in reactions to decrease and increase the oxidation state (redox reactions).

Compare, for example, sulfur and arsenic... Sulfur is located higher and to the right, which means that it is easier for it to complete its 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, you 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 of Mendeleev 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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Chemical properties of non-metals
According to the numerical values ​​of the relative electronegativities oxidative capacity of non-metals increases in the following order: Si, B, H, P, C, S, I, N, Cl, O, F.
Non-metals as oxidants
The oxidizing properties of non-metals are manifested when they interact:

· with metals: 2Na + Cl 2 = 2NaCl;

· with hydrogen: H 2 + F 2 = 2HF;

· with non-metals that have a 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.

Non-metals 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, that is, it can be a reducing agent. All other non-metals exhibit reducing properties. So, for example, chlorine does not combine directly with oxygen, but its oxides (Cl 2 O, ClO 2, Cl 2 O 2), in which chlorine exhibits a positive oxidation state, can be obtained indirectly. 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 one and the same non-metal 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, that is, the ability to donate electrons in chemical reactions.

Nonmetal compounds
Non-metals can form compounds with different intramolecular bonds.
Types of compounds of non-metals
General formulas of hydrogen compounds by groups of the periodic system of chemical elements are given in the table:

	RH 2	RH 3	RH 4	RH 3	H 2 R
Non-volatile hydrogen compounds				Volatile hydrogen compounds

With metals, hydrogen forms (with some exceptions) non-volatile compounds, which are non-molecular solids. 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 designated by the formula NH 4 OH and called ammonium hydroxide. It is also designated by the formula NH 3 ∙ H 2 O and is called ammonia hydrate.
With oxygen, non-metals 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), and in others, a lower one (for example, SO 2, N 2 O 3). Acidic oxides correspond to acids, and of the two oxygenic acids of one non-metal, the one in which it exhibits a higher oxidation state is stronger. For example, nitric acid HNO 3 is stronger than nitrous HNO 2, and sulfuric acid H 2 SO 4 is stronger than sulfuric H 2 SO 3.
Characteristics of oxygen compounds of non-metals

1. The properties of higher oxides (i.e., oxides containing an element of this 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 acids corresponding to these oxides.

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

4. In the main subgroups of the periodic table of chemical elements, the acidic properties of higher oxides of non-metals decrease in the direction from top to bottom.

Lecture 3. Non-metals

1. General characteristics of non-metallic elements

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 humans. Of carbon, hydrogen, oxygen, sulfur, phosphorus and nitrogen, all the most important organic matter, they are elements of life. Hydrogen and helium are the main elements of the Universe; all space objects, including our Sun, are made of them. It is impossible to imagine our life without compounds of non-metals, especially if we remember that a vital chemical compound - water - consists of hydrogen and oxygen.

Non-metals are chemical elements whose atoms accept electrons to complete an external energy level, while forming negatively charged ions.

Almost all non-metals have relatively small radii and big number electrons at the external energy level from 4 to 7, they are characterized by high values ​​of electronegativity and oxidizing properties.

1.1. Position of non-metallic elements in the periodic table of chemical elements of Mendeleev

If in the Periodic Table we draw a diagonal from boron to astatine, then on the right upwards along the diagonal there will be nonmetal elements, and from the bottom left - metals, these also include elements of all side subgroups, lanthanides and actinides. Elements located near the diagonal, for example, beryllium, aluminum, titanium, germanium, antimony, have a dual character and belong to metalloids. Non-metallic elements: s-element - hydrogen; p-elements of group 13 - boron; 14 groups - carbon and silicon; 15 groups - nitrogen, phosphorus and arsenic, 16 groups - oxygen, sulfur, selenium and tellurium and all elements of group 17 - fluorine, chlorine, bromine, iodine and astatine. Elements of group 18 - inert gases, occupy a special position, they have a completely completed outer electronic layer and occupy an intermediate position between metals and non-metals. They are sometimes referred to as non-metals, but formally, according to physical signs.

1.2. Electronic structure of non-metallic elements

Almost all non-metallic elements at the external energy level have a large number of electrons - from 4 to 7. Boron is an analogue of aluminum, it has only 3 electrons at the external energy level, but it has a small radius, firmly holds its electrons and has the properties of a non-metal. Let us especially note the electronic structure of hydrogen. It is an s-element, but it quite easily accepts one electron, forms a hydride ion, and exhibits the oxidizing properties of a metal.

The electronic configurations of valence electrons of non-metallic elements are given in the table:

1.3. Regularities in changing the properties of non-metallic elements

Let us consider some patterns in the change in the properties of non-metallic elements belonging to one period and one group based on the structure of their atoms.

In the period:

The charge of the nucleus increases,

The radius of the atom decreases

The number of electrons in the external energy level increases,

Electronegativity increases

Oxidizing properties are enhanced

Non-metallic properties are enhanced.

In a group:

The charge of the nucleus increases,

The radius of the atom is increasing

The number of electrons at the external energy level does not change,

Electronegativity decreases

Oxidizing properties weaken

Non-metallic properties are weakened.

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

Chemical elements - non-metals

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 humans. All the most important organic substances consist of carbon, hydrogen, oxygen, sulfur, phosphorus and nitrogen, they are the elements of life. Hydrogen and helium are the main elements of the Universe; all space objects, including our Sun, are made of them. It is impossible to imagine our life without compounds of non-metals, especially if we remember that a vital chemical compound - water - consists of hydrogen and oxygen.

If in the Periodic Table we draw a diagonal from beryllium to astatine, then on the right upwards along the diagonal there will be nonmetal elements, and on the bottom left - metals, these also include elements of all side subgroups, lanthanides and actinides. Elements located near the diagonal, for example, beryllium, aluminum, titanium, germanium, antimony, have a dual character and belong to metalloids. Non-metallic elements: s-element - hydrogen; p-elements of group 13 - boron; 14 groups - carbon and silicon; 15 groups - nitrogen, phosphorus and arsenic, 16 groups - oxygen, sulfur, selenium and tellurium and all elements of the 17th group - fluorine, chlorine, bromine, iodine and astatine... Elements of group 18 - inert gases, occupy a special position, they have a completely completed outer electron layer and occupy an intermediate position between metals and non-metals. They are sometimes referred to as non-metals, but formally, according to physical characteristics.

Nonmetals- these are chemical elements, the atoms of which accept electrons to complete the external energy level, while forming negatively charged ions.

The outer electron layer of nonmetal atoms contains from three to eight electrons.

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 ​​of electronegativity and oxidizing properties. Therefore, in comparison with metal atoms, non-metals are characterized by:

· Smaller atomic radius;

· Four or more electrons at the external energy level;

Hence, such an important property of nonmetal atoms is the tendency to accept missing up to 8 electrons, i.e. oxidizing properties. A qualitative characteristic of nonmetal atoms, i.e. Electronegativity can serve as a kind of measure of their non-metallicity, i.e. the property of atoms of chemical elements to polarize a chemical bond, to pull away common electron pairs;

The very first scientific classification of chemical elements was their division into metals and non-metals. This classification has not lost its significance at the present time. Non-metals are chemical elements whose atoms are characterized by the ability to accept electrons until the completion of the outer layer due to the presence, as a rule, of four or more electrons on the outer electron layer and the small radius of atoms in comparison with metal atoms.

This definition leaves aside the elements of Group VIII of the main subgroup - inert, or noble, gases, the atoms of which have a complete outer electron layer. The electronic configuration of the atoms of these elements is such that they cannot be attributed to either metals or non-metals. They are the objects that divide the elements into metals and non-metals, occupying a borderline position between them. Inert, or noble, gases ("nobility" is expressed in inertness) are sometimes referred to as non-metals, but only formally, according to physical characteristics. These substances remain gaseous down to very low temperatures. So, helium does not pass into a liquid state at t ° = -268.9 ° C.

Inertia in chemically these elements are relative. For xenon and krypton, compounds with fluorine and oxygen are known: KrF 2, XeF 2, XeF 4, etc. Undoubtedly, in the formation of these compounds, inert gases played the role of reducing agents. It follows from the definition of non-metals that their atoms are characterized by high values ​​of electronegativity. It varies from 2 to 4. Non-metals are elements of the main subgroups, mainly p-elements, with the exception of hydrogen - s-element.

All nonmetal elements (except hydrogen) occupy the upper right corner in the Periodic Table of Chemical Elements of D.I. However, special attention should be paid to the dual position of hydrogen in the Periodic Table: in the main subgroups of groups I and VII. This is no coincidence. On the one hand, a hydrogen atom, like alkali metal atoms, has one electron (1s 1 electronic configuration) on the outer (and only for it) electron layer, which it is able to donate, exhibiting the properties of a reducing agent.

In most of its compounds, hydrogen, like alkali metals, exhibits an oxidation state of +1. But the return of an electron by a hydrogen atom is more difficult than that of alkali metal atoms. On the other hand, the hydrogen atom, like the halogen atoms, lacks one electron to complete the outer electron layer, so the hydrogen atom can accept one electron, exhibiting the properties of an oxidizing agent and the oxidation state characteristic of halogen -1 in hydrides (compounds with metals, like metal compounds with halogens - halides). But the attachment of one electron to a hydrogen atom is more difficult than with halogens.

Under normal conditions, hydrogen H 2 is a gas. Its molecule, like halogens, is diatomic. Atoms of non-metals are dominated by oxidizing properties, that is, the ability to attach electrons. This ability is characterized by the value of electronegativity, which naturally changes in periods and subgroups. Fluorine is the most powerful oxidizing agent, its atoms in chemical reactions are not able to donate electrons, that is, to exhibit reducing properties. Other non-metals can exhibit reducing properties, albeit to a much weaker degree than metals; in periods and subgroups, their regenerative capacity changes in reverse order compared to oxidative.

• Non-metallic elements are located in the main subgroups III – VIII of groups PS D.I. Mendeleev, occupying its upper right corner.
• On the outer electron layer of atoms of non-metallic elements, there are from 3 to 8 electrons.
• The non-metallic properties of elements increase in periods and weaken in subgroups with an increase in the ordinal number of the element.
• Higher oxygen compounds of non-metals are acidic in nature (acidic oxides and hydroxides).
• The atoms of non-metallic elements are capable of both accepting electrons, exhibiting oxidative functions, and giving them away, exhibiting reductive functions.

Structure and physical properties of non-metals

In simple substances, the atoms of non-metals are bound covalent non-polar bond... Due to this, a more stable electronic system is formed than that of isolated atoms. In this case, single (for example, in hydrogen molecules H 2, halogens F 2, Br 2, I 2), double (for example, in sulfur molecules S 2), triple (for example, in nitrogen molecules N 2) covalent bonds are formed.

• No malleability
• No gloss
• Thermal conductivity (graphite only)
• The color is varied: yellow, yellowish-green, red-brown.
• Electrical Conductivity (Graphite and Black Phosphorus only.)

State of aggregation:

• liquid - Br 2;

Unlike metals, non-metals are simple substances, characterized by a wide variety of properties. Non-metals have a different state of aggregation under normal conditions:

• gases - H 2, O 2, O 3, N 2, F 2, Cl 2;
• liquid - Br 2;
• solids - modifications of sulfur, phosphorus, silicon, carbon, etc.

The spectrum of colors is much richer in non-metals: red - in phosphorus, red-brown - in bromine, yellow - in sulfur, yellow-green - in chlorine, violet - in iodine vapor. Elements - non-metals are more capable, in comparison with metals, of allotropy.

The ability of atoms of one chemical element to form several simple substances is called allotropy, and these simple substances are called allotropic modifications.

Simple substances - non-metals can have:

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

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

Many non-metallic elements form several simple substances - allotropic modifications... This property of atoms is called allotropy. Allotropy can be associated with a different composition of molecules (O 2, O 3), and with different structure crystals. Allotropic modifications of carbon are graphite, diamond, carbyne, fullerene. To identify the properties characteristic of all non-metals, one must pay attention to their location in the periodic table of elements and determine the configuration of the outer electron layer.

In the period:

• the charge of the nucleus increases;
• the radius of the atom decreases;
• the number of electrons in the outer layer increases;
• electronegativity increases;
• oxidizing properties are enhanced;
• non-metallic properties are enhanced.

In the main subgroup:

• the charge of the nucleus increases;
• the radius of the atom increases;
• the number of electrons on the outer layer does not change;
• electronegativity decreases;
• oxidizing properties are weakened;
• non-metallic properties are weakened.

For most metals, with rare exceptions (gold, copper and some others), a silvery-white color is characteristic. But for simple substances - non-metals, the gamut of colors is much more diverse: P, Se - yellow; B - brown; O 2 (g) - blue; Si, As (meth) - gray; P 4 - pale yellow; I - violet-black with a metallic sheen; Br 2 (g) - brown liquid; C1 2 (g) - yellow-green; F 2 (r) - pale green; S 8 (TV) - yellow. Crystals of non-metals are non-plastic, and any deformation causes the destruction of covalent bonds. Most non-metals do not have a metallic luster.

There are only 16 non-metal chemical elements! Not much when you consider that there are 114 known elements. Two non-metallic elements make up 76% of the mass of the earth's crust. These are oxygen (49%) and silicon (27%). The atmosphere contains 0.03% of the mass of oxygen in the earth's crust. Non-metals make up 98.5% of the mass of plants, 97.6% of the mass of the human body. Non-metals C, H, O, N, S are biogenic elements that form the most important organic substances of a living cell: proteins, fats, carbohydrates, nucleic acids. The air we breathe includes simple and complex substances, also formed by non-metallic elements (oxygen O 2, nitrogen N 2, carbon dioxide CO 2, water vapor H 2 O, etc.)

Oxidizing properties of simple substances - non-metals

For atoms of non-metals, and therefore, for the simple substances formed by them, they are characterized as oxidative and restorative properties.

1. Oxidizing properties of non-metals appear first of all when interacting with metals(metals are always reducing agents):

The oxidizing properties of chlorine Cl 2 are more pronounced than that of sulfur, therefore, the metal Fe, which has stable oxidation states of +2 and +3 in the compounds, is oxidized by it to a higher oxidation state.

1. Most non-metals exhibit oxidizing properties when interacting with hydrogen... As a result, volatile hydrogen compounds are formed.

2. Any non-metal acts as an oxidizing agent in reactions with those non-metals that have a lower electronegativity value:

The electronegativity of sulfur is greater than that of phosphorus, so it exhibits oxidizing properties here.

The electronegativity of fluorine is greater than that of all other chemical elements, therefore it exhibits the properties of an oxidizing agent. Fluorine F 2 is the strongest oxidizing agent among non-metals; it exhibits only oxidizing properties in reactions.

3. Non-metals also exhibit oxidizing properties in reactions with certain complex substances..

Let us first of all note the oxidizing properties of the non-metal oxygen in reactions with complex substances:

Not only oxygen, but also other non-metals can also be oxidizing agents in reactions with complex substances.- inorganic (1, 2) and organic (3, 4):

The strong oxidizing agent chlorine Cl 2 oxidizes iron (II) chloride to iron (III) chloride;

Chlorine Cl 2 as a stronger oxidizing agent displaces iodine I 2 in free form from the potassium iodide solution;

Methane halogenation is a characteristic reaction for alkanes;

A qualitative reaction to unsaturated compounds is their discoloration of bromine water.

Reducing properties of simple substances - non-metals

By revising reactions of non-metals with each other, that depending on the value of their electronegativity, one of them exhibits the properties of an oxidizing agent, and the other - the properties of a reducing agent.

1. In relation to fluorine, all non-metals (even oxygen) exhibit reducing properties.

2. Of course, non-metals, in addition to fluorine, serve as reducing agents when interacting with oxygen.

As a result of the reactions, nonmetal oxides: non-salt-forming and salt-forming acidic. And although halogens do not combine directly with oxygen, their oxides are known: Cl 2 +1 O -2, Cl 2 +4 O 2 -2, Cl 2 +7 O 7 -2, Br 2 +1 O -2, Br +4 O 2 -2, I 2 +5 O 5 -2, and others, which are obtained indirectly.

3. Many non-metals can act as a reducing agent in reactions with complex substances - oxidizing agents:

There are also reactions in which one and the same non-metal is both an oxidizing agent and a reducing agent. These are self-oxidation-self-healing (disproportionation) reactions:

Thus, most non-metals can act in chemical reactions both as an oxidizing agent and as a reducing agent (reducing properties are not inherent only to fluorine F 2).

Hydrogen compounds of non-metals

Unlike metals, non-metals form gaseous hydrogen compounds. Their composition depends on the oxidation state of the non-metals.

RH 4 → RH 3 → H 2 R → HR

A common property of all non-metals is the formation of volatile hydrogen compounds, in most of which the non-metal has the lowest oxidation state. Among the given formulas of substances there are many those whose properties, application and preparation you studied earlier: CH 4, NH 3, H 2 O, H 2 S, HCl.

It is known that these compounds can be obtained most simply directly the interaction of a non-metal with hydrogen, that is, by synthesis:

All hydrogen compounds of non-metals are formed by covalent polar bonds, have a molecular structure and, under normal conditions, are gases, except for water (liquid). Hydrogen compounds of non-metals are characterized by a different attitude to water. Methane and silane are practically insoluble in it. Ammonia, when dissolved in water, forms a weak base NH 3 H 2 O. When hydrogen sulfide, hydrogen selenide, hydrogen telluride, and also hydrogen halides are dissolved in water, acids are formed with the same formula as the hydrogen compounds themselves: H 2 S, H 2 Se, H 2 Te, HF, HCl, HBr, HI.

If we compare the acid-base properties of hydrogen compounds formed by non-metals of one period, for example, the second (NH 3, H 2 O, HF) or the third (PH 3, H 2 S, HCl), then we can conclude that their acidic properties are naturally enhanced and, accordingly, the weakening of the main ones. This is obviously due to the increase in polarity communication E-N(where E is a non-metal).

The acid-base properties of hydrogen compounds of non-metals of one subgroup also differ. For example, in the series of hydrogen halides HF, HCl, HBr, HI, the strength of the E-H bond decreases, since the bond length increases. In HCl, HBr, HI solutions, they dissociate almost completely - these are strong acids, and their strength increases from HF to HI. In this case, HF refers to weak acids, which is due to another factor - intermolecular interaction, the formation of hydrogen bonds ... H-F ... H-F .... Hydrogen atoms are bonded to fluorine F atoms not only of their own molecule, but also of the neighboring one.

Summarizing comparative characteristics acid-base properties of hydrogen compounds of non-metals, we conclude that the acidic and weakening of the basic properties of these substances by periods and main subgroups with an increase in the atomic numbers of their constituent elements.

Over the period in the PS of chemical elements, with an increase in the ordinal number of the element - non-metal, the acidic nature of the hydrogen compound increases.

SiH 4 → PH 3 → H 2 S → HCl

In addition to the properties considered, hydrogen compounds of non-metals in redox reactions always exhibit the properties of reducing agents, because in them the non-metal has lower degree oxidation.

Hydrogen

Hydrogen is the main element of the Universe. Many space objects (gas clouds, stars, including the Sun) are more than half composed of hydrogen. On Earth, it, including the atmosphere, hydrosphere and lithosphere, is only 0.88%. But this is by mass, and atomic mass hydrogen is very small. Therefore, its small content is only apparent, and out of every 100 atoms on Earth, 17 are hydrogen atoms.

In a free state, hydrogen exists in the form of H 2 molecules, the atoms are linked into a molecule covalent non-polar bond.

Hydrogen (H 2) is the lightest gas of all gaseous substances. Has the highest thermal conductivity and the most low temperature boiling (after helium). Slightly soluble in water. At a temperature of -252.8 ° C and atmospheric pressure, hydrogen turns into a liquid state.

1. The hydrogen molecule is very strong, which makes it inactive:

H 2 = 2H - 432 kJ

2. At ordinary temperatures, hydrogen reacts with active metals:

Ca + H 2 = CaH 2,

forming calcium hydride, and with F 2, forming hydrogen fluoride:

F 2 + H 2 = 2HF

3. When high temperatures get ammonia:

N 2 + 3H 2 = 2NH 3

and titanium hydride (metal in powder):

Ti + H 2 = TiH 2

4. When ignited, hydrogen reacts with oxygen:

2H 2 + O 2 = 2H 2 O + 484 kJ

5. Hydrogen has a regenerative ability:

CuO + H 2 = Cu + H 2 O

Elements of the main subgroup of the VII group of the periodic system, united under a common name halogens, fluorine (F), chlorine (Cl), bromine (Bg), iodine (I), astatine (At) (rare in nature) are typical non-metals. This is understandable, because their atoms contain on the external energy level seven electrons and they only need one electron to complete it. The atoms of these elements, when interacting with metals, receive an electron from the metal atoms. In this case, an ionic bond arises and salts are formed. Hence the common name "halogens", that is, "giving birth to salts."

very strong oxidants... Fluorine in chemical reactions exhibits only oxidizing properties, and it is characterized by an oxidation state of -1. The rest of the halogens can also exhibit reducing properties when interacting with more electronegative elements - fluorine, oxygen, nitrogen, while their oxidation states can take values ​​+1, +3, +5, +7. The reducing properties of halogens increase from chlorine to iodine, which is associated with an increase in the radii of their atoms: chlorine atoms are about half that of iodine.

Halogens are simple substances

All halogens exist in a free state as diatomic molecules with a covalent non-polar chemical bond between atoms. In the solid state F 2, Cl 2, Br 2, I 2 have molecular crystal lattices, which is confirmed by their physical properties.

With an increase in the molecular weight of halogens, the melting and boiling points increase, and the densities increase: bromine is a liquid, iodine is a solid, fluorine and chlorine are gases. This is due to the fact that with an increase in the size of atoms and molecules of halogens, the forces of intermolecular interaction between them increase. From F 2 to I 2, the intensity of the color of the halogens increases.

The chemical activity of halogens, as non-metals, weakens from fluorine to iodine, iodine crystals have a metallic luster. Each halogen is the strongest oxidizing agent in its period... The oxidizing properties of halogens are clearly manifested when they interact with metals. This produces salts. So, fluorine, even under normal conditions, reacts with most metals, and when heated, and with gold, silver, platinum, known for their chemical passivity. Aluminum and zinc in a fluorine atmosphere ignite:

The rest of the halogens react with metals when heated... Heated iron powder also ignites when exposed to chlorine. The experiment can be carried out as with antimony, but only iron filings must first be heated in an iron spoon, and then pour them in small portions into a flask with chlorine. Since chlorine is a strong oxidizing agent, iron (III) chloride is formed as a result of the reaction:

In bromine vapor red-hot copper wire burns out:

Iodine oxidizes metals more slowly, but in the presence of water, which is a catalyst, the reaction of iodine with aluminum powder proceeds very violently:

The reaction is accompanied by the release of violet iodine vapor.

On a decrease in the oxidizing and an increase in the reducing properties of halogens from fluorine to iodine can be judged by their ability to displace each other from solutions of their salts, and also it is clearly manifested when they interact with hydrogen. The equation for this reaction can be written in general view So:

If fluorine interacts with hydrogen in any conditions with an explosion, then a mixture of chlorine with hydrogen reacts only when ignited or irradiated with direct sunlight, bromine interacts with hydrogen when heated and without explosion. These reactions are exothermic. The reaction of the combination of iodine with hydrogen is weakly endothermic, it proceeds slowly even when heated.

As a result of these reactions, hydrogen fluoride HF, hydrogen chloride HCl, hydrogen bromide HBr and hydrogen iodide HI are formed, respectively.

Chlorine chemical properties in tables

Getting halogens

Fluorine and chlorine are obtained by electrolysis of melts or solutions of their salts. For example, the electrolysis process of sodium chloride melt can be reflected by the equation:

When chlorine is obtained by electrolysis of sodium chloride solution, in addition to chlorine, hydrogen and sodium hydroxide are also formed:

Oxygen (O)- the founder of the main subgroup of the VI group of the Periodic Table of the Elements. The elements of this subgroup - oxygen O, sulfur S, selenium Se, tellurium Te, polonium Po - have the general name "chalcogenes", which means "producing ores".

Oxygen is the most abundant element on our planet. It is part of water (88.9%), but it covers 2/3 of the surface of the globe, forming it aquatic shell- the hydrosphere. Oxygen is the second in quantity and the first in importance for life constituent part of the Earth's air shell - the atmosphere, where it accounts for 21% (by volume) and 23.15% (by mass). Oxygen is a part of numerous minerals of the hard shell of the earth's crust - the lithosphere: out of every 100 atoms of the earth's crust, 58 atoms account for oxygen.

Ordinary oxygen exists in the form of O 2. It is a colorless, odorless and tasteless gas. In the liquid state it has a light blue color, in the solid state it is blue. Gaseous oxygen is more soluble in water than nitrogen and hydrogen.

Oxygen interacts with almost all simple substances, except for halogens, noble gases, gold and platinum metals... Reactions of non-metals with oxygen occur very often with the release of a large amount of heat and are accompanied by ignition - combustion reactions. For example, the combustion of sulfur with the formation of SO 2, phosphorus - with the formation of P 2 O 5 or coal - with the formation of CO 2. Almost all reactions involving oxygen are exothermic. An exception is the interaction of nitrogen with oxygen: this is an endothermic reaction that occurs at temperatures above 1200 ° C or with an electric discharge:

Oxygen vigorously oxidizes not only simple, but also many complex substances, while the oxides of the elements from which they are built are formed:

The high oxidizing power of oxygen is the basis for the combustion of all types of fuel.

Oxygen also participates in the processes of slow oxidation of various substances at ordinary temperatures. The role of oxygen in the process of respiration in humans and animals is extremely important. Plants also absorb atmospheric oxygen. But if in the dark there is only the process of absorption of oxygen by plants, then another process opposite to it occurs in the light - photosynthesis, as a result of which plants absorb carbon dioxide and release oxygen.

In industry, oxygen is obtained from liquid air, and in the laboratory - by decomposition of hydrogen peroxide in the presence of a manganese dioxide catalyst MnO 2 :

and decomposition of potassium permanganate KMnO 4 when heated:

Oxygen chemical properties in tables

Oxygen application

Oxygen is used in the metallurgical and chemical industries to speed up (intensify) production processes. Pure oxygen is also used to obtain high temperatures, for example, in gas welding and metal cutting. In medicine, oxygen is used in cases of temporary breathing difficulties associated with certain diseases. Oxygen is also used in metallurgy as an oxidizer for rocket fuel, in aviation for breathing, for cutting metals, for welding metals, and for blasting operations. Oxygen is stored in blue painted steel cylinders under a pressure of 150 atm. In laboratory conditions, oxygen is stored in glass devices - gas meters.

Atoms sulfur (S), like oxygen atoms and all other elements of the main subgroup of group VI, contain at the external energy level 6 electrons, of which two electrons unpaired... However, in comparison with oxygen atoms, sulfur atoms have a larger radius and a lower electronegativity value, therefore they exhibit pronounced reducing properties, forming compounds with oxidation states +2, +4, +6. In relation to less negative elements (hydrogen, metals), sulfur exhibits oxidizing properties and acquires an oxidation state -2 .

Sulfur is a simple substance

Sulfur, like oxygen, is characterized by allotropy. There are many modifications of sulfur with a cyclic or linear structure of molecules of various compositions.

The most stable modification is known as rhombic sulfur, consisting of S 8 molecules. Its crystals are octahedra with cut corners. They are colored lemon yellow and translucent, melting point 112.8 ° C. All other modifications are converted into this modification at room temperature. Crystallization from the melt first produces monoclinic sulfur (needle crystals, melting point 119.3 ° C), which then becomes rhombic. When heated pieces of sulfur in a test tube, it melts, turning into a liquid yellow color... At a temperature of about 160 ° C, liquid sulfur begins to darken, becomes thick and viscous, does not pour out of the test tube, upon further heating it turns into a mobile liquid, but retains its former dark brown color. If you pour it into cold water, it solidifies in the form of a transparent rubbery mass. This is plastic sulfur. It can also be obtained in the form of threads. After a few days, it also turns into rhombic sulfur.

Sulfur is insoluble in water. Sulfur crystals sink in water, but the powder floats on the surface of the water, since small sulfur crystals are not wetted by water and are kept afloat by small air bubbles. This is a flotation process. Sulfur is slightly soluble in ethyl alcohol and diethyl ether, well soluble in carbon disulfide.

Under normal conditions sulfur reacts with all alkali and alkaline earth metals, copper, mercury, silver, for example:

This reaction underlies the removal and disposal of spilled mercury from, for example, a broken thermometer. Visible drops of mercury can be collected on a piece of paper or copper plastic. The mercury that has got into the cracks must be covered with sulfur powder. This process is called demercurization.

When heated, sulfur also reacts with other metals (Zn, Al, Fe), and only gold does not interact with it under any conditions. Sulfur also exhibits oxidizing properties with hydrogen, with which it reacts when heated:

Of non-metals, only nitrogen, iodine and noble gases do not react with sulfur. Sulfur burns with a bluish flame, forming sulfur oxide (IV):

This compound is commonly known as sulphurous gas.

Chemical properties of sulfur in tables

Sulfur is one of the most common elements: the earth's crust contains 4.7 · 10-2% sulfur by mass (15th place among other elements), and the Earth as a whole is much more (0.7%). The main mass of sulfur is located in the depths of the earth, in its mantle-layer located between the earth's crust and the core of the Earth. Here, at a depth of about 1200-3000 km, a thick layer of sulphides and metal oxides lies. In the earth's crust, sulfur is found both in a free state (native) and, mainly, in the form of compounds of sulfides and sulfates. The most common sulfides in the earth's crust are pyrite FeS2, chalcopyrite FeCuS2, lead luster (galena) PbS, zinc blende (sphalerite) ZnS. Large quantities Sulfur is found in the earth's crust in the form of sparingly soluble sulfates - gypsum CaSO4 2H2O, barite BaSO4; sulfates of magnesium, sodium and potassium are widespread in seawater.

It is interesting that in ancient times of the geological history of the Earth (about 800 million years ago) there were no sulfates in nature. They were formed as products of sulfide oxidation, when an oxygen atmosphere arose as a result of the vital activity of plants. Hydrogen sulfide H2S and sulfur dioxide SO2 are found in volcanic gases. therefore, native sulfur found in areas close to active volcanoes (Sicily, Japan) could be formed by the interaction of these two gases:

2H 2 S + SO 2 = 3S + 2H 2 O.

Other deposits of native sulfur are associated with the vital activity of microorganisms.

Microorganisms are involved in many chemical processes, which in general make up the sulfur cycle in nature. With their assistance, sulfides are oxidized to sulfates, sulfates are absorbed by living organisms, where sulfur is reduced and is part of proteins and other vital substances. When the dead remains of organisms decay, proteins are destroyed, and hydrogen sulfide is released, which is then oxidized either to elemental sulfur (this is how sulfur deposits are formed), or to sulfates. Interestingly, bacteria and algae that oxidize hydrogen sulfide to sulfur collect it in their cells. The cells of such microorganisms can be 95% pure sulfur.

The origin of sulfur can be established by the presence of its analogue - selenium in it: if selenium is found in native sulfur, then sulfur is of volcanic origin, if not - biogenic, since microorganisms avoid including selenium in their life cycle Also, biogenic sulfur contains more isotope 32S than the heavier 34S.

The biological significance of sulfur

Vital chemical element... It is part of proteins - one of the main chemical components of the cells of all living organisms. Especially a lot of sulfur is in the proteins of hair, horns, wool. In addition, sulfur is an integral part of biologically active substances organism: vitamins and hormones (for example, insulin). Sulfur is involved in the redox processes of the body. With a lack of sulfur in the body, fragility and fragility of bones and hair loss are observed.

Leguminous plants (peas, lentils), oatmeal, eggs are rich in sulfur.

Sulfur application

Sulfur is used in the manufacture of matches and paper, rubber and paints, explosives and drugs, plastics and cosmetics. In agriculture, it is used to control plant pests. However, the main consumer of sulfur is the chemical industry. About half of the sulfur produced in the world goes to the production of sulfuric acid.

Nitrogen

Nitrogen (N)- the first representative of the main subgroup of group V of the Periodic system. Its atoms contain five electrons on the external energy level, of which three are unpaired. It follows that the atoms of these elements can attach three electrons, completing the external energy level.

Nitrogen atoms can donate their outer electrons to more electronegative elements (fluorine, oxygen) and acquire oxidation states +3 and +5. Nitrogen atoms exhibit reducing properties in oxidation states +1, +2, +4.

In a free state, nitrogen exists in the water of the diatomic molecule N 2. In this molecule, two N atoms are linked by a very strong triple covalent bond, these bonds can be designated as follows:

Nitrogen is a colorless, odorless and tasteless gas.

Under normal conditions nitrogen interacts only with lithium, forming Li nitride 3 N:

It interacts with other metals only at high temperatures.

Also at high temperatures and pressures in the presence of a catalyst, nitrogen reacts with hydrogen to form ammonia:

At the temperature of the electric arc, it combines with oxygen to form nitric oxide (II):

Chemical properties of nitrogen in tables

Nitrogen application

Nitrogen obtained by distillation of liquid air is used in industry for the synthesis of ammonia and the production of nitric acid... In medicine, pure nitrogen is used as an inert medium for the treatment of pulmonary tuberculosis, and liquid nitrogen is used in the treatment of diseases of the spine, joints, etc.

Phosphorus

The chemical element phosphorus forms several allotropic modifications. Two of them are simple substances: white phosphorus and red phosphorus. White phosphorus has a molecular crystal lattice consisting of P 4 molecules. Insoluble in water, readily soluble in carbon disulfide. It easily oxidizes in air, and even ignites in a powdery state. White phosphorus is highly toxic. A special property is the ability to glow in the dark due to oxidation. Store it under water. Red phosphorus is a dark crimson powder. It does not dissolve in water or carbon disulfide. It oxidizes slowly in air and does not ignite spontaneously. It is non-toxic and does not glow in the dark. When red phosphorus is heated in a test tube, it turns into white phosphorus (concentrated vapors).

The chemical properties of red and white phosphorus are similar, but white phosphorus is more chemically active. So, both of them interact with metals, forming phosphides:

White phosphorus ignites spontaneously in air, and red burns when ignited. In both cases, phosphorus (V) oxide is formed, which is emitted in the form of thick white smoke:

Phosphorus does not directly react with hydrogen, phosphine PH 3 can be obtained indirectly, for example, from phosphides:

Phosphine is a very poisonous gas with an unpleasant odor. Flammable in air. It is this property of phosphine that explains the appearance of wandering bog fires.

Chemical properties of phosphorus in tables

Phosphorus use

Phosphorus is the most important biogenic element and at the same time is widely used in industry. Red phosphorus is used in the manufacture of matches. Together with finely crushed glass and glue, it is applied to lateral surface boxes. Friction of a match head, which contains potassium chlorate and sulfur, ignites.

Perhaps the first property of phosphorus that a person has put to his service is flammability. The flammability of phosphorus is very high and depends on the allotropic modification.

The most active chemical, toxic and flammable white ("yellow") phosphorus, therefore it is very often used (in incendiary bombs, etc.).

Red phosphorus is the main modification produced and consumed by industry. It is used in the production of matches, explosives, incendiary compounds, various types of fuel, as well as extreme pressure lubricants, as a getter in the production of incandescent lamps.

Phosphorus (in the form of phosphates) is one of the three most important nutrients involved in the synthesis of ATP. Most of the phosphoric acid goes to receive phosphate fertilizers- superphosphate, precipitate, ammophos, etc.

Phosphates are widely used:

• as complexing agents (water softeners),
• as part of metal surface passivators (protection against corrosion, for example, the so-called "mazhef" composition).

The ability of phosphates to form a durable three-dimensional polymer mesh used for the manufacture of phosphate and aluminophosphate binders.

Carbon

Carbon (C)- the first element of the main subgroup of the VI group of the Periodic system. Its atoms contain 4 electrons on the outer level, so they can accept four electrons, while acquiring the oxidation state -4 , i.e., exhibit oxidizing properties and donate their electrons to more electronegative elements, i.e., exhibit reducing properties, while acquiring an oxidation state +4.

Carbon is a simple substance

Carbon forms allotropic modifications diamond and graphite... Diamond is a transparent crystalline substance, the hardest of all natural substances... It serves as a standard of hardness, which according to the ten-point system is estimated with the highest score of 10. Such hardness of a diamond is due to the special structure of its atomic crystal lattice. In it, each carbon atom is surrounded by the same atoms located at the vertices of a regular tetrahedron.

Diamond crystals are usually colorless, but they come in blue, light blue, red and black. They have a very strong gloss due to their high light refractive and reflective properties. And due to their extremely high hardness, they are used for the manufacture of drills, drills, grinding tools, glass cutting.

The largest diamond deposits are located in South Africa, and in Russia they are mined in Yakutia.

Graphite is a dark gray, greasy crystalline substance with a metallic luster. Unlike diamond, graphite is soft (leaves a mark on paper) and opaque, conducts heat and electric current well. The softness of graphite is due to its layered structure. In the crystal lattice of graphite, carbon atoms lying in one plane are firmly bound into regular hexagons. The bonds between the layers are weak. It is very refractory. Graphite is used to make electrodes, solid lubricants, neutron moderators in nuclear reactors, and pencil rods. At high temperatures and pressures, artificial diamonds are obtained from graphite, which are widely used in technology.

Soot and charcoal have a structure similar to graphite. Charcoal is obtained by dry distillation of wood. This coal, due to its porous surface, has a remarkable ability to absorb gases and solutes. This property is called adsorption. The greater the porosity of the charcoal, the more efficient the adsorption. To increase the absorption capacity, charcoal is treated with hot steam. The carbon treated in this way is called activated or active. In pharmacies, it is sold in the form of black carbolene tablets.

Chemical properties of carbon

Diamond and graphite combine with oxygen at very high temperatures. Soot and coal interact with oxygen much more easily, burning in it. But in any case, the result of such an interaction is the same - carbon dioxide is formed:

With metals, carbon forms when heated carbides:

Aluminum carbide- light yellow transparent crystals. Known calcium carbide CaC 2 in the form of pieces gray... It is used by gas welders to produce acetylene:

Acetylene used for cutting and welding metals, burning it with oxygen in special torches.

If you act on aluminum carbide with water, you get another gas - methane CH 4:

Silicon

Silicon (Si) is the second element of the main subgroup of group IV of the periodic system. In nature, silicon is the second most common chemical element after oxygen. More than a quarter of the earth's crust consists of its compounds. The most common silicon compound is its dioxide SiO 2 - silica. In nature, it forms the mineral quartz and many varieties, such as rock crystal and its famous purple form - amethyst, as well as agate, opal, jasper, chalcedony, carnelian. Silicon dioxide is also common and quartz sand. The second type of natural silicon compounds are silicates. Among them, the most common are aluminosilicates - granite, different kinds clay, mica. The non-aluminum silicate is, for example, asbestos. Silicon oxide is essential for plant and animal life. It gives strength to plant stems and animal protective covers. Silicon gives smoothness and strength to human bones. Silicon is a part of the lowest living organisms - diatoms and radiolarians.

Silicon chemical properties

Silicon burns in oxygen forming silicon dioxide or silicon oxide (IV):

As a non-metal, when heated, it combines with metals to form silicides:

Silicides are readily decomposed by water or acids, while a gaseous hydrogen silicon compound is released - silane:

4HCl + Mg 2 Si → SiH 4 + 2MgCl 2

Unlike hydrocarbons, silane ignites spontaneously in air and burns to form silicon dioxide and water:

The increased reactivity of silane in comparison with methane CH 4 is explained by the fact that silicon has a larger atom size than carbon; therefore, the Si-H chemical bonds are weaker than CH bonds.

Silicon interacts with concentrated aqueous solutions of alkali, forming silicates and hydrogen:

Silicon is obtained, reducing it from dioxide with magnesium or carbon:

Silicon oxide (IV), or silicon dioxide, or silica SiO 2, like CO 2, is an acidic oxide. However, unlike CO 2, it has not a molecular, but an atomic crystal lattice. Therefore, SiO 2 is a solid and refractory substance. It does not dissolve in water and acids, except for hydrofluoric acid, but at high temperatures interacts with alkalis to form silicic acid salts - silicates:

Silicates can also be obtained by fusing silicon dioxide with metal oxides or carbonates:

Sodium and potassium silicates are called soluble glass. Their aqueous solutions are well known silicate glue. From solutions of silicates, the action of stronger acids on them - hydrochloric, sulfuric, acetic and even carbonic - silicic acid H 2 SiO 3 :

Hence, H 2 SiO 3 - very weak acid... It is insoluble in water and falls out of the reaction mixture in the form of a gelatinous precipitate, sometimes filling the entire volume of the solution compactly, turning it into a semi-solid mass, similar to jelly, jelly. When this mass dries, a highly porous substance is formed - silica gel, which is widely used as an adsorbent - absorbent of other substances.

Reference material for passing the test:

Mendeleev table

Solubility table