As you know, everything material in the Universe consists of atoms. An atom is the smallest unit of matter that carries its properties. In turn, the structure of an atom is made up of a magical trinity of microparticles: protons, neutrons and electrons.

Moreover, each of the microparticles is universal. That is, you cannot find two different protons, neutrons or electrons in the world. All of them are absolutely similar to each other. And the properties of the atom will depend only on the quantitative composition of these microparticles in the general structure of the atom.

For example, the structure of a hydrogen atom consists of one proton and one electron. Next in complexity, the helium atom is made up of two protons, two neutrons, and two electrons. A lithium atom is made up of three protons, four neutrons and three electrons, etc.

Structure of atoms (from left to right): hydrogen, helium, lithium

Atoms combine into molecules, and molecules combine into substances, minerals and organisms. The DNA molecule, which is the basis of all life, is a structure assembled from the same three magical building blocks of the universe as the stone lying on the road. Although this structure is much more complex.

Even more amazing facts are revealed when we try to take a closer look at the proportions and structure of the atomic system. It is known that an atom consists of a nucleus and electrons moving around it along a trajectory that describes a sphere. That is, it cannot even be called a movement in the usual sense of the word. The electron is rather located everywhere and immediately within this sphere, creating an electron cloud around the nucleus and forming an electromagnetic field.

Schematic representations of the structure of the atom

The nucleus of an atom consists of protons and neutrons, and almost the entire mass of the system is concentrated in it. But at the same time, the nucleus itself is so small that if you increase its radius to a scale of 1 cm, then the radius of the entire structure of the atom will reach hundreds of meters. Thus, everything that we perceive as dense matter consists of more than 99% of the energy bonds between physical particles alone and less than 1% of the physical forms themselves.

But what are these physical forms? What are they made of, and how material are they? To answer these questions, let's take a closer look at the structures of protons, neutrons, and electrons. So, we descend one more step into the depths of the microcosm - to the level of subatomic particles.

What is an electron made of?

The smallest particle of an atom is an electron. An electron has mass but no volume. In the scientific view, the electron does not consist of anything, but is a structureless point.

An electron cannot be seen under a microscope. It is observed only in the form of an electron cloud, which looks like a fuzzy sphere around the atomic nucleus. At the same time, it is impossible to say with accuracy where the electron is located at a moment in time. Devices are capable of capturing not the particle itself, but only its energy trace. The essence of the electron is not embedded in the concept of matter. It is rather like an empty form that exists only in and through movement.

No structure has yet been found in the electron. It is the same point particle as the quantum of energy. In fact, an electron is energy, however, this is its more stable form than the one represented by photons of light.

At the moment, the electron is considered indivisible. This is understandable, because it is impossible to divide something that has no volume. However, there are already developments in the theory, according to which the composition of an electron contains a trinity of such quasiparticles as:

• Orbiton - contains information about the orbital position of the electron;
• Spinon - responsible for the spin or torque;
• Holon - carries information about the charge of an electron.

However, as we see, quasi-particles have absolutely nothing in common with matter, and carry only information.

Photographs of atoms of different substances in an electron microscope

Interestingly, an electron can absorb energy quanta, such as light or heat. In this case, the atom moves to a new energy level, and the boundaries of the electron cloud expand. It also happens that the energy absorbed by an electron is so great that it can jump out of the atomic system and continue its movement as an independent particle. At the same time, it behaves like a photon of light, that is, it seems to cease to be a particle and begins to exhibit the properties of a wave. This has been proven in an experiment.

Young's experiment

In the course of the experiment, a stream of electrons was directed onto a screen with two slits cut into it. Passing through these slits, the electrons collided with the surface of another projection screen, leaving their mark on it. As a result of this “bombardment” by electrons, an interference pattern appeared on the projection screen, similar to that which would appear if waves, but not particles, passed through two slits.

Such a pattern occurs due to the fact that the wave, passing between the two slots, is divided into two waves. As a result of further movement, the waves overlap each other, and in some areas they cancel each other out. As a result, we get many stripes on the projection screen, instead of one, as it would be if the electron behaved like a particle.

The structure of the nucleus of an atom: protons and neutrons

Protons and neutrons make up the nucleus of an atom. And despite the fact that in the total volume the core occupies less than 1%, it is in this structure that almost the entire mass of the system is concentrated. But at the expense of the structure of protons and neutrons, physicists are divided in opinion, and at the moment there are two theories at once.

• Theory #1 - Standard

The Standard Model says that protons and neutrons are made up of three quarks connected by a cloud of gluons. Quarks are point particles, just like quanta and electrons. And gluons are virtual particles that ensure the interaction of quarks. However, neither quarks nor gluons have been found in nature, so this model is subject to severe criticism.

• Theory #2 - Alternative

But according to the alternative unified field theory developed by Einstein, the proton, like the neutron, like any other particle of the physical world, is an electromagnetic field rotating at the speed of light.

Electromagnetic fields of man and the planet

What are the principles of the structure of the atom?

Everything in the world - subtle and dense, liquid, solid and gaseous - is just the energy states of countless fields that permeate the space of the Universe. The higher the energy level in the field, the thinner and less perceptible it is. The lower the energy level, the more stable and tangible it is. In the structure of the atom, as well as in the structure of any other unit of the Universe, lies the interaction of such fields - different in energy density. It turns out that matter is only an illusion of the mind.

DEFINITION

Atom is the smallest chemical particle.

The variety of chemical compounds is due to the different combination of atoms of chemical elements into molecules and non-molecular substances. The ability of an atom to enter into chemical compounds, its chemical and physical properties are determined by the structure of the atom. In this regard, for chemistry, the internal structure of the atom and, first of all, the structure of its electron shell is of paramount importance.

Models of the structure of the atom

At the beginning of the 19th century, D. Dalton revived the atomistic theory, relying on the fundamental laws of chemistry known by that time (constancy of composition, multiple ratios and equivalents). The first experiments were carried out to study the structure of matter. However, despite the discoveries made (the atoms of the same element have the same properties, and the atoms of other elements have different properties, the concept of atomic mass was introduced), the atom was considered indivisible.

After receiving experimental evidence (late XIX - early XX century) of the complexity of the structure of the atom (photoelectric effect, cathode and X-rays, radioactivity), it was found that the atom consists of negatively and positively charged particles that interact with each other.

These discoveries gave impetus to the creation of the first models of the structure of the atom. One of the first models was proposed J. Thomson(1904) (Fig. 1): the atom was presented as a "sea of ​​positive electricity" with electrons oscillating in it.

After experiments with α-particles, in 1911. Rutherford proposed the so-called planetary model structure of the atom (Fig. 1), similar to the structure of the solar system. According to the planetary model, in the center of the atom there is a very small nucleus with a charge Z e, the size of which is approximately 1,000,000 times smaller than the size of the atom itself. The nucleus contains almost the entire mass of the atom and has a positive charge. Electrons move in orbits around the nucleus, the number of which is determined by the charge of the nucleus. The outer trajectory of the electrons determines the outer dimensions of the atom. The diameter of an atom is 10 -8 cm, while the diameter of the nucleus is much smaller -10 -12 cm.

Rice. 1 Models of the structure of the atom according to Thomson and Rutherford

Experiments on the study of atomic spectra showed the imperfection of the planetary model of the structure of the atom, since this model contradicts the line structure of atomic spectra. Based on the Rutherford model, Einstein's theory of light quanta and the quantum theory of radiation, Planck Niels Bohr (1913) formulated postulates, which contains atomic theory(Fig. 2): an electron can rotate around the nucleus not in any, but only in some specific orbits (stationary), moving along such an orbit, it does not emit electromagnetic energy, radiation (absorption or emission of a quantum of electromagnetic energy) occurs during the transition (jump-like) electron from one orbit to another.

Rice. 2. Model of the structure of the atom according to N. Bohr

The accumulated experimental material characterizing the structure of the atom showed that the properties of electrons, as well as other micro-objects, cannot be described on the basis of the concepts of classical mechanics. Microparticles obey the laws of quantum mechanics, which became the basis for creating modern model of the structure of the atom.

The main theses of quantum mechanics:

- energy is emitted and absorbed by bodies in separate portions - quanta, therefore, the energy of particles changes abruptly;

- electrons and other microparticles have a dual nature - it exhibits the properties of both particles and waves (particle-wave dualism);

— quantum mechanics denies the existence of certain orbits for microparticles (for moving electrons it is impossible to determine the exact position, because they move in space near the nucleus, one can only determine the probability of finding an electron in different parts of space).

The space near the nucleus, in which the probability of finding an electron is sufficiently high (90%), is called orbital.

quantum numbers. Pauli principle. Rules of Klechkovsky

The state of an electron in an atom can be described using four quantum numbers.

n is the principal quantum number. Characterizes the total energy of an electron in an atom and the number of the energy level. n takes on integer values ​​from 1 to ∞. The electron has the lowest energy at n=1; with increasing n - energy. The state of an atom, when its electrons are at such energy levels that their total energy is minimal, is called the ground state. States with higher values ​​are called excited. Energy levels are indicated by Arabic numerals according to the value of n. Electrons can be arranged in seven levels, therefore, in reality, n exists from 1 to 7. The main quantum number determines the size of the electron cloud and determines the average radius of the electron in the atom.

l is the orbital quantum number. It characterizes the energy reserve of electrons in the sublevel and the shape of the orbital (Table 1). Accepts integer values ​​from 0 to n-1. l depends on n. If n=1, then l=0, which means that at the 1st level there is a 1st sublevel.

me is the magnetic quantum number. Characterizes the orientation of the orbital in space. Accepts integer values ​​from –l through 0 to +l. Thus, when l=1 (p-orbital), m e takes on the values ​​-1, 0, 1, and the orientation of the orbital can be different (Fig. 3).

Rice. 3. One of the possible orientations in the p-orbital space

s is the spin quantum number. Characterizes the electron's own rotation around the axis. It takes the values ​​-1/2(↓) and +1/2 (). Two electrons in the same orbital have antiparallel spins.

The state of electrons in atoms is determined Pauli principle: an atom cannot have two electrons with the same set of all quantum numbers. The sequence of filling orbitals with electrons is determined by Klechkovsky's rules: orbitals are filled with electrons in ascending order of the sum (n + l) for these orbitals, if the sum (n + l) is the same, then the orbital with the lower value of n is filled first.

However, an atom usually contains not one, but several electrons, and in order to take into account their interaction with each other, the concept of the effective charge of the nucleus is used - an electron of the outer level is affected by a charge that is less than the charge of the nucleus, as a result of which the inner electrons screen the outer ones.

The main characteristics of an atom: atomic radius (covalent, metallic, van der Waals, ionic), electron affinity, ionization potential, magnetic moment.

Electronic formulas of atoms

All the electrons of an atom form its electron shell. The structure of the electron shell is depicted electronic formula, which shows the distribution of electrons over energy levels and sublevels. The number of electrons in a sublevel is indicated by a number, which is written to the upper right of the letter indicating the sublevel. For example, a hydrogen atom has one electron, which is located on the s-sublevel of the 1st energy level: 1s 1. The electronic formula of helium containing two electrons is written as follows: 1s 2.

For elements of the second period, electrons fill the 2nd energy level, which can contain no more than 8 electrons. First, the electrons fill the s-sublevel, then the p-sublevel. For example:

5 B 1s 2 2s 2 2p 1

The relationship of the electronic structure of the atom with the position of the element in the Periodic system

The electronic formula of an element is determined by its position in the Periodic system of D.I. Mendeleev. So, the number of the period corresponds to the elements of the second period, the electrons fill the 2nd energy level, which can contain no more than 8 electrons. First, the electrons fill In the elements of the second period, the electrons fill the 2nd energy level, which can contain no more than 8 electrons. First, the electrons fill the s-sublevel, then the p-sublevel. For example:

5 B 1s 2 2s 2 2p 1

For atoms of some elements, the phenomenon of "leakage" of an electron from an external energy level to the penultimate one is observed. Electron slip occurs in atoms of copper, chromium, palladium and some other elements. For example:

24 Cr 1s 2 2s 2 2p 6 3s 2 3p 6 3d 5 4s 1

energy level that can contain no more than 8 electrons. First, the electrons fill the s-sublevel, then the p-sublevel. For example:

5 B 1s 2 2s 2 2p 1

The group number for the elements of the main subgroups is equal to the number of electrons in the external energy level, such electrons are called valence electrons (they participate in the formation of a chemical bond). The valence electrons of the elements of the side subgroups can be electrons of the outer energy level and the d-sublevel of the penultimate level. The number of the group of elements of the side subgroups of III-VII groups, as well as for Fe, Ru, Os, corresponds to the total number of electrons in the s-sublevel of the outer energy level and the d-sublevel of the penultimate level

Tasks:

Draw the electronic formulas of phosphorus, rubidium and zirconium atoms. List the valence electrons.

Answer:

15 P 1s 2 2s 2 2p 6 3s 2 3p 3 Valence electrons 3s 2 3p 3

37 Rb 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 5s 1 Valence electrons 5s 1

40 Zr 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 4d 2 5s 2 Valence electrons 4d 2 5s 2

Lecture: The structure of the electron shells of atoms of the elements of the first four periods: s-, p- and d-elements

The structure of the atom

The 20th century is the time of the invention of the "model of the structure of the atom". Based on the provided structure, it was possible to develop the following hypothesis: around a nucleus that is sufficiently small in volume and size, electrons make movements similar to the movement of planets around the Sun. Subsequent study of the atom showed that the atom itself and its structure are much more complex than previously established. And at present, with enormous opportunities in the scientific field, the atom is not fully explored. Such components as an atom and molecules are considered to be objects of the microworld. Therefore, a person is not able to consider these parts on his own. In this world, completely different laws and rules are established, which differ from the macrocosm. Proceeding from this, the study of the atom is carried out on its model.

Any atom is assigned a serial number, fixed in the Periodic Table of Mendeleev D.I. For example, the serial number of the phosphorus atom (P) is 15.

So an atom is made up of protons (p + ) , neutrons (n 0 ) and electrons (e - ). Protons and neutrons form the nucleus of an atom, it has a positive charge. And the electrons moving around the nucleus "construct" the electron shell of the atom, which has a negative charge.

How many electrons are in an atom? It's easy to know. It is enough to look at the ordinal number of the element in the table.

So, the number of electrons in phosphorus is 15 . The number of electrons contained in the shell of an atom is strictly equal to the number of protons contained in the nucleus. So the protons in the nucleus of the phosphorus atom 15 .

The mass of protons and neutrons that make up the mass of the nucleus of an atom is the same. And electrons are 2000 times smaller. This means that the entire mass of the atom is concentrated in the nucleus, the mass of electrons is neglected. We can also find out the mass of the nucleus of an atom from the table. Look at the image of phosphorus in the table. Below we see the designation 30, 974 - this is the mass of the phosphorus nucleus, its atomic mass. When writing, we round this figure. Based on the foregoing, we write the structure of the phosphorus atom as follows:

(at the bottom left they wrote the charge of the nucleus - 15, at the top left the rounded value of the mass of the atom - 31).

The nucleus of a phosphorus atom:

(at the bottom left we write the charge: protons have a charge equal to +1, and neutrons are not charged, that is, charge 0; at the top left, the mass of a proton and a neutron, equal to 1, is a conventional unit of mass of an atom; the charge of an atom's nucleus is equal to the number of protons in the nucleus, which means p = 15, and the number of neutrons must be calculated: subtract the charge from the atomic mass, i.e. 31 - 15 = 16).

The electron shell of the phosphorus atom is 15 negatively charged electrons that balance positively charged protons. Therefore, an atom is an electrically neutral particle.

Energy levels

Fig.1

Next, we need to analyze in detail how electrons are distributed in an atom. Their movement is not chaotic, but is subject to a specific order. Some of the available electrons are attracted to the nucleus with a sufficiently large force, while others, on the contrary, are attracted weakly. The root cause of such behavior of electrons is hidden in varying degrees of remoteness of electrons from the nucleus. That is, an electron closer to the nucleus will become more strongly interconnected with it. These electrons simply cannot be detached from the electron shell. The farther the electron is from the nucleus, the easier it is to "pull" it out of the shell. Also, the energy of an electron increases as it moves away from the nucleus of an atom. The electron energy is determined by the main quantum number n, which is equal to any natural number (1,2,3,4…). Electrons having the same value of n form one electron layer, as if fencing off other electrons moving at a remote distance. Figure 1 shows the electron layers contained in the electron shell at the center of the atom's nucleus.

You can notice how the volume of the layer increases as you move away from the core. Therefore, the farther the layer is from the nucleus, the more electrons it contains.

The electron layer contains electrons that are similar in terms of energy. Because of this, such layers are often referred to as energy levels. How many levels can an atom contain? The number of energy levels is equal to the number of the period in the periodic table D.I. in which the element is located. For example, phosphorus (P) is in the third period, so the phosphorus atom has three energy levels.

Rice. 2

How to find out the maximum number of electrons located on one electron layer? For this we use the formula Nmax = 2n 2 , where n is the level number.

We get that the first level contains only 2 electrons, the second - 8, the third - 18, the fourth - 32.

Each energy level contains sublevels. Their letters are: s-, p-, d- and f-. Look at fig. 2:

Energy levels are marked with different colors, and sublevels with stripes of different thicknesses.

The thinnest sublevel is denoted by the letter s. 1s is the s-sublevel of the first level, 2s is the s-sublevel of the second level, and so on.

The p-sublevel appeared at the second energy level, the d-sublevel appeared at the third one, and the f-sublevel appeared at the fourth one.

Remember what you saw: the first energy level includes one s-sublevel, the second two s- and p-sublevels, the third three s-, p- and d-sublevels, and the fourth level four s-, p-, d- and f-sublevels.

On the Only 2 electrons can be in the s-sublevel, a maximum of 6 electrons in the p-sublevel, 10 electrons in the d-sublevel, and up to 14 electrons in the f-sublevel.

Electronic orbitals

The area (place) where an electron can be located is called an electron cloud or orbital. Keep in mind that we are talking about the probable region where the electron is located, since the speed of its movement is hundreds of thousands of times greater than the speed of the needle of a sewing machine. Graphically, this area is displayed as a cell:

One cell can contain two electrons. Judging by Figure 2, we can conclude that the s-sublevel, which includes no more than two electrons, can contain only one s-orbital, is denoted by one cell; The p-sublayer has three p-orbitals (3 slots), the d-sublayer has five d-orbitals (5 slots), and the f-sublayer has seven f-orbitals (7 slots).

The shape of the orbital depends on orbital quantum number (l - el) atom. Atomic energy level originates from s- an orbital that has l= 0. The presented orbital has a spherical shape. At the levels after s- orbitals are formed p- orbitals with l = 1. P Orbitals are shaped like dumbbells. There are only three orbitals with this shape. Each possible orbital contains no more than 2 electrons. Next are more complex structures. d-orbitals ( l= 2), and after them f-orbitals ( l = 3).

Rice. 3 The shape of the orbitals

Electrons in orbitals are shown as arrows. If the orbitals contain one electron each, then they are unidirectional - arrow up:

If there are two electrons in the orbital, then they have two directions: an arrow up and an arrow down, i.e. electrons are in opposite directions:

This structure of electrons is called valence.

There are three conditions for filling atomic orbitals with electrons:

1 condition: The principle of the minimum amount of energy. The filling of orbitals starts from the sublevel having the minimum energy. According to this principle, the sublevels are filled in the following order: take a place in a sub-level of a higher level, although the sub-level of a lower level is not filled. For example, the valence configuration of a phosphorus atom looks like this:

Rice. four

2 condition: Pauli principle. One orbital includes 2 electrons (electron pair) and no more. But the content of only one electron is also possible. It is called unpaired.

3 condition: Hund's rule. Each orbital of one sublevel is first filled with one electron, then a second electron is added to them. In life, we have seen a similar situation when unfamiliar bus passengers first occupy all the free seats one at a time, and then take two seats.

Electronic configuration of an atom in the ground and excited state

The energy of an atom in its ground state is the lowest. If atoms begin to receive energy from outside, for example, when a substance is heated, then they pass from the ground state to an excited one. This transition is possible in the presence of free orbitals to which electrons can move. But this is temporary, giving off energy, the excited atom returns to its ground state.

Let's consolidate our knowledge with an example. Consider the electronic configuration, i.e. the concentration of electrons in the orbitals of the phosphorus atom in the ground (unexcited state). Let us turn again to Fig. 4. So, remember that the phosphorus atom has three energy levels, which are represented by half-arcs: +15)))

Let's distribute the available 15 electrons into these three energy levels:

Such formulas are called electronic configurations. There are also electronic - graphic, they illustrate the placement of electrons inside the energy levels. The electronic-graphic configuration of phosphorus looks like this: 1s 2 2s 2 2p 6 3s 2 3p 3 (here the large numbers are the numbers of energy levels, the letters are the sublevels, and the small numbers are the number of electrons in the sublevel, if you add them up, you get the number 15).

In the excited state of the phosphorus atom 1, the electron moves from the 3s orbital to the 3d orbital, and the configuration looks like this: 1s 2 2s 2 2p 6 3s 1 3p 3 3d 1 .

Electrons

The concept of an atom originated in the ancient world to denote the particles of matter. In Greek, atom means "indivisible".

The Irish physicist Stoney, on the basis of experiments, came to the conclusion that electricity is carried by the smallest particles that exist in the atoms of all chemical elements. In 1891, Stoney proposed to call these particles electrons, which in Greek means "amber". A few years after the electron got its name, English physicist Joseph Thomson and French physicist Jean Perrin proved that electrons carry a negative charge. This is the smallest negative charge, which in chemistry is taken as a unit (-1). Thomson even managed to determine the speed of the electron (the speed of an electron in orbit is inversely proportional to the orbit number n. The radii of the orbits grow in proportion to the square of the orbit number. In the first orbit of the hydrogen atom (n=1; Z=1), the speed is ≈ 2.2 106 m / c, that is, about a hundred times less than the speed of light c=3 108 m/s.) and the mass of an electron (it is almost 2000 times less than the mass of a hydrogen atom).

The state of electrons in an atom

The state of an electron in an atom is a set of information about the energy of a particular electron and the space in which it is located. An electron in an atom does not have a trajectory of motion, i.e., one can only speak of the probability of finding it in the space around the nucleus.

It can be located in any part of this space surrounding the nucleus, and the totality of its various positions is considered as an electron cloud with a certain negative charge density. Figuratively, this can be imagined as follows: if it were possible to photograph the position of an electron in an atom in hundredths or millionths of a second, as in a photo finish, then the electron in such photographs would be represented as points. Overlaying countless such photographs would result in a picture of an electron cloud with the highest density where there will be most of these points.

The space around the atomic nucleus, in which the electron is most likely to be found, is called the orbital. It contains approximately 90% e-cloud, and this means that about 90% of the time the electron is in this part of space. Distinguished by shape 4 currently known types of orbitals, which are denoted by Latin letters s, p, d and f. A graphic representation of some forms of electronic orbitals is shown in the figure.

The most important characteristic of the motion of an electron in a certain orbit is the energy of its connection with the nucleus. Electrons with similar energy values ​​form a single electron layer, or energy level. Energy levels are numbered starting from the nucleus - 1, 2, 3, 4, 5, 6 and 7.

An integer n, denoting the number of the energy level, is called the main quantum number. It characterizes the energy of electrons occupying a given energy level. The electrons of the first energy level, closest to the nucleus, have the lowest energy. Compared with the electrons of the first level, the electrons of the next levels will be characterized by a large amount of energy. Consequently, the electrons of the outer level are the least strongly bound to the nucleus of the atom.

The largest number of electrons in the energy level is determined by the formula:

N = 2n2,

where N is the maximum number of electrons; n is the level number, or the main quantum number. Consequently, the first energy level closest to the nucleus can contain no more than two electrons; on the second - no more than 8; on the third - no more than 18; on the fourth - no more than 32.

Starting from the second energy level (n = 2), each of the levels is subdivided into sublevels (sublayers), which differ somewhat from each other in the binding energy with the nucleus. The number of sublevels is equal to the value of the main quantum number: the first energy level has one sublevel; the second - two; third - three; fourth - four sublevels. Sublevels, in turn, are formed by orbitals. Each valuen corresponds to the number of orbitals equal to n.

It is customary to designate sublevels in Latin letters, as well as the shape of the orbitals of which they consist: s, p, d, f.

Protons and neutrons

An atom of any chemical element is comparable to a tiny solar system. Therefore, such a model of the atom, proposed by E. Rutherford, is called planetary.

The atomic nucleus, in which the entire mass of the atom is concentrated, consists of particles of two types - protons and neutrons.

Protons have a charge equal to the charge of electrons, but opposite in sign (+1), and a mass equal to the mass of a hydrogen atom (it is accepted in chemistry as a unit). Neutrons carry no charge, they are neutral and have a mass equal to that of a proton.

Protons and neutrons are collectively called nucleons (from the Latin nucleus - nucleus). The sum of the number of protons and neutrons in an atom is called the mass number. For example, the mass number of an aluminum atom:

13 + 14 = 27

number of protons 13, number of neutrons 14, mass number 27

Since the mass of the electron, which is negligible, can be neglected, it is obvious that the entire mass of the atom is concentrated in the nucleus. Electrons represent e - .

Because the atom electrically neutral, it is also obvious that the number of protons and electrons in an atom is the same. It is equal to the serial number of the chemical element assigned to it in the Periodic system. The mass of an atom is made up of the mass of protons and neutrons. Knowing the serial number of the element (Z), i.e., the number of protons, and the mass number (A), equal to the sum of the numbers of protons and neutrons, you can find the number of neutrons (N) using the formula:

N=A-Z

For example, the number of neutrons in an iron atom is:

56 — 26 = 30

isotopes

Varieties of atoms of the same element that have the same nuclear charge but different mass numbers are called isotopes. Chemical elements found in nature are a mixture of isotopes. So, carbon has three isotopes with a mass of 12, 13, 14; oxygen - three isotopes with a mass of 16, 17, 18, etc. Usually given in the Periodic system, the relative atomic mass of a chemical element is the average value of the atomic masses of a natural mixture of isotopes of a given element, taking into account their relative abundance in nature. The chemical properties of the isotopes of most chemical elements are exactly the same. However, hydrogen isotopes differ greatly in properties due to the dramatic fold increase in their relative atomic mass; they have even been given individual names and chemical symbols.

Elements of the first period

Scheme of the electronic structure of the hydrogen atom:

Schemes of the electronic structure of atoms show the distribution of electrons over electronic layers (energy levels).

The graphical electronic formula of the hydrogen atom (shows the distribution of electrons over energy levels and sublevels):

Graphic electronic formulas of atoms show the distribution of electrons not only in levels and sublevels, but also in orbits.

In a helium atom, the first electron layer is completed - it has 2 electrons. Hydrogen and helium are s-elements; for these atoms, the s-orbital is filled with electrons.

All elements of the second period the first electron layer is filled, and the electrons fill the s- and p-orbitals of the second electron layer in accordance with the principle of least energy (first s, and then p) and the rules of Pauli and Hund.

In the neon atom, the second electron layer is completed - it has 8 electrons.

For atoms of elements of the third period, the first and second electron layers are completed, so the third electron layer is filled, in which electrons can occupy 3s-, 3p- and 3d-sublevels.

A 3s ​​electron orbital is completed at the magnesium atom. Na and Mg are s-elements.

For aluminum and subsequent elements, the 3p sublevel is filled with electrons.

The elements of the third period have unfilled 3d orbitals.

All elements from Al to Ar are p-elements. s- and p-elements form the main subgroups in the Periodic system.

Elements of the fourth - seventh periods

A fourth electron layer appears at the potassium and calcium atoms, the 4s sublevel is filled, since it has less energy than the 3d sublevel.

K, Ca - s-elements included in the main subgroups. For atoms from Sc to Zn, the 3d sublevel is filled with electrons. These are 3d elements. They are included in the secondary subgroups, they have a pre-external electron layer filled, they are referred to as transition elements.

Pay attention to the structure of the electron shells of chromium and copper atoms. In them, a “failure” of one electron from the 4s- to the 3d-sublevel occurs, which is explained by the greater energy stability of the resulting electronic configurations 3d 5 and 3d 10:

In the zinc atom, the third electron layer is completed - all the 3s, 3p and 3d sublevels are filled in it, in total there are 18 electrons on them. In the elements following zinc, the fourth electron layer continues to be filled, the 4p sublevel.

Elements from Ga to Kr are p-elements.

The outer layer (fourth) of the krypton atom is complete and has 8 electrons. But there can only be 32 electrons in the fourth electron layer; the 4d- and 4f-sublevels of the krypton atom still remain unfilled. The elements of the fifth period are filling the sub-levels in the following order: 5s - 4d - 5p. And there are also exceptions related to " failure» electrons, y 41 Nb, 42 Mo, 44 ​​Ru, 45 Rh, 46 Pd, 47 Ag.

In the sixth and seventh periods, f-elements appear, i.e., elements in which the 4f- and 5f-sublevels of the third outside electronic layer are filled, respectively.

4f elements are called lanthanides.

5f elements are called actinides.

The order of filling of electronic sublevels in the atoms of elements of the sixth period: 55 Cs and 56 Ba - 6s-elements; 57 La … 6s 2 5d x - 5d element; 58 Ce - 71 Lu - 4f elements; 72 Hf - 80 Hg - 5d elements; 81 T1 - 86 Rn - 6d elements. But even here there are elements in which the order of filling of electronic orbitals is “violated”, which, for example, is associated with greater energy stability of half and completely filled f-sublevels, i.e. nf 7 and nf 14. Depending on which sublevel of the atom is filled with electrons last, all elements are divided into four electronic families, or blocks:

• s-elements. The s-sublevel of the outer level of the atom is filled with electrons; s-elements include hydrogen, helium and elements of the main subgroups of groups I and II.

• p-elements. The p-sublevel of the outer level of the atom is filled with electrons; p-elements include elements of the main subgroups of III-VIII groups.

• d-elements. The d-sublevel of the preexternal level of the atom is filled with electrons; d-elements include elements of secondary subgroups of groups I-VIII, i.e., elements of intercalary decades of large periods located between s- and p-elements. They are also called transition elements.

• f-elements. The f-sublevel of the third outside level of the atom is filled with electrons; these include the lanthanides and antinoids.

The Swiss physicist W. Pauli in 1925 established that in an atom in one orbital there can be no more than two electrons having opposite (antiparallel) spins (translated from English - “spindle”), i.e. having such properties that can be conditionally imagined as the rotation of an electron around its imaginary axis: clockwise or counterclockwise.

This principle is called Pauli principle. If there is one electron in the orbital, then it is called unpaired, if there are two, then these are paired electrons, that is, electrons with opposite spins. The figure shows a diagram of the division of energy levels into sublevels and the order in which they are filled.

Very often, the structure of the electron shells of atoms is depicted using energy or quantum cells - they write down the so-called graphic electronic formulas. For this record, the following notation is used: each quantum cell is denoted by a cell that corresponds to one orbital; each electron is indicated by an arrow corresponding to the direction of the spin. When writing a graphical electronic formula, two rules should be remembered: Pauli principle and F. Hund's rule, according to which electrons occupy free cells first one at a time and at the same time have the same spin value, and only then they pair, but the spins, according to the Pauli principle, will already be oppositely directed.

Hund's rule and Pauli's principle

Hund's rule- the rule of quantum chemistry, which determines the order of filling the orbitals of a certain sublayer and is formulated as follows: the total value of the spin quantum number of electrons of this sublayer should be maximum. Formulated by Friedrich Hund in 1925.

This means that in each of the orbitals of the sublayer, one electron is first filled, and only after the unfilled orbitals are exhausted, a second electron is added to this orbital. In this case, there are two electrons with half-integer spins of the opposite sign in one orbital, which pair (form a two-electron cloud) and, as a result, the total spin of the orbital becomes equal to zero.

Other wording: Below in energy lies the atomic term for which two conditions are satisfied.

1. Multiplicity is maximum
2. When the multiplicities coincide, the total orbital momentum L is maximum.

Let's analyze this rule using the example of filling the orbitals of the p-sublevel p- elements of the second period (that is, from boron to neon (in the diagram below, horizontal lines indicate orbitals, vertical arrows indicate electrons, and the direction of the arrow indicates the orientation of the spin).

Klechkovsky's rule

Klechkovsky's rule - as the total number of electrons in atoms increases (with an increase in the charges of their nuclei, or the ordinal numbers of chemical elements), atomic orbitals are populated in such a way that the appearance of electrons in higher-energy orbitals depends only on the principal quantum number n and does not depend on all other quantum numbers. numbers, including those from l. Physically, this means that in a hydrogen-like atom (in the absence of interelectron repulsion) the orbital energy of an electron is determined only by the spatial remoteness of the electron charge density from the nucleus and does not depend on the features of its motion in the field of the nucleus.

Klechkovsky's empirical rule and the sequence of sequences of a somewhat contradictory real energy sequence of atomic orbitals arising from it only in two cases of the same type: for atoms Cr, Cu, Nb, Mo, Ru, Rh, Pd, Ag, Pt, Au, there is a “failure” of an electron with s - sublevel of the outer layer to the d-sublevel of the previous layer, which leads to an energetically more stable state of the atom, namely: after filling the orbital 6 with two electrons s