The lesson is devoted to the formation of ideas about the complex structure of the atom. The state of electrons in an atom is considered, the concepts of "atomic orbital and electron cloud", the forms of orbitals (s--, p-, d-orbitals) are introduced. Also considered are aspects such as the maximum number of electrons at energy levels and sublevels, the distribution of electrons over energy levels and sublevels in atoms of elements of the first four periods, valence electrons of s-, p- and d-elements. A graphical diagram of the structure of the electronic layers of atoms (electron-graphic formula) is given.

Topic: The structure of the atom. Periodic law D.I. Mendeleev

Lesson: The structure of the atom

Translated from Greek, the word " atom" means "indivisible". However, phenomena have been discovered that demonstrate the possibility of its division. These are the emission of x-rays, the emission of cathode rays, the phenomenon of the photoelectric effect, the phenomenon of radioactivity. Electrons, protons, and neutrons are the particles that make up an atom. They're called subatomic particles.

Tab. one

In addition to protons, the nucleus of most atoms contains neutrons that carry no charge. As can be seen from Table. 1, the mass of the neutron practically does not differ from the mass of the proton. Protons and neutrons make up the nucleus of an atom and are called nucleons (nucleus - nucleus). Their charges and masses in atomic mass units (a.m.u.) are shown in Table 1. When calculating the mass of an atom, the mass of an electron can be neglected.

Mass of an atom ( mass number) is equal to the sum of the masses of the protons and neutrons that make up its nucleus. The mass number is denoted by the letter BUT. From the name of this quantity, it can be seen that it is closely related to the atomic mass of the element rounded to an integer. A=Z+N

Here A- mass number of an atom (the sum of protons and neutrons), Z- nuclear charge (number of protons in the nucleus), N is the number of neutrons in the nucleus. According to the doctrine of isotopes, the concept of "chemical element" can be given the following definition:

chemical element A group of atoms with the same nuclear charge is called.

Some elements exist as multiple isotopes. "Isotopes" means "occupying the same place." Isotopes have the same number of protons, but differ in mass, i.e., the number of neutrons in the nucleus (number N). Because neutrons have little to no effect on the chemical properties of elements, all isotopes of the same element are chemically indistinguishable.

Isotopes are called varieties of atoms of the same chemical element with the same nuclear charge (that is, with the same number of protons), but with a different number of neutrons in the nucleus.

Isotopes differ from each other only in mass number. This is indicated either by a superscript in the right corner, or in a line: 12 C or C-12 . If an element contains several natural isotopes, then in the periodic table D.I. Mendeleev indicates its average atomic mass, taking into account the prevalence. For example, chlorine contains 2 natural isotopes 35 Cl and 37 Cl, the content of which is 75% and 25%, respectively. Thus, the atomic mass of chlorine will be equal to:

BUTr(Cl)=0,75 . 35+0,25 . 37=35,5

For artificially synthesized heavy atoms, one atomic mass value is given in square brackets. This is the atomic mass of the most stable isotope of that element.

Basic models of the structure of the atom

Historically, the Thomson model of the atom was the first in 1897.

Rice. 1. Model of the structure of the atom by J. Thomson

The English physicist J. J. Thomson suggested that atoms consist of a positively charged sphere in which electrons are interspersed (Fig. 1). This model is figuratively called "plum pudding", a bun with raisins (where "raisins" are electrons), or "watermelon" with "seeds" - electrons. However, this model was abandoned, since experimental data were obtained that contradicted it.

Rice. 2. Model of the structure of the atom by E. Rutherford

In 1910, the English physicist Ernst Rutherford, with his students Geiger and Marsden, conducted an experiment that gave amazing results that were inexplicable from the point of view of the Thomson model. Ernst Rutherford proved by experience that in the center of the atom there is a positively charged nucleus (Fig. 2), around which, like planets around the Sun, electrons revolve. The atom as a whole is electrically neutral, and the electrons are held in the atom due to the forces of electrostatic attraction (Coulomb forces). This model had many contradictions and, most importantly, did not explain why electrons do not fall on the nucleus, as well as the possibility of absorption and emission of energy by it.

The Danish physicist N. Bohr in 1913, taking Rutherford's model of the atom as a basis, proposed a model of the atom in which electron-particles revolve around the atomic nucleus in much the same way as the planets revolve around the Sun.

Rice. 3. Planetary model of N. Bohr

Bohr suggested that electrons in an atom can only exist stably in orbits at strictly defined distances from the nucleus. These orbits he called stationary. An electron cannot exist outside stationary orbits. Why this is so, Bohr could not explain at the time. But he showed that such a model (Fig. 3) makes it possible to explain many experimental facts.

Currently used to describe the structure of the atom quantum mechanics. This is a science, the main aspect of which is that the electron has the properties of a particle and a wave at the same time, i.e., wave-particle duality. According to quantum mechanics, the region of space in which the probability of finding an electron is greatest is calledorbital. The farther the electron is from the nucleus, the lower its interaction energy with the nucleus. Electrons with close energies form energy level. Number of energy levels equals period number, in which this element is located in the table D.I. Mendeleev. There are various shapes of atomic orbitals. (Fig. 4). The d-orbital and f-orbital have a more complex shape.

Rice. 4. Shapes of atomic orbitals

There are exactly as many electrons in the electron shell of any atom as there are protons in its nucleus, so the atom as a whole is electrically neutral. Electrons in an atom are arranged so that their energy is minimal. The farther the electron is from the nucleus, the more orbitals and the more complex they are in shape. Each level and sublevel can only hold a certain number of electrons. The sublevels, in turn, consist of orbitals.

At the first energy level, closest to the nucleus, there can be one spherical orbital ( 1 s). At the second energy level - a spherical orbital, large in size and three p-orbitals: 2 s2 ppp. On the third level: 3 s3 ppp3 dddd.

In addition to movement around the nucleus, electrons also have movement, which can be represented as their movement around their own axis. This rotation is called spin ( in lane from English. "spindle"). Only two electrons with opposite (antiparallel) spins can be in one orbital.

Maximum number of electrons per energy level is determined by the formula N=2 n 2.

Where n is the main quantum number (energy level number). See table. 2

Tab. 2

Depending on which orbital the last electron is in, they distinguish s-, p-, d-elements. Elements of the main subgroups belong to s-, p-elements. In the side subgroups are d-elements

Graphic diagram of the structure of the electronic layers of atoms (electronic graphic formula).

To describe the arrangement of electrons in atomic orbitals, the electronic configuration is used. To write it in a line, orbitals are written in the legend ( s--, p-, d-,f-orbitals), and in front of them are numbers indicating the number of the energy level. The larger the number, the further the electron is from the nucleus. In upper case, above the designation of the orbital, the number of electrons in this orbital is written (Fig. 5).

Rice. 5

Graphically, the distribution of electrons in atomic orbitals can be represented as cells. Each cell corresponds to one orbital. There will be three such cells for the p-orbital, five for the d-orbital, and seven for the f-orbital. One cell can contain 1 or 2 electrons. According to Gund's rule, electrons are distributed in orbitals of the same energy (for example, in three p-orbitals), first one at a time, and only when there is already one electron in each such orbital, the filling of these orbitals with second electrons begins. Such electrons are called paired. This is explained by the fact that in neighboring cells, electrons repel each other less, as similarly charged particles.

See fig. 6 for atom 7 N.

Rice. 6

The electronic configuration of the scandium atom

21 sc: 1 s 2 2 s 2 2 p 6 3 s 2 3 p 6 4 s 2 3 d 1

Electrons in the outer energy level are called valence electrons. 21 sc refers to d-elements.

Summing up the lesson

At the lesson, the structure of the atom, the state of electrons in the atom were considered, the concept of "atomic orbital and electron cloud" was introduced. Students learned what the shape of orbitals is ( s-, p-, d-orbitals), what is the maximum number of electrons at energy levels and sublevels, the distribution of electrons over energy levels, what is s-, p- and d-elements. A graphical diagram of the structure of the electronic layers of atoms (electron-graphic formula) is given.

Bibliography

1. Rudzitis G.E. Chemistry. Fundamentals of General Chemistry. Grade 11: textbook for educational institutions: basic level / G.E. Rudzitis, F.G. Feldman. - 14th ed. - M.: Education, 2012.

2. Popel P.P. Chemistry: 8th grade: a textbook for general educational institutions / P.P. Popel, L.S. Krivlya. - K .: Information Center "Academy", 2008. - 240 p.: ill.

3. A.V. Manuilov, V.I. Rodionov. Fundamentals of chemistry. Internet tutorial.

Homework

1. No. 5-7 (p. 22) Rudzitis G.E. Chemistry. Fundamentals of General Chemistry. Grade 11: textbook for educational institutions: basic level / G.E. Rudzitis, F.G. Feldman. - 14th ed. - M.: Education, 2012.

2. Write electronic formulas for the following elements: 6 C, 12 Mg, 16 S, 21 Sc.

3. Elements have the following electronic formulas: a) 1s 2 2s 2 2p 4 .b) 1s 2 2s 2 2p 6 3s 2 3p 1. c) 1s 2 2s 2 2p 6 3s 2 3p 6 3d 6 4s 2 . What are these elements?

The composition of a molecule. That is, by what atoms the molecule is formed, in what quantity, by what bonds these atoms are connected. All this determines the property of the molecule, and, accordingly, the property of the substance that these molecules form.

For example, the properties of water: transparency, fluidity, the ability to cause rust are due precisely to the presence of two hydrogen atoms and one oxygen atom.

Therefore, before proceeding to the study of the properties of molecules (that is, the properties of substances), it is necessary to consider the “building blocks” by which these molecules are formed. Understand the structure of the atom.

How is an atom arranged?

Atoms are particles that, when combined with each other, form molecules.

The atom itself is made up of positively charged nucleus (+) and negatively charged electron shell (-). In general, the atom is electrically neutral. That is, the charge of the nucleus is equal in absolute value to the charge of the electron shell.

The nucleus is formed by the following particles:

• Protons. One proton carries a +1 charge. Its mass is 1 amu (atomic mass unit). These particles are necessarily present in the nucleus.

• Neutrons. The neutron has no charge (charge = 0). Its mass is 1 amu. Neutrons may not be in the nucleus. It is not a required component of the atomic nucleus.

Thus, protons are responsible for the total charge of the nucleus. Since one neutron has a charge of +1, the charge of the nucleus is equal to the number of protons.

The electron shell, as the name implies, is formed by particles called electrons. If we compare the nucleus of an atom with a planet, then electrons are its satellites. Revolving around the nucleus (for now let's imagine that in orbits, but in fact in orbits), they form an electron shell.

• Electron is a very small particle. Its mass is so small that it is taken as 0. But the charge of an electron is -1. That is, the modulus is equal to the charge of the proton, differs in sign. Since one electron carries a charge of -1, the total charge of the electron shell is equal to the number of electrons in it.

One important consequence, since an atom is a particle that does not have a charge (the charge of the nucleus and the charge of the electron shell are equal in absolute value, but opposite in sign), that is, electrically neutral, therefore, the number of electrons in an atom is equal to the number of protons.

How do atoms of different chemical elements differ from each other?

Atoms of different chemical elements differ from each other in the charge of the nucleus (that is, the number of protons, and, consequently, the number of electrons).

How to find out the charge of the nucleus of an atom of an element? The brilliant domestic chemist D. I. Mendeleev, having discovered the periodic law, and having developed a table named after him, gave us the opportunity to do this. His discovery was far ahead of the curve. When it was not yet known about the structure of the atom, Mendeleev arranged the elements in the table in order of increasing nuclear charge.

That is, the serial number of an element in the periodic system is the charge of the nucleus of an atom of a given element. For example, oxygen has a serial number of 8, respectively, the charge of the nucleus of the oxygen atom is +8. Accordingly, the number of protons is 8, and the number of electrons is 8.

It is the electrons in the electron shell that determine the chemical properties of the atom, but more on that later.

Now let's talk about the mass.

One proton is one unit of mass, one neutron is also one unit of mass. Therefore, the sum of neutrons and protons in the nucleus is called mass number. (The electrons do not affect the mass in any way, since we neglect its mass and consider it equal to zero).

The atomic mass unit (a.m.u.) is a special physical quantity for designating small masses of particles that form atoms.

All these three atoms are atoms of one chemical element - hydrogen. Because they have the same nuclear charge.

How will they differ? These atoms have different mass numbers (due to the different number of neutrons). The first atom has a mass number of 1, the second has 2, and the third has 3.

Atoms of the same element that differ in the number of neutrons (and hence mass numbers) are called isotopes.

The presented hydrogen isotopes even have their own names:

• The first isotope (mass number 1) is called protium.
• The second isotope (mass number 2) is called deuterium.
• The third isotope (with a mass number of 3) is called tritium.

Now the next reasonable question is why if the number of neutrons and protons in the nucleus is an integer, their mass is 1 amu, then in the periodic system the mass of an atom is a fractional number. For sulfur, for example: 32.066.

Answer: an element has several isotopes, they differ from each other in mass numbers. Therefore, the atomic mass in the periodic table is the average value of the atomic masses of all isotopes of an element, taking into account their occurrence in nature. This mass, given in the periodic system, is called relative atomic mass.

For chemical calculations, indicators of just such an “averaged atom” are used. Atomic mass is rounded to the nearest integer.

The structure of the electron shell.

The chemical properties of an atom are determined by the structure of its electron shell. The electrons around the nucleus are not arranged anyhow. Electrons are localized in electron orbitals.

Electronic orbital- the space around the atomic nucleus, where the probability of finding an electron is greatest.

An electron has one quantum parameter called spin. If we take the classical definition from quantum mechanics, then spin is the intrinsic angular momentum of the particle. In a simplified form, this can be represented as the direction of rotation of a particle around its axis.

An electron is a particle with a half-integer spin, an electron can have either +½ or -½ spin. Conventionally, this can be represented as a clockwise and counterclockwise rotation.

No more than two electrons with opposite spins can be in one electron orbital.

The generally accepted designation of an electronic dwelling is a cell or a dash. The electron is indicated by an arrow: the up arrow is an electron with a positive spin +½, the down arrow ↓ is an electron with a negative spin -½.

An electron that is alone in an orbital is called unpaired. Two electrons in the same orbital are called paired.

Electronic orbitals are divided into four types depending on the shape: s, p, d, f. Orbitals of the same shape form a sublevel. The number of orbitals at a sublevel is determined by the number of possible locations in space.

1. s orbital.

The s orbital is spherical:

In space, the s-orbital can only be located in one way:

Therefore, the s-sublevel is formed by only one s-orbital.

1. p-orbital.

The p orbital is shaped like a dumbbell:

In space, the p-orbital can only be located in three ways:

Therefore, the p-sublevel is formed by three p-orbitals.

1. d-orbital.

The d-orbital has a complex shape:

In space, the d-orbital can be located in five different ways. Therefore, the d-sublevel is formed by five d-orbitals.

1. f-orbital

The f-orbital has an even more complex shape. In space, the f-orbital can be placed in seven different ways. Therefore, the f-sublevel is formed by seven f-orbitals.

The electron shell of an atom is like a puff pastry. It also has layers. Electrons located on different layers have different energies: on layers closer to the nucleus - less, on those far from the nucleus - more. These layers are called energy levels.

Filling of electron orbitals.

The first energy level has only the s-sublevel:

At the second energy level, there is an s-sublevel and a p-sublevel appears:

At the third energy level, there is an s-sublevel, a p-sublevel, and a d-sublevel appears:

At the fourth energy level, in principle, an f-sublevel is added. But in the school course, f-orbitals are not filled, so we can not depict the f-sublevel:

The number of energy levels in an atom of an element is period number. When filling electron orbitals, the following principles should be followed:

1. Each electron tries to occupy the position in the atom where its energy will be minimal. That is, first the first energy level is filled, then the second, and so on.

To describe the structure of the electron shell, the electronic formula is also used. The electronic formula is a short one-line record of the distribution of electrons by sublevels.

1. At the sublevel, each electron first fills a vacant orbital. And each has spin +½ (up arrow).

And only after there is one electron in each sublevel orbital, the next electron becomes paired - that is, it occupies an orbital that already has an electron:

1. d-sublevel is filled in a special way.

The fact is that the energy of the d-sublevel is higher than the energy of the s-sublevel of the NEXT energy layer. And as we know, the electron tries to take that position in the atom, where its energy will be minimal.

Therefore, after filling the 3p sublevel, the 4s sublevel is filled first, after which the 3d sublevel is filled.

And only after the 3d sublevel is completely filled, the 4p sublevel is filled.

It is the same with the 4th energy level. After the 4p sublevel is filled, the 5s sublevel is filled next, followed by the 4d sublevel. And after it only 5p.

1. And there is one more point, one rule regarding the filling of the d-sublevel.

Then there is a phenomenon called failure. In case of failure, one electron from the s-sublevel of the next energy level literally falls to the d-electron.

Ground and excited states of the atom.

The atoms whose electronic configurations we have now built are called atoms in basic condition. That is, this is a normal, natural, if you like, state.

When an atom receives energy from outside, excitation can occur.

Excitation is the transition of a paired electron to an empty orbital, within the outer energy level.

For example, for a carbon atom:

Excitation is characteristic of many atoms. This must be remembered, because excitation determines the ability of atoms to bind to each other. The main thing to remember is the condition under which excitation can occur: a paired electron and an empty orbital in the outer energy level.

There are atoms that have several excited states:

Electronic configuration of the ion.

Ions are particles that atoms and molecules turn into by gaining or losing electrons. These particles have a charge, because they either "not enough" electrons, or their excess. Positively charged ions are called cations, negative - anions.

The chlorine atom (has no charge) gains an electron. The electron has a charge of 1- (one minus), respectively, a particle is formed that has an excess negative charge. Chlorine anion:

Cl 0 + 1e → Cl –

The lithium atom (also having no charge) loses an electron. An electron has a charge of 1+ (one plus), a particle is formed, with a lack of a negative charge, that is, its charge is positive. lithium cation:

Li 0 – 1e → Li +

Turning into ions, atoms acquire such a configuration that the external energy level becomes "beautiful", that is, completely filled. This configuration is the most thermodynamically stable, so there is a reason for atoms to turn into ions.

And therefore, the atoms of the elements of the VIII-A group (the eighth group of the main subgroup), as stated in the next paragraph, are noble gases, such are chemically inactive. They have the following structure in the ground state: the outer energy level is completely filled. Other atoms, as it were, tend to acquire the configuration of these most noble gases, and therefore turn into ions and form chemical bonds.

(Lecture notes)

The structure of the atom. Introduction.

The object of study in chemistry is the chemical elements and their compounds. chemical element A group of atoms with the same positive charge is called. Atom is the smallest particle of a chemical element that retains it Chemical properties. Connecting with each other, atoms of one or different elements form more complex particles - molecules. A collection of atoms or molecules form chemicals. Each individual chemical substance is characterized by a set of individual physical properties, such as boiling and melting points, density, electrical and thermal conductivity, etc.

1. The structure of the atom and the Periodic system of elements

DI. Mendeleev.

Knowledge and understanding of the regularities of the order of filling the Periodic system of elements D.I. Mendeleev allows us to understand the following:

1. the physical essence of the existence in nature of certain elements,

2. the nature of the chemical valency of the element,

3. the ability and "ease" of an element to give or receive electrons when interacting with another element,

4. the nature of the chemical bonds that a given element can form when interacting with other elements, the spatial structure of simple and complex molecules, etc., etc.

The structure of the atom.

An atom is a complex microsystem of elementary particles in motion and interacting with each other.

In the late 19th and early 20th centuries, it was found that atoms are composed of smaller particles: neutrons, protons and electrons. The last two particles are charged particles, the proton carries a positive charge, the electron is negative. Since the atoms of an element in the ground state are electrically neutral, this means that the number of protons in an atom of any element is equal to the number of electrons. The mass of atoms is determined by the sum of the masses of protons and neutrons, the number of which is equal to the difference between the mass of atoms and its serial number in the periodic system of D.I. Mendeleev.

In 1926, Schrodinger proposed to describe the motion of microparticles in the atom of an element using the wave equation he derived. When solving the Schrödinger wave equation for the hydrogen atom, three integer quantum numbers appear: n, ℓ and m ℓ , which characterize the state of an electron in three-dimensional space in the central field of the nucleus. quantum numbers n, ℓ and m ℓ take integer values. Wave function defined by three quantum numbers n, ℓ and m ℓ and obtained as a result of solving the Schrödinger equation is called an orbital. An orbital is a region of space in which an electron is most likely to be found. belonging to an atom of a chemical element. Thus, the solution of the Schrödinger equation for the hydrogen atom leads to the appearance of three quantum numbers, the physical meaning of which is that they characterize three different types of orbitals that an atom can have. Let's take a closer look at each quantum number.

Principal quantum number n can take any positive integer values: n = 1,2,3,4,5,6,7… It characterizes the energy of the electronic level and the size of the electronic "cloud". It is characteristic that the number of the main quantum number coincides with the number of the period in which the given element is located.

Azimuthal or orbital quantum numberℓ can take integer values ​​from ℓ = 0….up to n – 1 and determines the moment of electron motion, i.e. orbital shape. For various numerical values ​​of ℓ, the following notation is used: ℓ = 0, 1, 2, 3, and are denoted by symbols s, p, d, f, respectively for ℓ = 0, 1, 2 and 3. In the periodic table of elements there are no elements with a spin number ℓ = 4.

Magnetic quantum numberm ℓ characterizes the spatial arrangement of electron orbitals and, consequently, the electromagnetic properties of the electron. It can take values ​​from - ℓ to + ℓ , including zero.

The shape or, more precisely, the symmetry properties of atomic orbitals depend on quantum numbers ℓ and m ℓ . "electronic cloud", corresponding to s- orbitals has, has the shape of a ball (at the same time ℓ = 0).

Fig.1. 1s orbital

Orbitals defined by quantum numbers ℓ = 1 and m ℓ = -1, 0 and +1 are called p-orbitals. Since m ℓ in this case has three different values, then the atom has three energetically equivalent p-orbitals (the main quantum number for them is the same and can have the value n = 2,3,4,5,6 or 7). p-Orbitals have axial symmetry and have the form of three-dimensional eights, oriented along the x, y and z axes in an external field (Fig. 1.2). Hence the origin of the symbols p x , p y and p z .

Fig.2. p x , p y and p z -orbitals

In addition, there are d- and f-atomic orbitals, for the first ℓ = 2 and m ℓ = -2, -1, 0, +1 and +2, i.e. five AO, for the second ℓ = 3 and m ℓ = -3, -2, -1, 0, +1, +2 and +3, i.e. 7 AO.

fourth quantum m s called the spin quantum number, was introduced to explain some subtle effects in the spectrum of the hydrogen atom by Goudsmit and Uhlenbeck in 1925. The spin of an electron is the angular momentum of a charged elementary particle of an electron, the orientation of which is quantized, i.e. strictly limited to certain angles. This orientation is determined by the value of the spin magnetic quantum number (s), which for an electron is ½ , therefore, for an electron, according to the quantization rules m s = ± ½. In this regard, to the set of three quantum numbers, one should add the quantum number m s . We emphasize once again that four quantum numbers determine the order in which Mendeleev’s periodic table of elements is constructed and explain why there are only two elements in the first period, eight in the second and third, 18 in the fourth, and so on. However, in order to explain the structure of multielectron of atoms, the order of filling of electronic levels as the positive charge of an atom increases, it is not enough to have an idea about the four quantum numbers that "govern" the behavior of electrons when filling electronic orbitals, but you need to know some more simple rules, namely, Pauli's principle, Gund's rule and Klechkovsky's rules.

According to the Pauli principle in the same quantum state, characterized by certain values ​​of four quantum numbers, there cannot be more than one electron. This means that one electron can, in principle, be placed in any atomic orbital. Two electrons can be in the same atomic orbital only if they have different spin quantum numbers.

When filling three p-AOs, five d-AOs and seven f-AOs with electrons, one should be guided not only by the Pauli principle but also by the Hund rule: The filling of the orbitals of one subshell in the ground state occurs with electrons with the same spins.

When filling subshells (p, d, f) the absolute value of the sum of spins must be maximum.

Klechkovsky's rule. According to the Klechkovsky rule, when fillingd and forbital by electrons must be respectedprinciple of minimum energy. According to this principle, electrons in the ground state fill the orbits with minimum energy levels. The sublevel energy is determined by the sum of quantum numbersn + ℓ = E .

Klechkovsky's first rule: first fill those sublevels for whichn + ℓ = E minimal.

Klechkovsky's second rule: in case of equalityn + ℓ for several sublevels, the sublevel for whichn minimal .

Currently, 109 elements are known.

2. Ionization energy, electron affinity and electronegativity.

The most important characteristics of the electronic configuration of an atom are the ionization energy (EI) or ionization potential (IP) and the atom's electron affinity (SE). The ionization energy is the change in energy in the process of detachment of an electron from a free atom at 0 K: A = + + ē . The dependence of the ionization energy on the atomic number Z of the element, the size of the atomic radius has a pronounced periodic character.

Electron affinity (SE) is the change in energy that accompanies the addition of an electron to an isolated atom with the formation of a negative ion at 0 K: A + ē = A - (the atom and ion are in their ground states). In this case, the electron occupies the lowest free atomic orbital (LUAO) if the VZAO is occupied by two electrons. SE strongly depends on their orbital electronic configuration.

Changes in EI and SE correlate with changes in many properties of elements and their compounds, which is used to predict these properties from the values ​​of EI and SE. Halogens have the highest absolute electron affinity. In each group of the periodic table of elements, the ionization potential or EI decreases with increasing element number, which is associated with an increase in atomic radius and with an increase in the number of electron layers, and which correlates well with an increase in the element's reducing power.

Table 1 of the Periodic Table of the Elements gives the values ​​of EI and SE in eV/atom. Note that the exact SE values ​​are known only for a few atoms; their values ​​are underlined in Table 1.

Table 1

The first ionization energy (EI), electron affinity (SE) and electronegativity χ) of atoms in the periodic table.

χ	0.747 2. 1 0 0, 3 7																	1,2 2
χ	0.54 1. 55	-0.3 1. 1 3											0.2 0. 91	1.2 5	-0. 1 0, 55	1.47 0. 59	3.45 0. 64	1 ,60
χ	0. 7 4 1. 89	-0.3 1 . 3 1 1 . 6 0											0. 6	1.63	0.7	2.07	3.61
χ	2.3 6	- 0 .6						1.26(α)				-0.9 1 . 39	0. 18	1.2	0. 6	2.07	3.36
χ	2.4 8											-0.6 1 . 56	0. 2			2.2
χ	2.6 7	2, 2 1						Os

χ - Pauling electronegativity

r- atomic radius, (from "Laboratory and seminar classes in general and inorganic chemistry", N.S. Akhmetov, M.K. Azizova, L.I. Badygina)

An atom is the smallest particle of matter. Its study began in ancient Greece, when the attention of not only scientists, but also philosophers was riveted to the structure of the atom. What is the electronic structure of an atom, and what basic information is known about this particle?

The structure of the atom

Already ancient Greek scientists guessed the existence of the smallest chemical particles that make up any object and organism. And if in the XVII-XVIII centuries. chemists were sure that the atom is an indivisible elementary particle, then at the turn of the 19th-20th centuries, they managed to prove experimentally that the atom is not indivisible.

An atom, being a microscopic particle of matter, consists of a nucleus and electrons. The nucleus is 10,000 times smaller than an atom, but almost all of its mass is concentrated in the nucleus. The main characteristic of the atomic nucleus is that it has a positive charge and is made up of protons and neutrons. Protons are positively charged, while neutrons have no charge (they are neutral).

They are connected to each other by the strong nuclear force. The mass of a proton is approximately equal to the mass of a neutron, but at the same time it is 1840 times greater than the mass of an electron. Protons and neutrons have a common name in chemistry - nucleons. The atom itself is electrically neutral.

An atom of any element can be denoted by an electronic formula and an electronic graphic formula:

Rice. 1. Electron-graphic formula of the atom.

The only element in the Periodic Table that does not contain neutrons is light hydrogen (protium).

An electron is a negatively charged particle. The electron shell consists of electrons moving around the nucleus. Electrons have properties to be attracted to the nucleus, and between each other they are influenced by the Coulomb interaction. To overcome the attraction of the nucleus, the electrons must receive energy from an external source. The farther the electron is from the nucleus, the less energy is needed for this.

Atom Models

For a long time, scientists have sought to understand the nature of the atom. At an early stage, the ancient Greek philosopher Democritus made a great contribution. Although now his theory seems banal and too simple to us, at a time when the concept of elementary particles was just beginning to emerge, his theory of pieces of matter was taken quite seriously. Democritus believed that the properties of any substance depend on the shape, mass and other characteristics of atoms. So, for example, near fire, he believed, there are sharp atoms - therefore, fire burns; water has smooth atoms, so it can flow; in solid objects, in his view, the atoms were rough.

Democritus believed that absolutely everything consists of atoms, even the human soul.

In 1904, J. J. Thomson proposed his model of the atom. The main provisions of the theory boiled down to the fact that the atom was represented as a positively charged body, inside of which there were electrons with a negative charge. Later this theory was refuted by E. Rutherford.

Rice. 2. Thomson's model of the atom.

Also in 1904, the Japanese physicist H. Nagaoka proposed an early planetary model of the atom by analogy with the planet Saturn. According to this theory, electrons are united in rings and revolve around a positively charged nucleus. This theory turned out to be wrong.

In 1911, E. Rutherford, having done a series of experiments, concluded that the atom in its structure is similar to the planetary system. After all, electrons, like planets, move in orbits around a heavy positively charged nucleus. However, this description contradicted classical electrodynamics. Then the Danish physicist Niels Bohr in 1913 introduced the postulates, the essence of which was that the electron, being in some special states, does not radiate energy. Thus, Bohr's postulates showed that classical mechanics is inapplicable to atoms. The planetary model described by Rutherford and supplemented by Bohr was called the Bohr-Rutherford planetary model.

Rice. 3. Bohr-Rutherford planetary model.

Further study of the atom led to the creation of such a section as quantum mechanics, with the help of which many scientific facts were explained. Modern ideas about the atom have developed from the Bohr-Rutherford planetary model. Evaluation of the report

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