Lesson summary "Change in the number of electrons at the external energy level of atoms of chemical elements"

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 level. In our article we will look at the electronics 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

When passing 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 levels of the atom farthest 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 electrons are most likely to be found is called an 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 the number of electrons in the external energy level for the 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 in 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 for them there is the following rule: the number of external negative particles is equal to 1 (for chromium and copper) or 2. This is explained by the fact that as the charge of the atomic nucleus increases, the internal d-sublevel is first filled and the external energy levels remain unchanged.

Why do the properties of elements of short periods change?

The 1st, 2nd, 3rd and 7th periods are considered small. The 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 that can be easily removed from the nucleus. In this case, a positively charged metal ion is formed.

Amphoteric elements, for example, aluminum or zinc, fill their outer 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 to the octet, attracting electrons from other atoms. For example, the nonmetal with the highest electronegativity, fluorine, has last layer 7 electrons and always takes one electron not only from metals, but also from active non-metallic elements: oxygen, chlorine, nitrogen. Small periods, like large ones, end with inert gases, whose monatomic molecules have completely completed outer energy levels up to 8 electrons.

Features of the structure of atoms of long periods

The even rows of periods 4, 5, and 6 consist of elements whose outer shells accommodate only one or two electrons. As we said earlier, they fill the d- or f-sublevels of the penultimate layer with electrons. Usually these are typical metals. Physical and Chemical properties they change very slowly. Odd rows contain elements whose outer energy levels are filled with electrons according to the following scheme: metals - amphoteric element - nonmetals - 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, the atoms of which have a completely 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 positions of electrons on the external energy level of atoms of elements. It turned out that for elements of the main subgroups, for example, lithium, sodium, potassium, rubidium and cesium, the last electron is located on the s-sublevel. Group 7 elements of the main subgroup (halogens) fill their p-sublevel with negative particles.

For representatives of side subgroups, such as chromium, filling the d-sublevel with electrons will be typical. And for elements included in the families, the accumulation of negative charges occurs 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 outer energy levels of atoms have chemical elements, and determined their role in interatomic interactions.

“Types of chemical bonds” - Crystals are hard, refractory, odorless, insoluble in water. EO in the period increases EO in the group increases the MOST electronegative element fluorine. Substances are fusible and often have an odor. AN IONIC BOND formed as a result of electrostatic attraction. The atomic frame has high strength.

“Metallic chemical bond” - The best conductors are copper and silver. Mercury, silver, palladium, and aluminum have high reflectivity. Differences between metallic bonds and ionic and covalent bonds. A metallic bond has features similar to a covalent bond. A metallic bond has something in common with: An ionic bond – the formation of ions. Gold products.

“Chemistry “Chemical Bond”” - Substances with covalent bonds. Covalent bond parameters. Hydrogen chemical bond. Two types of crystal lattices. Metals form metallic crystal lattices. Ionic bonding is an electrostatic attraction between ions. Sharp boundaries between different types there are no chemical bonds. Covalent bond.

“Covalent polar bond” - Electron pairs. Atoms. View chemical bond. Write down the electronic and structural formulas. Formation of the concept of covalent chemical bonds. Metals and non-metals. Elements. Poles. Communication type. Covalent polar chemical bond. Electronegativity series. Increased electronegativity. Common electron pairs.

“Hydrogen chemical bond” - The appearance of a new absorption band in the electronic spectra. Complexes with elements of the 6th group. Properties of covalent chemical bonds. States of molecular complexes of composition DA. Complexes of two types. Dispersive energy. Donor-acceptor bond. Symmetrical. Dependence of energy on distance. Two molecules are described by the Hamiltonians HA and HB.

“Types and characteristics of chemical bonds” - Ionic bond. Covalent polar. Covalent polar connection. Covalent bond. Substances with a molecular crystal lattice. Metal connection. Hydrogen bond. Properties of substances. Connection. Molecular and atomic crystal lattices. Properties of substances with metallic bonds. Ionic crystal lattices.

There are a total of 23 presentations in the topic

Chemistry lesson in 8th grade. "_____"______ 20_____

Change in the number of electrons at the external energy level of atoms of chemical elements.

Target. Consider changes in the properties of atoms of chemical elements in PSHE D.I. Mendeleev.

Educational. Explain the patterns of changes in the properties of elements within small periods and main subgroups; determine the reasons for changes in metallic and non-metallic properties in periods and groups.

Developmental. Develop the ability to compare and find patterns of changes in properties in PSHE D.I. Mendeleev.

Educational. Foster a culture of academic work in the classroom.

During the classes.

1. Org. moment.

2. Repetition of the studied material.

Independent work.

Option 1.

6-10. Indicate the number of energy levels in the atoms of the following elements.

Option 2.

1-5. Indicate the number of neutrons in the nucleus of an atom.

6-10. Indicate the number of electrons in the outer energy level.

11-15. The indicated electronic formula of the atom corresponds to the element.

3. Studying a new topic.

Exercise. Distribute the electrons among the energy levels of the following elements: Mg, S, Ar.

The completed electronic layers have increased robustness and stability. Atoms that have 8 electrons in their outer energy level - inert gases - are stable.

An atom will always be stable if it has 8ē at its external energy level.

How can atoms of these elements reach the 8-electron outer level?

2 ways to complete:

Donate electrons

Accept electrons.

Metals are elements that donate electrons; at their outer energy level they have 1-3 ē.

Nonmetals are elements that accept electrons; their outer energy level is 4-7ē.

Changing properties in PSHE.

Within one period, as the atomic number of an element increases, the metallic properties weaken and the nonmetallic properties increase.

1. The number of electrons at the external energy level increases.

2. The radius of the atom decreases

3. The number of energy levels is constant

In the main subgroups, non-metallic properties decrease, and metallic properties increase.

1. The number of electrons at the external energy level is constant;

2. The number of energy levels increases;

3. The radius of the atom increases.

Thus, francium is the strongest metal, fluorine is the strongest non-metal.

4. Consolidation.

Exercises.

1. Arrange these chemical elements in order of increasing metallic properties:

A) Al, Na, Cl, Si, P

B) Mg, Ba, Ca, Be

B) N, Sb, Bi, As

D) Cs, Li, K, Na, Rb

2. Arrange these chemical elements in order of increasing non-metallic properties:

B) C, Sn, Ge, Si

B) Li, O, N, B, C

D) Br, F, I, Cl

3. Underline the symbols for chemical metals:

A) Cl, Al, S, Na, P, Mg, Ar, Si

B) Sn, Si, Pb, Ge, C

Arrange in order of decreasing metallic properties.

4. Underline the symbols of the chemical elements of non-metals:

A) Li, F, N, Be, O, B, C

B) Bi, As, N, Sb, P

Arrange in order of decreasing nonmetallic properties.

Homework. Page 61- 63. Ex. 4 page 66

Each period of D.I. 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 pure form earlier than other analogues. What is the reason for the inertness of helium, neon, argon, krypton, xenon and radon? The fact is that atoms of inert gases have eight electrons at the outermost levels from the nucleus (helium has two). Eight electrons at the outer level is the limiting number for each element of D.I. Mendeleev’s Periodic Table, except 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 D.I. Mendeleev’s Periodic Table strive.

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

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 to a greater extent It has metallic properties. It is clear, therefore, that the most typical metals in D.I. Mendeleev’s Periodic Table are the elements of the main subgroup of group I (group IA). Conversely, atoms of non-metal elements tend to accept those missing before the completion of the external energy level. From the above we can draw the following conclusion. 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 nonmetallic properties of elements, characterized by the ease of accepting electrons to the external level, are enhanced.

The most typical non-metals are the elements of the main subgroup of group VII (group VIIA) of D. I. Mendeleev’s Periodic Table. The outer level of the atoms of these elements contains seven electrons. Up to eight electrons at the external level, i.e., to the stable state of atoms, they lack one electron. They easily attach them, exhibiting non-metallic properties.

How do atoms of elements of the main subgroup of group IV (group IVA) of D.I. Mendeleev’s periodic system behave? After all, they have four electrons on the outer level, and it would seem that they don’t care whether they give or take four electrons. It turned out that the ability of atoms to donate or accept 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 shrink. Therefore, it becomes increasingly difficult to give up external electrons and, conversely, it becomes increasingly easier to accept the missing up to eight electrons.

Within the same subgroup, the radius of an atom increases with increasing charge of the atomic nucleus, since when constant number electrons in the outer level (it is equal to the group number), the number of energy levels increases (it is equal to the period number). Therefore, it becomes increasingly easier for the atom to give up its outer electrons.

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

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

Let’s imagine that two atoms “meet”: a Group IA metal atom and a Group VIIA nonmetal atom. A metal atom has a single electron at its outer energy level, while a non-metal atom just lacks one electron for its outer level to be complete.

A metal atom will easily give up its electron, farthest from the nucleus and weakly bound to it, to a non-metal atom, which will give it free place at its external energy level.

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

Both atoms will fulfill their " cherished dream" - will receive the much-coveted eight electrons at the outer energy level. But what happens next? Oppositely charged ions, in full accordance with the law of attraction of opposite charges, will immediately unite, i.e., a chemical bond will arise between them.

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

The process of converting atoms into ions is depicted in the diagram and figure:

For example, ionic bond is also formed by the interaction of calcium and oxygen atoms:

This transformation of atoms into ions always occurs during the interaction of atoms of typical metals and typical non-metals.

In conclusion, let us consider the algorithm (sequence) of reasoning when writing the 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 D.I. Mendeleev’s Periodic Table, a metal. It is easier for its atom to give away two outer electrons than to accept the missing six:

2. Chlorine is an element of the main subgroup of group VII (group VIIA) of D.I. Mendeleev’s table, a non-metal. It is easier for its atom to accept one electron, which it lacks to complete the outer energy level, than to give away seven electrons from the outer level:

3. First, let's find the least common multiple between the charges of the resulting ions; it is equal to 2 (2×1). Then we determine how many calcium atoms need to be taken so that they can give up two electrons (i.e., you need to take 1 Ca atom), and how many chlorine atoms need to be taken so that they can accept two electrons (i.e., you 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 - analogues of molecular formulas.

Numbers showing the number of atoms, molecules or formula units are called coefficients, and 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 changes in the properties of elements. In the second part of the paragraph we present the key words.

Key words and phrases

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

Work with computer

1. Refer to the electronic application. Study the lesson material and complete the assigned tasks.
2. Find email addresses on the Internet that can serve as additional sources that reveal the content of keywords and phrases in the paragraph. Offer your help to the teacher in preparing a new lesson - send a message by 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 ionic bonds between 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 D. I. Mendeleev’s Periodic Table.
4. Using additional sources of information, explain why inert gases came to be called noble gases.

Chemistry lesson in 8th grade. "_____"______ 20_____

Change in the number of electrons at the external energy level of atoms of chemical elements.

Target. Consider changes in the properties of atoms of chemical elements in PSHE D.I. Mendeleev.

Educational. Explain the patterns of changes in the properties of elements within small periods and main subgroups; determine the reasons for changes in metallic and non-metallic properties in periods and groups.

Developmental. Develop the ability to compare and find patterns of changes in properties in PSHE D.I. Mendeleev.

Educational. Foster a culture of academic work in the classroom.

During the classes.

Org. moment.

Repetition of learned material.

Independent work.

Option 1.

Option 2.

1-5. Indicate the number of neutrons in the nucleus of an atom.

11-15. The indicated electronic formula of the atom corresponds to the element.

Studying a new topic.

Exercise. Distribute the electrons among the energy levels of the following elements: Mg, S, Ar.

The completed electronic layers have increased robustness and stability. Atoms that have 8 electrons in their outer energy level - inert gases - are stable.

An atom will always be stable if it has 8ē at its external energy level.

How can atoms of these elements reach the 8-electron outer level?

2 ways to complete:

Donate electrons

Accept electrons.

Metals are elements that donate electrons; at their outer energy level they have 1-3 ē.

Nonmetals are elements that accept electrons; their outer energy level is 4-7ē.

Changing properties in PSHE.

Within one period with increasing serial number element, the metallic properties are weakened, and the non-metallic properties are enhanced.

The number of electrons at the outer energy level increases.

The radius of the atom decreases

The number of energy levels is constant.

In the main subgroups non-metallic properties decrease, and metallic properties increase.

The number of electrons in the outer energy level is constant;

The number of energy levels increases;

The radius of the atom increases.

Thus, francium is the strongest metal, fluorine is the strongest non-metal.

Consolidation.

Exercises.

Arrange these chemical elements in order of increasing metallic properties:

A) Al, Na, Cl, Si, P

B) Mg, Ba, Ca, Be

B) N, Sb, Bi, As

D) Cs, Li, K, Na, Rb

Arrange these chemical elements in order of increasing nonmetallic properties:

B) C, Sn, Ge, Si

B) Li, O, N, B, C

D) Br, F, I, Cl

Underline the chemical metal symbols:

A) Cl, Al, S, Na, P, Mg, Ar, Si

B) Sn, Si, Pb, Ge, C

Arrange in order of decreasing metallic properties.

Underline the symbols of the chemical elements of nonmetals:

A) Li, F, N, Be, O, B, C

B) Bi, As, N, Sb, P

Arrange in order of decreasing nonmetallic properties.

Homework. Page 61- 63. Ex. 4 page 66

Option 1.

1-5. Indicate the number of neutrons in the nucleus of an atom.

6-10. Indicate the number of energy levels in the atoms of the following elements.

11-15. The indicated electronic formula of the atom corresponds to the element.

Option 2.

1-5. Indicate the number of neutrons in the nucleus of an atom.

6-10. Indicate the number of electrons in the outer energy level.

11-15. The indicated electronic formula of the atom corresponds to the element.