Each period of DI Mendeleev's Periodic Table ends with an inert, or noble, gas.

The most common of the inert (noble) gases in the Earth's atmosphere is argon, which was isolated in its pure form earlier than other analogues. What is the reason for the inertness of helium, neon, argon, krypton, xenon and radon?

The fact that atoms of inert gases have eight electrons at the outermost levels farthest from the nucleus (helium has two). Eight electrons at the outer level is the limiting number for each element of DI Mendeleev's Periodic Table, except for hydrogen and helium. This is a kind of ideal of the strength of the energy level, to which the atoms of all other elements of the Periodic Table of D.I.Mendeleev strive.

Atoms can achieve such a position of electrons in two ways: by donating electrons from the external level (in this case, the external incomplete level disappears, and the penultimate level, which was completed in the previous period, becomes external) or by accepting electrons, which are not enough until the coveted eight. Atoms that have a smaller number of electrons on the outer level donate them to atoms that have more electrons on the outer level. It is easy to donate one electron, when it is the only one on the external level, to the atoms of the elements of the main subgroup of group I (group IA). It is more difficult to donate two electrons, for example, to the atoms of the elements of the main subgroup of group II (IIA group). It is even more difficult to donate your three outer electrons to the atoms of Group III (Group IIIA) elements.

Atoms of metal elements have a tendency to give up electrons from the outer level.... And the easier the atoms of a metal element give up their outer electrons, the more pronounced its metallic properties are. It is therefore clear that the most typical metals in the Periodic Table of D. I. Mendeleev are elements of the main subgroup of group I (group IA). And vice versa, atoms of non-metallic elements have a tendency to accept the missing before the completion of the external energy level. From what has been said, the following conclusion can be drawn. Within the period, with an increase in the charge of the atomic nucleus, and, accordingly, with an increase in the number of external electrons, the metallic properties of chemical elements weaken. The non-metallic properties of elements, characterized by the ease of accepting electrons to the external level, are enhanced at the same time.

The most typical non-metals are the elements of the main subgroup of group VII (group VIIA) of the Periodic Table of D. I. Mendeleev. At the outer level of the atoms of these elements, there are seven electrons. Up to eight electrons at the external level, that is, to a stable state of atoms, they lack one electron each. They easily attach them, showing non-metallic properties.

And how do the atoms of the elements of the main subgroup of the IV group (IVA group) of the periodic table of D.I.Mendeleev behave? After all, they have four electrons at the outer level, and it would seem that they do not care whether to give or receive four electrons. It turned out that the ability of atoms to give or receive electrons is influenced not only by the number of electrons at the outer level, but also by the radius of the atom. Within the period, the number of energy levels of atoms of elements does not change, it is the same, but the radius decreases, as the positive charge of the nucleus (the number of protons in it) increases. As a result, the attraction of electrons to the nucleus increases, and the radius of the atom decreases, the atom seems to be compressed. Therefore, it becomes more and more difficult to donate external electrons and, conversely, it becomes easier to accept the missing electrons up to eight.

Within the same subgroup, the radius of the atom increases with an increase in the charge of the atomic nucleus, since with a constant number of electrons at the outer level (it is equal to the number of the group), the number of energy levels increases (it is equal to the number of the period). Therefore, it becomes increasingly easier for the atom to donate external electrons.

In the Periodic Table of D. I. Mendeleev, with an increase in the serial number, the properties of atoms of chemical elements change as follows.

What is the result of the acceptance or release of electrons by the atoms of chemical elements?

Let's imagine that two atoms “meet”: a metal atom of group IA and a non-metal atom of group VIIA. The metal atom has a single electron on the external energy level, and the non-metal atom just lacks just one electron for its external level to be complete.

The metal atom will easily give up its farthest from the nucleus and weakly bound to it electron to the nonmetal atom, which will give it a free space on its external energy level.

Then a metal atom, devoid of one negative charge, will acquire a positive charge, and a non-metal atom, thanks to the resulting electron, will turn into a negatively charged particle - an ion.

Both atoms will fulfill their "cherished dream" - they will receive the much-desired eight electrons on the external energy level. But what happens next? Oppositely charged ions in full accordance with the law of attraction of opposite charges will immediately combine, that is, a chemical bond will arise between them.

The chemical bond formed between ions is called ionic.

Let us consider the formation of this chemical bond using the example of the well-known compound of sodium chloride (table salt):

The process of transformation of atoms into ions is shown in the diagram and figure:

For example, an ionic bond is also formed when calcium and oxygen atoms interact:

This transformation of atoms into ions always occurs when the atoms of typical metals and typical non-metals interact.

In conclusion, let us consider an algorithm (sequence) of reasoning when writing a scheme for the formation of an ionic bond, for example, between calcium and chlorine atoms.

1. Calcium is an element of the main subgroup of group II (HA group) of the Periodic Table of DI Mendeleev, metal. It is easier for its atom to donate two external electrons than to accept the missing six:

2. Chlorine is an element of the main subgroup of group VII (group VIIA) of the DI Mendeleev's table, non-metal. It is easier for its atom to accept one electron, which it lacks until the completion of the external energy level, than to donate seven electrons from the external level:

3. First, we find the smallest common multiple between the charges of the formed ions, it is equal to 2 (2 × 1). Then we determine how many calcium atoms need to be taken in order for them to give up two electrons (that is, we need to take 1 Ca atom), and how many chlorine atoms need to be taken so that they can take two electrons (that is, we need to take 2 Cl atoms) ...

4. Schematically, the formation of an ionic bond between calcium and chlorine atoms can be written as follows:

To express the composition of ionic compounds, formula units are used - analogs of molecular formulas.

The numbers showing the number of atoms, molecules or formula units are called coefficients, and the numbers showing the number of atoms in a molecule or ions in a formula unit are called indices.

In the first part of the paragraph, we made a conclusion about the nature and reasons for changing the properties of elements. In the second part of this section, we will list the keywords.

Key words and phrases

1. Atoms of metals and non-metals.
2. Ions are positive and negative.
3. Ionic chemical bond.
4. Odds and indices.

Work with computer

1. Please refer to the electronic attachment. Study the material in the lesson and complete the suggested assignments.
2. Search the Internet for e-mail addresses that can serve as additional sources for revealing the content of the keywords and phrases in the paragraph. Offer to help the teacher prepare a new lesson by reporting on the keywords and phrases in the next paragraph.

Questions and tasks

1. Compare the structure and properties of atoms: a) carbon and silicon; b) silicon and phosphorus.
2. Consider the schemes for the formation of an ionic bond between the atoms of chemical elements: a) potassium and oxygen; b) lithium and chlorine; c) magnesium and fluorine.
3. Name the most typical metal and the most typical non-metal of DI Mendeleev's Periodic Table.
4. Using additional sources of information, explain why inert gases have come to be called noble.

What happens to the atoms of elements during chemical reactions? What do the properties of elements depend on? One answer can be given to both of these questions: the reason lies in the structure of the external. In our article we will consider the electronic of metals and non-metals and find out the relationship between the structure of the external level and the properties of the elements.

Special properties of electrons

During the passage of a chemical reaction between the molecules of two or more reagents, changes occur in the structure of the electronic shells of atoms, while their nuclei remain unchanged. First, let's get acquainted with the characteristics of electrons located at the most distant levels of the atom from the nucleus. Negatively charged particles are arranged in layers at a certain distance from the nucleus and from each other. The space around the nucleus, where it is most possible to find electrons, is called the electron orbital. About 90% of the negatively charged electron cloud is condensed in it. The electron itself in an atom exhibits the property of duality; it can simultaneously behave both as a particle and as a wave.

Rules for filling the electron shell of an atom

The number of energy levels at which the particles are located is equal to the number of the period where the element is located. What does the electronic composition indicate? It turned out that at the external energy level for s- and p-elements of the main subgroups of small and large periods corresponds to the group number. For example, lithium atoms of the first group, which have two layers, have one electron on the outer shell. Sulfur atoms contain six electrons at the last energy level, since the element is located in the main subgroup of the sixth group, etc. If we are talking about d-elements, then the following rule exists for them: the number of external negative particles is 1 (for chromium and copper) or 2. This is explained by the fact that as the charge of the atomic nucleus increases, the inner d-sublevel is first filled and the outer energy levels remain unchanged.

Why do the properties of elements of small periods change?

Periods 1, 2, 3 and 7 are considered small. A smooth change in the properties of elements as nuclear charges increase, ranging from active metals to inert gases, is explained by a gradual increase in the number of electrons at the external level. The first elements in such periods are those whose atoms have only one or two electrons, which can easily be detached from the nucleus. In this case, a positively charged metal ion is formed.

Amphoteric elements, for example, aluminum or zinc, fill their external energy levels with a small number of electrons (1 for zinc, 3 for aluminum). Depending on the conditions of the chemical reaction, they can exhibit both the properties of metals and non-metals. Non-metallic elements of small periods contain from 4 to 7 negative particles on the outer shells of their atoms and complete it up to an octet, attracting the electrons of other atoms. For example, a non-metal with the highest electronegativity index - fluorine, has 7 electrons on the last layer and always takes one electron not only from metals, but also from active non-metallic elements: oxygen, chlorine, nitrogen. Small periods, as well as large ones, end with inert gases, whose monatomic molecules have completely completed external energy levels up to 8 electrons.

Features of the structure of atoms of long periods

Even rows of 4, 5, and 6 periods consist of elements, the outer shells of which contain only one or two electrons. As we said earlier, they fill the d- or f-sublevels of the penultimate layer with electrons. These are usually typical metals. Their physical and chemical properties change very slowly. Odd rows contain elements in which the external energy levels are filled with electrons according to the following scheme: metals - amphoteric element - non-metals - inert gas. We have already observed its manifestation in all small periods. For example, in the odd row of the 4th period, copper is a metal, zinc is amphoteric, then from gallium to bromine there is an increase in non-metallic properties. The period ends with krypton, whose atoms have a fully completed electron shell.

How to explain the division of elements into groups?

Each group - and there are eight of them in the short form of the table, is also divided into subgroups, called main and secondary. This classification reflects the different position of the electrons on the external energy level of the atoms of the elements. It turned out that in the elements of the main subgroups, for example, lithium, sodium, potassium, rubidium and cesium, the last electron is located at the s-sublevel. Elements of the 7th group of the main subgroup (halogens) fill their p-sublevel with negative particles.

For representatives of side subgroups, such as chromium, filling with electrons of the d-sublevel will be typical. And the elements of the family accumulate negative charges at the f-sublevel of the penultimate energy level. Moreover, the group number, as a rule, coincides with the number of electrons capable of forming chemical bonds.

In our article, we found out what structure the external energy levels of atoms of chemical elements have, and determined their role in interatomic interactions.

E. N. FRENKEL

Chemistry tutorial

A guide for those who do not know, but wants to know and understand chemistry

Part I. Elements of general chemistry
(the first level of difficulty)

Continuation. For the beginning, see No. 13, 18, 23/2007

Chapter 3. Elementary information about the structure of the atom.
D. I. Mendeleev's periodic law

Consider what an atom is, what an atom is made of, whether an atom changes in chemical reactions.

An atom is an electrically neutral particle consisting of a positively charged nucleus and negatively charged electrons.

The number of electrons may change during chemical processes, but the charge of the nucleus always remains unchanged... Knowing the distribution of electrons in an atom (the structure of an atom), it is possible to predict many properties of a given atom, as well as the properties of simple and complex substances, of which it is part.

The structure of the atom, i.e. the composition of the nucleus and the distribution of electrons around the nucleus can be easily determined by the position of the element in the periodic table.

In the periodic system of D.I. Mendeleev, chemical elements are arranged in a certain sequence. This sequence is closely related to the structure of the atoms of these elements. Each chemical element in the system is assigned serial number, in addition, for it you can specify the period number, group number, type of subgroup.

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Knowing the exact "address" of a chemical element - group, subgroup and period number, one can uniquely determine the structure of its atom.

Period Is a horizontal row of chemical elements. There are seven periods in the modern periodic system. The first three periods - small since they contain 2 or 8 elements:

1st period - H, Not - 2 elements;

2nd period - Li ... Ne - 8 elements;

3rd period - Na ... Ar - 8 elements.

Other periods - big... Each of them contains 2-3 rows of elements:

4th period (2 rows) - K ... Kr - 18 elements;

6th period (3 rows) - Сs ... Rn - 32 elements. This period includes a number of lanthanides.

Group- a vertical row of chemical elements. There are eight groups in total. Each group consists of two subgroups: main subgroup and side subgroup... For instance:

The main subgroup is formed by chemical elements of small periods (for example, N, P) and large periods (for example, As, Sb, Bi).

A side subgroup is formed by chemical elements of only long periods (for example, V, Nb,
Ta).

Visually, these subgroups are easy to distinguish. The main subgroup is "high", it starts from the 1st or 2nd period. A side subgroup - "low", starts from the 4th period.

So, each chemical element of the periodic system has its own address: period, group, subgroup, serial number.

For example, vanadium V is a chemical element of the 4th period, group V, side subgroup, serial number 23.

Task 3.1. Indicate the period, group and subgroup for chemical elements with serial numbers 8, 26, 31, 35, 54.

Task 3.2. Indicate the serial number and name of the chemical element, if it is known that it is located:

a) in the 4th period, group VI, side subgroup;

b) in the 5th period, IV group, main subgroup.

How can you connect information about the position of an element in the periodic table with the structure of its atom?

An atom consists of a nucleus (it has a positive charge) and electrons (they have a negative charge). In general, the atom is electrically neutral.

Positive nuclear charge is equal to the ordinal number of a chemical element.

The nucleus of an atom is a complex particle. Almost all the mass of an atom is concentrated in the nucleus. Since a chemical element is a collection of atoms with the same nuclear charge, the following coordinates are indicated near the symbol of the element:

From this data, you can determine the composition of the nucleus. The nucleus is made up of protons and neutrons.

Proton p has a mass of 1 (1.0073 amu) and a charge of +1. Neutron n has no charge (neutral), and its mass is approximately equal to the mass of a proton (1.0087 amu).

The charge of the nucleus is determined by the protons. And the number of protons is(largest) nuclear charge, i.e. ordinal number.

Number of neutrons N determined by the difference between the quantities: "core mass" A and "serial number" Z... So, for an aluminum atom:

N = A – Z = 27 –13 = 14n,

Task 3.3. Determine the composition of the nuclei of atoms if a chemical element is in:

a) 3rd period, VII group, main subgroup;

b) 4th period, IV group, side subgroup;

c) 5th period, I group, main subgroup.

Attention! When determining the mass number of an atomic nucleus, it is necessary to round off the atomic mass indicated in the periodic system. This is done because the masses of the proton and neutron are practically integer, and the mass of electrons can be neglected.

Let's determine which of the nuclei given below belong to the same chemical element:

A (20 R + 20n),

B (19 R + 20n),

IN 20 R + 19n).

The nuclei A and B belong to the atoms of the same chemical element, since they contain the same number of protons, that is, the charges of these nuclei are the same. Studies show that the mass of an atom does not significantly affect its chemical properties.

Isotopes are atoms of the same chemical element (the same number of protons) that differ in mass (different numbers of neutrons).

Isotopes and their chemical compounds differ from each other in physical properties, but the chemical properties of isotopes of one chemical element are the same. So, isotopes of carbon-14 (14 C) have the same chemical properties as carbon-12 (12 C), which are included in the tissues of any living organism. The difference is manifested only in radioactivity (isotope 14 C). Therefore, isotopes are used for the diagnosis and treatment of various diseases, for scientific research.

Let's return to the description of the structure of the atom. As you know, the atomic nucleus does not change in chemical processes. What is changing? The total number of electrons in the atom and the distribution of electrons turn out to be variable. General number of electrons in a neutral atom It is not difficult to determine - it is equal to the ordinal number, i.e. the charge of the atomic nucleus:

Electrons have a negative charge of –1, and their mass is negligible: 1/1840 of the mass of a proton.

Negatively charged electrons repel each other and are at different distances from the nucleus. Wherein electrons, which have approximately equal energy reserves, are at approximately equal distance from the nucleus and form an energy level.

The number of energy levels in an atom is equal to the number of the period in which the chemical element is located. Energy levels are conventionally designated as follows (for example, for Al):

Task 3.4. Determine the number of energy levels in the atoms of oxygen, magnesium, calcium, lead.

Each energy level can contain a limited number of electrons:

On the first - no more than two electrons;

On the second, no more than eight electrons;

On the third, no more than eighteen electrons.

These numbers show that, for example, at the second energy level there can be 2, 5 or 7 electrons, but there cannot be 9 or 12 electrons.

It is important to know that regardless of the energy level number on external level(the latter) cannot have more than eight electrons. The outer eight-electron energy level is the most stable and is called complete. Such energy levels are found in the most inactive elements - noble gases.

How to determine the number of electrons at the outer level of the remaining atoms? There is a simple rule for this: number of external electrons equals:

For elements of main subgroups - group number;

For elements of secondary subgroups, it cannot be more than two.

For example (fig. 5):

Task 3.5. Indicate the number of external electrons for chemical elements with serial numbers 15, 25, 30, 53.

Task 3.6. Find chemical elements in the periodic table, the atoms of which have a complete external level.

It is very important to correctly determine the number of external electrons, because it is with them that the most important properties of the atom are associated. So, in chemical reactions, atoms tend to acquire a stable, complete external level (8 e). Therefore, atoms, on the external level of which there are few electrons, prefer to give them away.

Chemical elements whose atoms are only capable of donating electrons are called metals... Obviously, there should be few electrons at the outer level of the metal atom: 1, 2, 3.

If there are many electrons on the external energy level of an atom, then such atoms tend to accept electrons until the completion of the external energy level, that is, up to eight electrons. Such elements are called non-metals.

Question. Chemical elements of secondary subgroups belong to metals or non-metals? Why?

Answer. Metals and non-metals of the main subgroups in the periodic table are separated by a line that can be drawn from boron to astatine. Above this line (and on the line) are non-metals, below - metals. All elements of side subgroups are below this line.

Task 3.7. Determine if metals or non-metals include: phosphorus, vanadium, cobalt, selenium, bismuth. Use the position of the element in the periodic table of chemical elements and the number of electrons at the outer level.

In order to compose the distribution of electrons over the remaining levels and sublevels, one should use the following algorithm.

1. Determine the total number of electrons in an atom (by ordinal number).

2. Determine the number of energy levels (by period number).

3. Determine the number of external electrons (by the type of subgroup and group number).

4. Indicate the number of electrons at all levels except the penultimate one.

For example, according to paragraphs 1-4 for a manganese atom, it is determined:

Total 25 e; distributed (2 + 8 + 2) = 12 e; which means that the third level is: 25 - 12 = 13 e.

We obtained the distribution of electrons in the manganese atom:

Task 3.8. Work out the algorithm by drawing up diagrams of the structure of atoms for elements No. 16, 26, 33, 37. Indicate whether these are metals or non-metals. Explain the answer.

Composing the above schemes of the structure of the atom, we did not take into account that the electrons in the atom occupy not only levels, but also certain sublevels each level. The types of sublevels are designated by Latin letters: s, p, d.

The number of possible sublevels is equal to the level number. The first level consists of one
s-sublevel. The second level consists of two sublevels - s and R... The third level - of three sublevels - s, p and d.

Each sublevel can contain a strictly limited number of electrons:

at the s-sublevel - no more than 2e;

on the p-sublevel - no more than 6e;

on the d-sublevel - no more than 10e.

Sublevels of one level are filled in in a strictly defined order: spd.

In this way, R-sublevel cannot start filling if it is not filled s-sublevel of a given energy level, etc. Based on this rule, it is easy to compose the electronic configuration of the manganese atom:

Generally electronic configuration of an atom manganese is written like this:

25 Mn 1 s 2 2s 2 2p 6 3s 2 3p 6 3d 5 4s 2 .

Task 3.9. Make electronic configurations of atoms for chemical elements Nos. 16, 26, 33, 37.

Why is it necessary to compose the electronic configurations of atoms? In order to determine the properties of these chemical elements. It should be remembered that only valence electrons.

Valence electrons are at an external energy level and unfinished
d-sublevel of the pre-external level.

Let us determine the number of valence electrons for manganese:

or abbreviated: Мn ... 3 d 5 4s 2 .

What can be determined by the formula for the electronic configuration of an atom?

1. Which element is it - metal or non-metal?

Manganese is a metal; the outer (fourth) level contains two electrons.

2. What process is typical for metal?

Manganese atoms in reactions always only donate electrons.

3. What electrons and how much will the manganese atom give?

In reactions, a manganese atom gives up two external electrons (they are farthest from the nucleus and are less attracted to it), as well as five pre-external d-electrons. The total number of valence electrons is seven (2 + 5). In this case, eight electrons will remain on the third level of the atom, i.e. a completed outer level is formed.

All these considerations and conclusions can be reflected using the diagram (Fig. 6):

The resulting conditional atomic charges are called oxidation states.

Considering the structure of the atom, in a similar way, one can show that the typical oxidation states for oxygen are –2, and for hydrogen +1.

Question. With which of the chemical elements can manganese form compounds, if we take into account the above oxidation states?

On the other hand, only with oxygen, because its atom has the opposite oxidation state in charge. The formulas of the corresponding manganese oxides (here, the oxidation states correspond to the valences of these chemical elements):

The structure of the manganese atom suggests that manganese cannot have a higher oxidation state, because in this case, the stable, now completed pre-external level would have to be affected. Therefore, the oxidation state +7 is the highest, and the corresponding oxide Mn 2 O 7 is the highest manganese oxide.

To consolidate all these concepts, consider the structure of the tellurium atom and some of its properties:

As a non-metal, the Te atom can accept 2 electrons before the completion of the outer level and give up the "extra" 6 electrons:

Task 3.10. Draw the electronic configurations of Na, Rb, Cl, I, Si, Sn atoms. Determine the properties of these chemical elements, the formulas of their simplest compounds (with oxygen and hydrogen).

Practical conclusions

1. Only valence electrons participate in chemical reactions, which can be located only at the last two levels.

2. Metal atoms can only donate valence electrons (all or some), assuming positive oxidation states.

3. Atoms of non-metals can accept electrons (missing - up to eight), while acquiring negative oxidation states, and donate valence electrons (all or several), while they acquire positive oxidation states.

Let us now compare the properties of chemical elements of one subgroup, for example, sodium and rubidium:
Na ... 3 s 1 and Rb ... 5 s 1 .

What is common in the structure of the atoms of these elements? At the outer level of each atom, one electron is active - these are active metals. Metallic activity associated with the ability to donate electrons: the easier an atom donates electrons, the more pronounced its metallic properties.

What keeps electrons in an atom? Attracting them to the core. The closer the electrons are to the nucleus, the more they are attracted by the nucleus of the atom, the more difficult it is to “tear off” them.

Based on this, we will answer the question: which element - Na or Rb - gives up an external electron more easily? Which element is the more active metal? Obviously rubidium, because its valence electrons are farther from the nucleus (and are less strongly held by the nucleus).

Conclusion. In the main subgroups, from top to bottom, metallic properties are enhanced since the radius of the atom increases, and the valence electrons are less attracted to the nucleus.

Let's compare the properties of chemical elements of group VIIa: Cl ... 3 s 2 3p 5 and I ... 5 s 2 5p 5 .

Both chemical elements are non-metals, because until the completion of the outer level, one electron is missing. These atoms will actively attract the missing electron. In this case, the more the missing electron attracts a non-metal atom, the more pronounced its non-metallic properties (the ability to accept electrons).

Due to what is the attraction of the electron? Due to the positive charge of the atomic nucleus. In addition, the closer the electron is to the nucleus, the stronger their mutual attraction, the more active the non-metal.

Question. Which element has more pronounced non-metallic properties: chlorine or iodine?

Answer. Obviously, chlorine, because its valence electrons are located closer to the nucleus.

Conclusion. The activity of non-metals in subgroups from top to bottom decreases since the radius of the atom increases and it becomes more and more difficult for the nucleus to attract the missing electrons.

Let's compare the properties of silicon and tin: Si ... 3 s 2 3p 2 and Sn ... 5 s 2 5p 2 .

At the outer level of both atoms, four electrons each. Nevertheless, these elements in the periodic table are on opposite sides of the line connecting boron and astatine. Therefore, silicon, the symbol of which is located above the B – At line, exhibits stronger nonmetallic properties. On the other hand, tin, the symbol of which is below the B – At line, exhibits stronger metallic properties. This is because in the tin atom, four valence electrons are distant from the nucleus. Therefore, attaching the missing four electrons is difficult. At the same time, the release of electrons from the fifth energy level occurs quite easily. For silicon, both processes are possible, with the first (electron reception) prevailing.

Conclusions for chapter 3. The fewer external electrons in an atom and the farther they are from the nucleus, the more pronounced the metallic properties are.

The more external electrons in the atom and the closer they are to the nucleus, the more pronounced non-metallic properties are.

Based on the conclusions formulated in this chapter, a "characteristic" can be drawn up for any chemical element of the periodic table.

Algorithm for describing properties
chemical element by its position
in the periodic system

1. Draw up a diagram of the structure of the atom, i.e. determine the composition of the nucleus and the distribution of electrons by energy levels and sublevels:

Determine the total number of protons, electrons and neutrons in an atom (by ordinal number and relative atomic mass);

Determine the number of energy levels (by period number);

Determine the number of external electrons (by the type of subgroup and group number);

Indicate the number of electrons at all energy levels, except for the penultimate;

2. Determine the number of valence electrons.

3. Determine which properties - metal or non-metal - are more pronounced in a given chemical element.

4. Determine the number of donated (received) electrons.

5. Determine the highest and lowest oxidation states of a chemical element.

6. Compile for these oxidation states the chemical formulas of the simplest compounds with oxygen and hydrogen.

7. Determine the nature of the oxide and draw up an equation for its reaction with water.

8. For the substances specified in paragraph 6, draw up the equations of typical reactions (see Chapter 2).

Task 3.11. According to the above scheme, compose descriptions of the atoms of sulfur, selenium, calcium and strontium and the properties of these chemical elements. What are the general properties of their oxides and hydroxides?

If you have completed exercises 3.10 and 3.11, then it is easy to notice that not only the atoms of the elements of one subgroup, but also their compounds, have common properties and a similar composition.

D. I. Mendeleev's periodic law:the properties of chemical elements, as well as the properties of simple and complex substances formed by them, are periodically dependent on the charge of the nuclei of their atoms.

The physical meaning of the periodic law: the properties of chemical elements are periodically repeated because the configurations of valence electrons (the distribution of electrons of the outer and penultimate levels) are periodically repeated.

So, chemical elements of the same subgroup have the same distribution of valence electrons and, therefore, similar properties.

For example, chemical elements of the fifth group have five valence electrons. Moreover, in the atoms of chemical elements of main subgroups- all valence electrons are at the external level: ... ns 2 np 3, where n- period number.

At atoms elements of secondary subgroups there are only 1 or 2 electrons on the external level, the rest are on d-sub-level of the pre-external level: ... ( n – 1)d 3 ns 2, where n- period number.

Task 3.12. Make short electronic formulas for atoms of chemical elements No. 35 and 42, and then make up the distribution of electrons in these atoms according to the algorithm. Make sure your prediction comes true.

Exercises for Chapter 3

1. Formulate the definitions of the concepts "period", "group", "subgroup". What do the chemical elements have in common: a) period; b) a group; c) a subgroup?

2. What are isotopes? What properties - physical or chemical - are the same for isotopes? Why?

3. Formulate the periodic law of D.I. Mendeleev. Explain its physical meaning and illustrate with examples.

4. What is the manifestation of the metallic properties of chemical elements? How do they change in the group and in the period? Why?

5. What is the manifestation of the non-metallic properties of chemical elements? How do they change in the group and in the period? Why?

6. Make short electronic formulas of chemical elements No. 43, 51, 38. Confirm your assumptions by describing the structure of the atoms of these elements according to the above algorithm. Specify properties for these elements.

7. By short electronic formulas

a) ... 4 s 2 4p 1;

b) ... 4 d 1 5s 2 ;

at 3 d 5 4s 1

determine the position of the corresponding chemical elements in the periodic table of D.I. Mendeleev. Name these chemical elements. Confirm your assumptions by describing the structure of the atoms of these chemical elements according to the algorithm. Indicate the properties of these chemical elements.

To be continued

MBOU "Gymnasium No. 1 of the city of Novopavlovsk"

Chemistry grade 8

Topic:

"Change in the number of electrons

on the external energy level

atoms of chemical elements "

Teacher: Tatiana Alekseevna Komarova

Novopavlovsk

Date: ___________

Lesson– 9

Lesson topic: Change in the number of electrons on the external energy

the level of atoms of chemical elements.

Lesson objectives:

To form a concept about the metallic and non-metallic properties of elements at the atomic level;

Show the reasons for the change in the properties of elements in periods and groups based on the structure of their atoms;

Provide an initial understanding of ionic bonding.

Equipment: PSKhE, table "Ionic bond".

During the classes

Organizing time.

Knowledge check

Characteristics of chemical elements according to the table (3 people)

Atomic structure (2 persons)

Learning new material

Consider the following questions:

1 ... Atoms, of which chemical elements, have complete energy levels?

These are the atoms of inert gases, which are located in the main subgroup of the 8th group.

The completed electronic layers have increased robustness and stability.

Atoms of group VIII (He Ne Ar Kr Xe Rn) contain 8е - at the external level, which is why they are inert, i.e. ... chemically inactive, do not interact with other substances, i.e. their atoms have increased stability and stability. That is, all chemical elements (having a different electronic structure) tend to get completed external energy level , 8e -.

Example:

11 +12 +9 +17

2 8 1 2 8 2 2 7 2 8 7

1 s 2 2s 2 p 6 3 s 1 1s 2 2s 2 p 6 3 s 2 1s 2 2s 2 p 5 1s 2 2s 2 p 6 3 s 2 p 5

How do you think the atoms of these elements can reach eight electrons on the outer level?

If (suppose) the hand closes the last level at Na and Mg, complete levels are obtained. Therefore, these electrons must be donated from the external electronic level! Then, with the donation of electrons, the pre-outer layer from 8e -, becomes external.

And for the elements F and Cl, you should take 1 missing electron to your energy level, than give 7e -. And so, there are 2 ways to achieve a completed energy level:

A) Recoil of ("extra") electrons from the outer layer.

B) Acceptance of ("missing") electrons to the external level.

2. The concept of metallicity and non-metallicity at the atomic level:

Metals Are elements whose atoms donate their outer electrons.

Non-metals - these are elements whose atoms take electrons to the external energy level.

The easier the Me atom gives up its electrons, the more pronounced it is. metallic properties.

The easier the HeMe atom accepts the missing electrons to the outer layer, the more pronounced it is. non-metallic properties.

3. Change of Ме and НеМе properties of CHE atoms. in periods and groups in the PSCE.

In periods:

Example: Na (1e -) Mg (2e -) - write down the structure of the atom.

What element do you think has stronger metallic properties? Na or Mg? Which is easier to give 1e - or 2e -? (Of course 1e -, therefore, Na has more pronounced metallic properties).

Example: Al (3e -) Si (4e -) etc.

Over the period, the number of electrons at the outer level increases from left to right.

(brighter metallic properties are expressed in Al).

Of course, the ability to donate electrons over the period will decrease, i.e. the metallic properties will weaken.

Thus, the strongest Me are located at the beginning of periods.

How will the ability to attach electrons change? (will increase)

Example:

14 r +17 r

2 8 4 2 8 7

It is easier to accept 1 missing electron (y Cl) than 4e for Si.

Conclusion:

From left to right, non-metallic properties will increase, and metallic properties will weaken.

One more reason for the enhancement of HEM properties is a decrease in the radius of the atom with a constant number of levels.

Because within the first period, the number of energy levels for atoms does not change, but the number of external electrons e - and the number of protons p - in the nucleus increase. As a result, the attraction of electrons to the nucleus increases (Coulomb's law), and the radius ( r) of the atom decreases, the atom seems to be compressed.

General conclusion:

Within one period with the growth of the ordinal number ( N) of the element, the metallic properties of the elements weaken, and the non-metallic properties increase, because:

The number e is growing - at the external level it is equal to the group number and the number of protons in the nucleus.

Atom radius decreases

The number of energy levels is constant.

4. Consider the vertical dependence of changes in the properties of elements (within the main subgroups) in groups.

Example: VII group main subgroup (halogens)

9 +17

2 7 2 8 7

1 s 2 2s 2 p 5 1s 2 2s 2 p 6 3s 2 p 5

The number e - at the outer levels of these elements is the same, but the number of energy levels is different,

for F -2e -, and Cl - 3e - /

Which atom has a larger radius? (- for chlorine, because there are 3 energy levels).

The closer e are to the nucleus, the stronger they are attracted to it.

The atom of which element will be easier to attach e - y F or Cl?

(F - easier to attach 1 missing electron), because it has a smaller radius, which means that the force of attraction of the electron to the nucleus is greater than that of Cl.

Coulomb's law

The force of interaction of two electric charges is inversely proportional to the square

the distance between them, i.e. the greater the distance between the atoms, the less the force

attraction of two opposite charges (in this case, electrons and protons).

F is stronger than Cl ˃Br˃J, etc.

Conclusion:

In groups (main subgroups), non-metallic properties are reduced, and metallic properties are enhanced, because:

one). The number of electrons at the outer level of atoms is the same (and is equal to the group number).

2). The number of energy levels in atoms is growing.

3). The radius of the atom increases.

Verbally, according to the PSCE table, consider I - group main subgroup. Conclude that the strongest metal is Fr francium, and the strongest non-metal is F fluorine.

Ionic bond.

Consider what happens to the atoms of the elements if they reach the octet (i.e. 8e -) at the outer level:

Let's write out the formulas of the elements:

Na 0 +11 2e - 8e - 1e - Mg 0 +12 2e - 8e - 2e - F 0 +9 2e - 7e - Cl 0 +17 2e - 8e - 7e -

Na x +11 2e - 8e - 0e - Mg x +12 2e - 8e - 0e - F x +9 2e - 8e - Cl x +17 2e - 8e - 8e -

The top row of formulas contains the same number of protons and electrons, since these are the formulas of neutral atoms (there is a zero charge "0" - this is the oxidation state).

Bottom row - different number p + and e -, i.e. these are formulas for charged particles.

Let's calculate the charge of these particles.

Na +1 +11 2e - 8e - 0e - 2 + 8 = 10, 11-10 = 1, oxidation state +1

F - +9 2e - 8e - 2 + 8 = 10, 9-10 = -1, oxidation state -1

Mg +2 +12 2е - 8e - 0e - 2 + 8 = 10, 12-10 = -2, oxidation state -2

As a result of the addition - recoil of electrons, charged particles are obtained, which are called ions.

Me atoms upon recoil e - acquires "+" (positive charge)

NotMe atoms accepting "foreign" electrons are charged "-" (negative charge)

The chemical bond that forms between ions is called ionic.

The ionic bond arises between strong Me and strong Me.

Examples.

a) the formation of an ionic bond. Na + Cl -

N a Cl + -

11 + +17 +11 +17

2 8 1 2 8 7 2 8 2 8 8

1 e -

The process of converting atoms into ions:

Na 0 + Cl 0 Na + + Cl - Na + Cl -

atom atom ion ion ionic compound

2e -

b) Ca O 2+ 2-

20 +8 +20 +8

2 8 8 2 2 6 2 8 8 2 8

Ca a 0 - 2e - Ca 2+ 2 1

Lesson summary

Literature:

1. Chemistry grade 8. textbook for general education

institutions / OS. Gabrielyan. Bustard 2009

2. Gabrielyan O.S. Handbook of the teacher.

Chemistry grade 8, Bustard, 2003

At the external energy level of the atoms of iron, cobalt and nickel there are 2 electrons each. At the d-sublevel of the penultimate energy level, iron, cobalt and nickel have 6, 7, and 8 electrons, respectively. The characteristic oxidation states of metals of the iron family are +2 and +3 (compounds are known in which they exhibit oxidation states of +1, +4 and +6, for example, potassium ferrate K 2 FeO 4, but there are few such compounds and they are not typical). For iron, compounds with the oxidation state (+3) are more stable, and for nickel and cobalt - (+2). Therefore, Fe 2+ is a rather strong reducing agent, while Ni 2+ and Co 2+ do not possess these properties to an appreciable extent; cobalt and nickel compounds are quite stable in air. In the +3 oxidation state, iron, cobalt, nickel exhibit oxidizing properties, the oxidizing ability increases in the series Fe 3+ - Ni 3+ - Co 3+.

Iron, cobalt and nickel are very similar in properties to each other (ferromagnetic, catalytic activity, ability to form colored ions, complexation). However, there are also differences between them: iron in its magnetic properties is allocated in a triad, the reducing activity of iron is much higher than that of cobalt and nickel, which are closer to tin than to iron in terms of their electrode potentials.

When heated, metals of the iron family interact vigorously with metalloids, for example, chlorine, bromine, oxygen, sulfur, etc. Chemically pure iron, cobalt and nickel do not change when exposed to air and water. However, ordinary iron contains various impurities, so it corrodes in humid air. The resulting rust layer is brittle and porous, it does not prevent the metal from contacting the environment and does not protect it from further oxidation. At high temperatures, iron interacts with water, displacing hydrogen from it. Iron dissolves easily in dilute acids; cobalt, nickel - much more difficult.

At a high concentration of acids in the cold, iron is passivated, becoming covered with the thinnest oxide film. Oxides of all three metals (FeO, CoO, NiO) are insoluble in water. Their hydrates are obtained by the action of alkali on soluble salts. Oxide hydrates exhibit basic properties. Fe (OH) 2 hydroxide, interacting with atmospheric oxygen and water, is rapidly oxidized:

4Fе (OH) 2 + O 2 + 2H 2 O = 4Fe (OH) 3.

Oxidation of Co 2+ and especially Ni 2+ ions is a little more difficult. Of the oxides and hydroxides of Fe, Co, Ni, only Fe 2 O 3 and Fe (OH) 3 are amphoteric with a predominance of basic properties. Cobalt and nickel oxides and hydroxides are strong oxidizing agents; when interacting with acids, they are reduced to salts of bivalent metals:

Co 2 O 3 + 6HC1 = 2CoC1 2 + Cl 2 + 3H 2 O;

4Ni (OH) 3 + 4H 2 SO 4 = 4NiSO 4 + О 2 + 10H 2 O

Fe 3+ compounds are weak oxidizing agents and, under the action of reducing agents, transform into Fe 2+ derivatives:

H 2 S + Fe 2 (SO 4) 3 = S + 2FeSO 4 + H 2 SO 4

Many simple and complex ions of the elements iron, cobalt and nickel are colored. Thus, hydrated Co 2+ ions are pink, Ni 2+ are green, Fe 3+ in an aqueous solution has a brown-yellow color due to hydrolysis.