Brief lecture on the structure of electronic shells of atoms college. Electronic shell of an atom. State of electrons in an atom

The outstanding Danish physicist Niels Bohr (Fig. 1) suggested that electrons in an atom can move not in any, but in strictly defined orbits.

Rice. 1. Bohr Niels Hendrich David (1885-1962)

In this case, the electrons in an atom differ in their energy. As experiments show, some of them are attracted to the nucleus more strongly, others - less. The main reason for this is the different distances of electrons from the nucleus of an atom. The closer the electrons are to the nucleus, the more tightly they are bound to it and the more difficult it is to tear them out of the electron shell. Thus, as the electron moves away from the nucleus of the atom, the energy reserve of the electron increases.

Electrons moving near the nucleus seem to block (screen) the nucleus from other electrons, which are attracted to the nucleus less strongly and move at a greater distance from it. This is how electronic layers are formed.

Each electron layer consists of electrons with similar energy values; Therefore, electronic layers are also called energy levels.

The nucleus is at the center of each element's atom, and the electrons, which form the electron shell, are arranged in layers around the nucleus.

The number of electron layers in an element's atom is equal to the number of the period in which the element is located.

For example, sodium Na is an element of the 3rd period, which means that its electron shell includes 3 energy levels. The bromine atom Br has 4 energy levels, since bromine is located in the 4th period (Fig. 2).

Sodium atom model: Bromine atom model:

The maximum number of electrons at an energy level is calculated by the formula: 2n2, where n is the number of the energy level.

Thus, the maximum number of electrons per:

3rd layer - 18, etc.

For elements of the main subgroups, the number of the group to which the element belongs is equal to the number of outer electrons of the atom.

The outer electrons are the electrons of the last electron layer.

For example, the sodium atom has 1 outer electron (since it is an element of the IA subgroup). The bromine atom has 7 electrons in the last electron layer (this is an element of subgroup VIIA).

Structure of electronic shells of elements of periods 1-3

In a hydrogen atom, the nuclear charge is +1, and this charge is neutralized by a single electron (Fig. 3).

The next element after hydrogen is helium, also an element of the 1st period. Consequently, in a helium atom there is 1 energy level, which contains two electrons (Fig. 4). This is the maximum possible number of electrons for the first energy level.

Element #3 is lithium. There are 2 electron layers in a lithium atom, since it is an element of the 2nd period. On the 1st layer in a lithium atom there are 2 electrons (this layer is completed), and on the 2nd layer there is 1 electron. The beryllium atom has 1 more electron than the lithium atom (Fig. 5).

Similarly, one can depict the atomic structure diagrams of the remaining elements of the second period (Fig. 6).

In the atom of the last element of the second period - neon - the last energy level is complete (it has 8 electrons, which corresponds to the maximum value for the 2nd layer). Neon is an inert gas that does not enter into chemical reactions, therefore its electron shell is very stable.

American chemist Gilbert Lewis gave an explanation for this and put forward octet rule, according to which the eight-electron layer is stable(with the exception of 1 layer: since it can contain no more than 2 electrons, a two-electron state will be stable for it).

After neon comes the element of the 3rd period - sodium. The sodium atom has 3 electron layers, on which 11 electrons are located (Fig. 7).

Rice. 7. Scheme of the structure of the sodium atom

Sodium is in group 1, its valence in compounds is equal to I, like lithium. This is due to the fact that there is 1 electron in the outer electron layer of the sodium and lithium atoms.

The properties of elements repeat periodically because the atoms of elements periodically repeat the number of electrons in their outer electron layer.

The structure of the atoms of the remaining elements of the third period can be represented by analogy with the structure of the atoms of the elements of the 2nd period.

The structure of electronic shells of elements of the 4th period

The fourth period includes 18 elements, among them there are elements of both the main (A) and secondary (B) subgroups. A peculiarity of the structure of atoms of elements of side subgroups is that their outer (internal) rather than outer electronic layers are sequentially filled.

The fourth period begins with potassium. Potassium is an alkali metal that exhibits valency I in compounds. This is quite consistent with the following structure of its atom. As a 4th period element, the potassium atom has 4 electron layers. The last (fourth) electron layer of potassium contains 1 electron, the total number of electrons in a potassium atom is 19 (the serial number of this element) (Fig. 8).

Rice. 8. Scheme of the structure of the potassium atom

Potassium is followed by calcium. The calcium atom will have 2 electrons on its outer electron layer, just like beryllium and magnesium (they are also elements of the II A subgroup).

The next element after calcium is scandium. This is an element of the secondary (B) subgroup. All elements of secondary subgroups are metals. A feature of the structure of their atoms is the presence of no more than 2 electrons in the last electronic layer, i.e. the penultimate electron layer will be sequentially filled with electrons.

Thus, for scandium we can imagine the following model of atomic structure (Fig. 9):

Rice. 9. Scheme of the structure of the scandium atom

This distribution of electrons is possible because on the third layer the maximum permissible number of electrons is 18, i.e. eight electrons on the 3rd layer is a stable, but not complete, state of the layer.

For ten elements of secondary subgroups of the 4th period from scandium to zinc, the third electron layer is sequentially filled.

The structure of a zinc atom can be represented as follows: there are two electrons on the outer electronic layer, and 18 on the outer one (Fig. 10).

Rice. 10. Scheme of the structure of the zinc atom

The elements following zinc belong to the elements of the main subgroup: gallium, germanium, etc. up to krypton. In the atoms of these elements, the 4th (i.e., outer) electron layer is sequentially filled. In an atom of the noble gas krypton there will be an octet on the outer shell, i.e. a stable state.

Summing up the lesson

In this lesson, you learned how the electron shell of an atom is structured and how to explain the phenomenon of periodicity. We got acquainted with models of the structure of the electronic shells of atoms, with the help of which we can predict and explain the properties of chemical elements and their compounds.

Sources

http://www.youtube.com/watch?t=7&v=xgPDyORYV_Q

http://www.youtube.com/watch?t=416&v=BBmhmB4ans4

http://www.youtube.com/watch?t=10&v=6Y19QgS5V5E

http://www.youtube.com/watch?t=3&v=B6XEB6_gbdI

presentation source - http://www.myshared.ru/slide/834600/#

Abstract http://interneturok.ru/ru/school/chemistry/8-klass

Let us mentally take an atom of any chemical element. In what states are the electrons in it? From the previous paragraph it is clear that for each electron it is necessary to know the values ​​of four quantum numbers that characterize its state. But we don't yet know how many electrons are in each state. Which conditions are more and which are less likely? The answer to these questions is provided by two important principle(law). The first of them was discovered in 1925 by the Swiss physicist W. Pauli (1900-1958) and named after him - Pauli's principle.

All electrons in an atom are in different states, i.e. characterized by different sets of four quantum numbers.

In this case, the concept of “principle” denotes one of the fundamental laws of nature, which makes the atom what it is - a microparticle of matter with an individual electronic structure for each chemical element. The role of the Pauli principle in nature becomes clearer if we imagine that it does not operate. Then the electronic environment of the atomic nucleus loses its structural definiteness. All electrons roll into one most favorable state.

It should be noted that this law is valid for all fermions.

A corollary follows from the Pauli principle that determines the capacity of the orbital, i.e. the number of electrons that can form a single electron cloud. By choosing any of the orbitals, we fix the first three quantum numbers. For example, for orbital 2 p 2: p = 2, /= 1, mj= 0. But you can also change the spin quantum number m s We get two sets of quantum numbers:

Therefore, an orbital can hold no more than two electrons, and atoms can contain one- and two-electron clouds.

Two electrons in the same orbital are called an electron pair.

Knowing the capacity of the orbital, it is easy to understand that the capacity of the energy sublevel is equal to twice the number of orbitals (Table 5.1).

Table 5.1

Structure of sublevels in atoms

The collection of electrons of the same energy sublevel is called the subshell of an atom.

The capacity of the energy level is the sum of the capacity of the sublevels (Table 5.2). In the first column of the table, in addition to the values ​​of the principal quantum number, the letter designations for the electron shells of the atom are given.

Table 5.2

Structure of energy levels in atoms

The collection of electrons of the same energy level is called the shell of an atom.

The actual filling (“population”) of orbitals, sublevels and levels with electrons is determined by the second principle - principle of least energy.

The ground (stable) state of the atom corresponds to the minimum total energy of electrons.

States of an atom with increased energy are called excited. An atom in an excited state is unstable in the sense that in a very short time (~10 -8 s) it goes into the ground state, emitting energy quanta.

Any physical system is more stable the lower its potential energy. Therefore, we invariably observe that an thrown body hits the ground or rolls down a hill, a bent spring straightens, etc. Also, the electron shells of atoms are in a stable state if the total energy supply of electrons is minimal. We already know the set of possible energy states of an atom (see Fig. 5.7). Let us consider how the corresponding sublevels and levels are populated by electrons. In this case, the Pauli principle is strictly followed, which has priority over the principle of least energy and is not violated. We will depict the electronic structure of atoms using energy diagrams and electronic formulas. The energy diagram is part of the overall sequence of sublevels (see Fig. 5.7), containing populated sublevels. The electron formula lists the populated sublevels in order of increasing energy, indicating the number of electrons with superscripts. The first two elements of the periodic table can be represented by diagrams I and II. The diagram shows that the position of the l* level in the helium atom is lower than in the hydrogen atom, since helium has a larger nuclear charge and electrons are more strongly attracted to the nucleus. The capacity of the first energy level in the helium atom is exhausted.

The second energy level of the elements following helium is populated. Let's consider the energy diagrams of the three closest elements - lithium, beryllium and boron (diagrams III, IV and V).

In lithium and beryllium, the sublevel is populated 2s. The fifth electron of the boron atom begins to populate sublevel 2 R in accordance with the Pauli principle. For carbon and nitrogen atoms, the population of this sublevel continues (diagrams VI and VII).

The structure of these elements reveals another important pattern in the formation of electron shells - Hund's rule (1927).

The basic 7 state of the atom corresponds to the occupation of the maximum number of energetically equivalent orbitals by electrons. In this case, the electrons have the same spin quantum numbers (all +1/2 or all -1/2).

When considering the energy diagram of an atom, it seems that the transfer of an electron between identical orbitals 2 R does not change its energy. In fact, when electrons move through different orbitals, the repulsion between them decreases, due to which the potential energy still decreases. Electrons that occupy individual orbitals are called unpaired. Further, when studying the nature of chemical bonds, we will see that the valence of atoms is determined by the number of unpaired electrons. Nitrogen has three unpaired electrons and is truly trivalent. It is enough to recall the formula of ammonia NH 3. Carbon, according to the diagram, is divalent. However, when a relatively small energy is absorbed, one electron is transferred from sublevel 25 to sublevel 2 rub. Carbon goes into an excited state with the electronic formula s 2 2s ( 2p s . In this state it has four unpaired electrons. A free atom can remain in an excited state for only a very short time. But, once inside the molecule, the atom receives additional electrons to fill the orbitals. After this, the possibility of transition to the ground state is excluded, and the carbon atom remains tetravalent. In fact, the energy expended on electron excitation is compensated by the energy of formation of additional chemical bonds.

Occupation of the 2p-orbis of gals with second electrons occurs in oxygen, fluorine and neon (diagrams VIII, IX, X). In this case, the number of remaining unpaired electrons and, accordingly, the valence of atoms decreases successively. This corresponds to basic knowledge about the properties of oxygen, fluorine and neon: oxygen is divalent, fluorine is monovalent, and neon does not form chemical bonds, i.e. its valency is zero.

We saw that in elements from lithium to neon the second energy level is populated by electrons, and that is why they make up

• 2nd period of the periodic table. Next to neon, sodium begins to populate the third energy level, and then forms
• 3rd period as sublevels 35 and 3 are populated R. Energy diagrams and electronic formulas of elements from sodium to argon can be presented in abbreviated form by designating the repeating set of neon electrons as . The meaning of the abbreviated electronic formula is that it indicates only the valence electrons of the atom. The remaining electrons that make up electron core of an atom, for chemistry are of secondary importance. As an example, let's write abbreviated formulas and diagrams for sodium, silicon and argon (diagrams XI, XII and XIII).

The number of chemical elements in the 2nd and 3rd periods is determined by the total capacity of the 5- and /^-sublevels, which is eight electrons. Thus, the presence of exactly eight groups in the periodic table receives a physical explanation. The reason for the observed similarity of chemical elements in the groups also becomes clear. Comparing the energy diagrams of elements of the same group - lithium and sodium, carbon and silicon, etc. - we notice that they are characterized by the same population of the outer energy level. This implies, first of all, the same valency of the atoms, which determines the similarity of chemical properties. But the electronic structures of atoms, taken as a whole, are different. The number of electron shells increases from period to period, which entails an increase in the radii of the atoms. Therefore, as already noted, along with the similarity, there is also a certain directionality in the change in properties.

From electronic formulas and energy diagrams of atoms it is obvious that in groups IA and PA electrons fill the outer 5-sublevel, and in groups I HA-V111A - the outer p-sublevel. This provides a basis for classifying chemical elements into blocks. The first two groups are considered as block of s-elements, and groups from ША to VIIIA - as block of p-elements.

In the periodic table there are also groups with the same numbers, but with the addition of the symbol “B”. How is the existence of these groups explained? From Fig. 5.6 it is obvious that the sublevel 3d in energy it is between sublevels 45 and 4 R. In the periodic table, the 4th period, like the previous ones, begins with two 5-elements - potassium ([Ar]45 l) and calcium (fAr]4l 2). After calcium, non-sublevel settlement begins R, as in the 2nd and 3rd periods, and the sublevel 3d, whose capacity is 10 electrons. Electrons at the ^-sublevel appear one after another in scandium and the elements following it, including zinc. They are included in block of d-elements. The numbering of groups of d-elements is based on the fact that in groups III through VIII there is an equal number of electrons in the two upper sublevels of both p-elements (5- and p-sublevels) and d-elements (5- and d- sublevels). Groups IB and PV are numbered according to the population of the outer 5-sublevel, similar to 5-elements.

The fourth period is completed by p-elements following zinc. Their filled 3g/-nolevel is energetically stabilized and becomes lower than the sublevel As. This is explained by the different rate of decrease in the energy of the orbitals of the 45- and 3^/-sublevels as the charge of the atomic nucleus increases (Fig. 5.9).

Rice. 5.9.

Example 5.1. Write abbreviated electronic formulas for iron and krypton.

Solution. For both iron and krypton, the closest precursor noble gas is argon (Z = 18). Iron (Z = 26) has eight electrons left to fill the upper 45- and 36-sublevels. We write the formula 45 2 3rf 6. Krypton (Z = 36) adds 10 more electrons, which completely populate the sublevels 3d And Ar. Filled 3d- we put the sublevel in the formula up to the 45th sublevel: [Ar]3 10 45 2 4/? 6.

The fifth period of the periodic table is similar in structure to the fourth. Both of them contain 18 chemical elements. In the 5th period, rubidium and strontium belong to the 5-block elements, 10 elements from yttrium to cadmium belong to the d-block and the remaining six elements from indium to xenon belong to the R- block.

This is followed by the longest 6th and 7th periods, containing 32 elements. In the 6th period, a family of 14 chemical elements is added - from lanthanum to ytterbium, called lanthanides, and in the 7th - a similar family actinides - from actinium to nobelium. In their atoms, the 4/- and 5/-sublevels are filled with electrons, respectively. Lanthanides and actinides make up the block of /-elements. Due to the special characteristics of the orbitals of the /-sublevels, all lanthanides and all actinides exhibit great similarity in chemical properties.

Example 5.2. What explains that families of /-elements contain 14 chemical elements?

Solution. In accordance with the formula 2/+1 sublevel f(1=3) consists of seven orbitals. Therefore, its capacity is 14 electrons, and the gradual filling of the /-sublevel occurs in 14 chemical elements.

Thus, a brief review of the electronic structure of atoms in general terms reveals the physical basis for the periodicity of changes in the properties of chemical elements and, consequently, the periodic law of D. I. Mendeleev. Briefly, we can say that the periodic law is a consequence of the Pauli principle and the principle of least energy.

An atom is the smallest particle of matter, consisting of a nucleus and electrons. The structure of the electronic shells of atoms is determined by the position of the element in the Periodic Table of Chemical Elements by D.I. Mendeleev.

Electron and electron shell of an atom

An atom, which is generally neutral, consists of a positively charged nucleus and a negatively charged electron shell (electron cloud), with the total positive and negative charges being equal in absolute value. When calculating the relative atomic mass, the mass of electrons is not taken into account, since it is negligible and 1840 times less than the mass of a proton or neutron.

Rice. 1. Atom.

An electron is a completely unique particle that has a dual nature: it has both the properties of a wave and a particle. They continuously move around the core.

The space around the nucleus where the probability of finding an electron is most likely is called an electron orbital, or electron cloud. This space has a specific shape, which is designated by the letters s-, p-, d-, and f-. The S-electron orbital has a spherical shape, the p-orbital has the shape of a dumbbell or a three-dimensional figure eight, the shapes of the d- and f-orbitals are much more complex.

Rice. 2. Shapes of electron orbitals.

Around the nucleus, electrons are arranged in electron layers. Each layer is characterized by its distance from the nucleus and its energy, which is why electronic layers are often called electronic energy levels. The closer the level is to the nucleus, the lower the energy of the electrons in it. One element differs from another in the number of protons in the nucleus of the atom and, accordingly, in the number of electrons. Consequently, the number of electrons in the electron shell of a neutral atom is equal to the number of protons contained in the nucleus of this atom. Each subsequent element has one more proton in its nucleus, and one more electron in its electron shell.

The newly entering electron occupies the orbital with the lowest energy. However, the maximum number of electrons per level is determined by the formula:

where N is the maximum number of electrons, and n is the number of the energy level.

The first level can only have 2 electrons, the second can have 8 electrons, the third can have 18 electrons, and the fourth level can have 32 electrons. The outer level of an atom cannot contain more than 8 electrons: as soon as the number of electrons reaches 8, the next level, further from the nucleus, begins to be filled.

Structure of electronic shells of atoms

Each element stands in a certain period. A period is a horizontal collection of elements arranged in order of increasing charge of the nuclei of their atoms, which begins with an alkali metal and ends with an inert gas. The first three periods in the table are small, and the next, starting from the fourth period, are large, consisting of two rows. The number of the period in which the element is located has a physical meaning. It means how many electronic energy levels there are in an atom of any element of a given period. Thus, the element chlorine Cl is in the 3rd period, that is, its electron shell has three electronic layers. Chlorine is in group VII of the table, and in the main subgroup. The main subgroup is the column within each group that begins with period 1 or 2.

Thus, the state of the electron shells of the chlorine atom is as follows: the atomic number of the chlorine element is 17, which means that the atom has 17 protons in the nucleus and 17 electrons in the electron shell. At level 1 there can only be 2 electrons, at level 3 - 7 electrons, since chlorine is in the main subgroup of group VII. Then at level 2 there are: 17-2-7 = 8 electrons.

Topics of the Unified State Examination codifier: The structure of the electronic shells of atoms of elements of the first four periods: s-, p- and d-elements. Electronic configuration of atoms and ions. Ground and excited states of atoms.

One of the first models of the structure of the atom - “ pudding model " - developed D.D. Thomson in 1904. Thomson discovered the existence of electrons, for which he received the Nobel Prize. However, science at that time could not explain the existence of these same electrons in space. Thomson proposed that the atom consists of negative electrons placed in a uniformly positively charged “soup” that compensates for the charge of the electrons (another analogy is raisins in pudding). The model is, of course, original, but incorrect. But Thomson’s model was an excellent start for further work in this area.

And further work turned out to be effective. Thomson's student, Ernest Rutherford, based on experiments on the scattering of alpha particles on gold foil, proposed a new, planetary model of the structure of the atom.

According to Rutherford's model, an atom consists of a massive, positively charged nucleus and particles with a small mass - electrons, which, like planets around the Sun, fly around the nucleus and do not fall on it.

Rutherford's model turned out to be the next step in studying the structure of the atom. However, modern science uses a more advanced model proposed by Niels Bohr in 1913. We will dwell on it in more detail.

Atom is the smallest, electrically neutral, chemically indivisible particle of matter, consisting of a positively charged nucleus and a negatively charged electron shell.

In this case, the electrons do not move in a certain orbit, as Rutherford assumed, but rather chaotically. The collection of electrons that move around the nucleus is called electron shell .

A languid core, as Rutherford proved, is massive and positively charged, located in the central part of the atom. The structure of the nucleus is quite complex and is studied in nuclear physics. The main particles that it consists of are: protons And neutrons. They are connected by nuclear forces ( strong interaction).

Let's look at the main characteristics protons, neutrons And electrons:

	Proton	Neutron	Electron
Weight	1.00728 amu	1.00867 amu	1/1960 amu
Charge	+ 1 elemental charge	0	- 1 elemental charge

1 amu (atomic mass unit) = 1.66054 10 -27 kg

1 elementary charge = 1.60219 10 -19 C

And the most important thing. The periodic table of chemical elements, structured by Dmitry Ivanovich Mendeleev, obeys a simple and understandable logic: atomic number is the number of protons in the nucleus of that atom . Moreover, Dmitry Ivanovich did not hear about any protons in the 19th century. All the more brilliant is his discovery and ability, and scientific instinct, which allowed him to step forward a century and a half in science.

Hence, nuclear charge Z equals number of protons, i.e. atom numberin the Periodic Table of Chemical Elements.

An atom is a charged particle, therefore the number of protons is equal to the number of electrons: N e = N p = Z.

Atomic mass ( mass number A ) is equal to the total mass of large particles that make up the atom - protons and neutrons. Since the mass of a proton and netron is approximately equal to 1 atomic mass unit, the formula can be used: M = N p + N n

Mass number listed in the Periodic Table of Chemical Elements in the cell of each element.

Note! When solving USE problems, the mass number of all atoms, except chlorine, is rounded to the nearest whole number according to the rules of mathematics. The mass number of the chlorine atom in the Unified State Examination is considered to be 35.5.

Collected in the Periodic Table chemical elements - atoms with the same nuclear charge. However, can the number of other particles in these atoms change? Quite. For example, atoms with different numbers of neutrons are called isotopes of this chemical element. The same element can have several isotopes.

Try to answer the questions. The answers to them are at the end of the article:

1. Do isotopes of the same element have the same or different mass numbers?
2. Do isotopes of the same element have the same or different numbers of protons?

The chemical properties of atoms are determined by the structure of the electron shell and the charge of the nucleus. Thus, the chemical properties of isotopes of one element are practically the same.

Since atoms of the same element can exist in the form of different isotopes, the name often indicates the mass number, for example, chlorine-35, and the following form of notation of atoms is accepted:

A few more questions:

3. Determine the number of neutrons, protons and electrons in the bromine-81 isotope.

4. Determine the number of neutrons in the isotope chlorine-37.

Structure of the electron shell

According to the quantum model of the structure of the atom by Niels Bohr, electrons in an atom can only move along certain (stationary ) orbits, removed from the nucleus at a certain distance and characterized by a certain energy. Another name for stationary orbits is electronic layersor energy levels .

Electronic levels can be designated by numbers - 1, 2, 3, ..., n. The layer number increases as it moves away from the core. The level number corresponds to the principal quantum number n.

In one layer, electrons can move along different trajectories. The orbital trajectory is characterized by electronic sublayer . The type of sublevel characterizes orbital quantum number l = 0,1, 2, 3..., or the corresponding letters - s, p, d, g and etc.

Within one sublevel (electron orbitals of the same type), options for the arrangement of orbitals in space are possible. The more complex the geometry of the orbitals of a given sublevel, the more options for their location in space. Total number of orbitals sublevel of this type l can be determined by the formula: 2 l +1. Each orbital can contain no more than two electrons.

Orbital type	s	p	d	f	g
Orbital quantum number value l	0	1	2	3	4
Number of atomic orbitals of a given type 2 l+1	1	3	5	7	9
Maximum number of electrons in orbitals of a given type	2	6	10	14	18

We get a summary table:

Level number, n	Sublevel	Number	Maximum number of electrons
1	1s	1	2
2	2s	1	2
	2p	3	6
	3s	1	2
	3p	3	6
	3d	5	10
	4s	1	2
	4p	3	6
	4d	5	10
	4f	7

The filling of energy orbitals with electrons occurs according to some basic rules. Let's look at them in detail.

Pauli's principle (Pauli's exclusion): can be located in the same atomic orbital no more than two electrons with opposite spins (spin is a quantum mechanical characteristic of electron motion).

RuleHunda. In atomic orbitals with the same energy, electrons are located one at a time with parallel spins. Those. orbitals of one sublevel are filled as follows: First, one electron is distributed to each orbital. Only when one electron is distributed in all orbitals of a given sublevel do we occupy the orbitals with second electrons with opposite spins.

Thus, the sum of the spin quantum numbers of such electrons on one energy sublevel (shell) will be maximum.

For example, the filling of the 2p orbital with three electrons will occur like this: and not like this:

The principle of minimum energy. Electrons fill the lowest energy orbitals first. The energy of an atomic orbital is equivalent to the sum of the principal and orbital quantum numbers: n + l . If the sum is the same, then the orbital with the smaller principal quantum number is filled first. n .

JSC	1s	2s	2p	3s	3p	3d	4s	4p	4d	4f	5s	5p	5d	5f	5 g
n	1	2	2	3	3	3	4	4	4	4	5	5	5	5	5
l	0	0	1	0	1	2	0	1	2	3	0	1	2	3	4
n + l	1	2	3	3	4	5	4	5	6	7	5	6	7	8	9

Thus, energy series of orbitals looks like that:

1 s < 2 s < 2 p < 3 s < 3 p < 4 s < 3 d < 4 p < 5 s < 4 d < 5 p < 6 s < 4 f~ 5 d < 6 p < 7 s <5 f~ 6 d…

The electronic structure of an atom can be represented in different forms − energy diagram, electronic formula etc. Let's look at the main ones.

Energy diagram of an atom is a schematic representation of orbitals taking into account their energy. The diagram shows the arrangement of electrons in energy levels and sublevels. The filling of orbitals occurs according to quantum principles.

For example, energy diagram for a carbon atom:

Electronic formula is a record of the distribution of electrons among the orbitals of an atom or ion. First, the level number is indicated, then the orbital type. The superscript to the right of the letter shows the number of electrons in the orbital. Orbitals are listed in order of occupancy. Record 1s 2 means that there are 2 electrons in the 1st level of the s-sublevel.

For example, the electronic formula of carbon looks like this: 1s 2 2s 2 2p 2 .

For brevity of notation, instead of energy orbitals being completely filled with electrons, sometimes use the symbol for the nearest noble gas (group VIIIA element) having the appropriate electronic configuration.

For example, electronic formula nitrogen can be written like this: 1s 2 2s 2 2p 3 or like this: 2s 2 2p 3.

1s 2 =

1s 2 2s 2 2p 6 =

1s 2 2s 2 2p 6 3s 2 3p 6 = and so on.

Electronic formulas of elements of the first four periods

Let us consider filling the shell of elements of the first four periods with electrons. U hydrogen the very first energy level, the s-sublevel, is filled, with 1 electron located:

+1H 1s 1 1s

U helium The 1s orbital is completely filled:

+2He 1s 2 1s

Since the first energy level can hold a maximum of 2 electrons, lithium filling of the second energy level begins, starting from the orbital with the minimum energy - 2s. In this case, the first energy level is filled first:

+3Li 1s 2 2s 1 1s 2s

U beryllium 2s-sublevel is filled:

+4Be 1s 2 2s 2 1s 2s

+5B 1s 2 2s 2 2p 1 1s 2s 2p

At the next element, carbon, the next electron, according to Hund’s rule, fills a vacant orbital, and does not move into a partially occupied one:

+6C 1s 2 2s 2 2p 2 1s 2s 2p

Try to create electronic and electronic graphic formulas for the following elements, and then you can test yourself using the answers at the end of the article:

5. Nitrogen

6. Oxygen

7. Fluorine

U not shefilling of the second energy level is completed:

+10Ne 1s 2 2s 2 2p 6 1s 2s 2p

U sodium filling of the third energy level begins:

+11Na 1s 2 2s 2 2p 6 3s 1 1s 2s 2p 3s

From sodium to argon, the filling of the 3rd level occurs in the same order as the filling of the 2nd energy level. I propose to compile electronic formulas of elements from magnesium before argon check the answers yourself.

8. Magnesium

9. Aluminum

10. Silicon

11. Phosphorus

12. Sulfur

13. Chlorine

14. Argon

But starting from the 19th element, potassium, sometimes confusion begins - it fills not a 3d orbital, but a 4s. We mentioned earlier in this article that the filling of energy levels and sublevels with electrons occurs according to energy series of orbitals , and not in order. I recommend repeating it again. So the formula potassium:

+19K 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 1s 2s 2p3s 3p4s

To write further electronic formulas in the article we will use the abbreviated form:

+19K4s 1 4s

U calcium 4s sublevel is filled:

+20Ca4s 2 4s

Element has 21, scandia, according to the energy series of orbitals, filling begins 3d-sublevel:

+21Sc 3d 14s 2 4s 3d

Further filling 3d-sublevel occurs according to quantum rules, from titanium before vanadium :

+22Ti 3d 24s 2 4s 3d

+23V 3d 34s 2 4s 3d

However, for the next element the order of filling the orbitals is violated. Electronic configuration chromium like this:

+24Cr 3d 54s 1 4s 3d

What's the matter? But the fact is that with the “traditional” order of filling the orbitals (accordingly, incorrect in this case - 3d 4 4s 2) exactly one cell in d-sublevel would remain unfilled. It turned out that such filling is energetically less profitable. A more profitable, When d-the orbital is filled completely, at least with single electrons. This extra electron goes from 4s-sublevel. And small energy costs for jumping an electron from 4s-sublevel more than covers the energy effect from filling all 3d- orbiters. This effect is called - failure or electron slip. And it is observed when d-the orbital is underfilled by 1 electron (one electron per cell or two).

In the following elements, the “traditional” order of filling the orbitals returns again. Configuration manganese :

+25Mn 3d 54s 2

Similarly for cobalt And nickel. But at copper we are watching again electron failure - the electron again jumps from 4s-sublevel on 3d- sublevel:

+29Cu 3d 104s 1

On zinc, filling of the 3d sublevel is completed:

+30Zn 3d 104s 2

The following elements, from Gaul before krypton, the 4p sublevel is filled according to quantum rules. For example, the electronic formula Gaul :

+31Ga 3d 104s 2 4p 1

We will not give formulas for the remaining elements; you can compose them yourself and check yourself on the Internet.

Some important concepts:

External energy level is the energy level in an atom with maximum number that has electrons. For example, y copper (3d 104s 1) external energy level is the fourth.

Valence electrons - electrons in an atom that can participate in the formation of a chemical bond. For example, in chromium ( +24Cr 3d 54s 1) not only the electrons of the outer energy level are valence ( 4s 1), but also unpaired electrons on 3d-sublevel, because they can form chemical bonds.

Ground and excited states of the atom

The electronic formulas that we compiled before correspond to the ground energy state of the atom . This is the most energetically favorable state of the atom.

However, in order to form, an atom in most situations must have unpaired (single) electrons . And chemical bonds are energetically very beneficial for the atom. Consequently, the more unpaired electrons in an atom, the more bonds it can form, and, as a result, it will move into a more favorable energy state.

Therefore, if there is free energy orbitals at this level paired pairs of electrons can steam , and one of the electrons of the paired pair can move to the vacant orbital. Thus the number of unpaired electrons increases, and the atom can form more chemical bonds, which is very beneficial from an energy point of view. This state of the atom is called excited and are indicated by an asterisk.

For example, in the ground state boron has the following energy level configuration:

+5B 1s 2 2s 2 2p 1 1s 2s 2p

At the second level (external) there is one paired electron pair, one single electron and a pair of free (vacant) orbitals. Consequently, there is a possibility for an electron to move from a pair to a vacant orbital, we get excited state boron atom (indicated by an asterisk):

+5B* 1s 2 2s 1 2p 2 1s 2s 2p

Try to create an electronic formula yourself that corresponds to the excited state of atoms. Don't forget to check your answers!

15. Carbon

16. Beryllium

17. Oxygen

Electronic formulas of ions

Atoms can give and receive electrons. By donating or accepting electrons, they become ions .

Ions are charged particles. Excess charge is indicated index in the upper right corner.

If an atom gives away electrons, then the total charge of the resulting particle will be positive (remember that the number of protons in an atom is equal to the number of electrons, and when electrons are lost, the number of protons will be greater than the number of electrons). Positively charged ions are cations . For example: sodium cation is formed as follows:

+11Na 1s 2 2s 2 2p 6 3s 1 -1е = +11Na+1s 2 2s 2 2p 6 3s 0

If an atom accepts electrons, then acquires negative charge . Negatively charged particles are anions . For example, the chlorine anion will be formed as follows:

+17Cl 1s 2 2s 2 2p 6 3s 2 3p 5 +1e = +17Cl – 1s 2 2s 2 2p 6 3s 2 3p 6

Thus, the electronic formulas of ions can be obtained adding or removing electrons from an atom. note , when cations are formed, electrons leave with external energy level . When anions are formed, electrons come to outer energy level .