The structure of the electron shells of atoms is important for chemistry, they determine Chemical properties substances. The most important characteristic of the motion of an electron in a certain orbit is the energy of its connection with the nucleus. Electrons in an atom differ in a certain energy, and, as experiments show, some are attracted to the nucleus more strongly, others weaker. This is explained by the remoteness of electrons from the nucleus. The closer the electrons to the nucleus, the greater their bond with the nucleus, but the less energy. As the distance from the nucleus of the atom, the force of attraction of the electron to the nucleus decreases, and the energy supply increases. This is how electron layers are formed in the electron shell of an atom. Electrons with similar energy values ​​form a single electron layer, or energy level. The energy of electrons in an atom and the energy level is determined by the main quantum number n and takes integer values ​​1, 2, 3, 4, 5, 6 and 7. Than more value n, the greater the energy of the electron in the atom. The maximum number of electrons that can be in a particular energy level is determined by the formula:

Where N is the maximum number of electrons per level;

n is the number of the energy level.

It has been established that no more than two electrons are located on the first shell, no more than eight on the second, no more than 18 on the third, and no more than 32 on the fourth. We will not consider the filling of more distant shells. It is known that the external energy level can contain no more than eight electrons, it is called complete. Electronic layers that do not contain the maximum number of electrons are called incomplete.

The number of electrons in the outer energy level of the electron shell of an atom is equal to the group number for the chemical elements of the main subgroups.

As previously mentioned, the electron does not move in an orbit, but in an orbit and has no trajectory.

The space around the nucleus where a given electron is most likely to be is called that electron's orbital, or electron cloud.

Orbitals, or sub-levels, as they are also called, can have different shape, and their number corresponds to the level number, but does not exceed four. The first energy level has one sublevel (s), the second has two (s,p), the third has three (s,p,d), and so on. Electrons of different sublevels of the same level have different shapes of the electron cloud: spherical (s), dumbbell-shaped (p), and more complex configurations (d) and (f). Scientists agreed to call the spherical atomic orbital s-orbital. It is the most stable and is located quite close to the core.

The greater the energy of an electron in an atom, the faster it rotates, the more the region of its stay is extended, and, finally, it turns into a dumbbell-shaped p-orbital:

An electron cloud of this form can occupy three positions in an atom along the coordinate axes of space x, y and z. This is easily explained: after all, all electrons are negatively charged, so the electron clouds repel each other and tend to be located as far as possible from each other.

So, p There can be three orbitals. Their energy, of course, is the same, but their location in space is different.

Draw a diagram of the sequential filling of energy levels with electrons

Now we can draw up a diagram of the structure of the electron shells of atoms:

1. Determine the total number of electrons on the shell by the element's serial number.

2. Determine the number of energy levels in the electron shell. Their number is equal to the number of the period in the table of D. I. Mendeleev, in which the element is located.

3. Determine the number of electrons at each energy level.

4. Using Arabic numerals to designate the level and designating the orbitals with the letters s and p, and the number of electrons in a given orbital Arabic numeral at the top right above the letter, we depict the structure of atoms with more complete electronic formulas. Scientists agreed to designate each atomic orbital as a quantum cell - a square on the energy diagram:

On the s A sublevel can contain one atomic orbital

and on p-there may already be three sublevels -

(according to the three coordinate axes):

Orbitals d- and f- sublevels in an atom can already be five and seven, respectively:

The nucleus of a hydrogen atom has a charge of +1, so only one electron moves around its nucleus at a single energy level. Let's write down the electronic configuration of the hydrogen atom

To establish a connection between the structure of the atom of a chemical element and its properties, we will consider a few more chemical elements.

The next element after hydrogen is helium. The nucleus of a helium atom has a charge of +2, so a helium atom contains two electrons in the first energy level:

Since the first energy level can contain no more than two electrons, it is considered complete.

Element number 3 - lithium. The lithium nucleus has a charge of +3, therefore, there are three electrons in the lithium atom. Two of them are at the first energy level, and the third electron begins to fill the second energy level. First, the s-orbital of the first level is filled, then the s-orbital of the second level. The electron in the second level is weaker bound to the nucleus than the other two.

For a carbon atom, it is already possible to assume three possible schemes for filling electron shells in accordance with electron-graphic formulas:

An analysis of the atomic spectrum shows that the latter scheme is correct. Using this rule, it is not difficult to draw up a diagram of the electronic structure for the nitrogen atom:

This scheme corresponds to the formula 1s22s22p3. Then the pairwise placement of electrons into 2p orbitals begins. Electronic formulas of the remaining atoms of the second period:

The filling of the second energy level of the neon atom ends, and the construction of the second period of the system of elements is completed.

Find the chemical sign of lithium in the periodic system, from lithium to neon Ne, the charge of atomic nuclei naturally increases. The second layer is gradually filled with electrons. With an increase in the number of electrons in the second layer, the metallic properties of the elements gradually weaken and are replaced by non-metallic ones.

The third period, like the second, begins with two elements (Na, Mg), in which the electrons are located on the s-sublevel of the outer electron layer. This is followed by six elements (from Al to Ar), in which the p-sublevel of the outer electron layer is formed. The structure of the outer electronic layer of the corresponding elements of the second and third periods is similar. In other words, with an increase in the charge of the nucleus, the electronic structure of the outer layers of atoms is periodically repeated. If the elements have the same external energy levels, then the properties of these elements are similar. For example, argon and neon contain eight electrons at the outer level, and therefore they are inert, that is, they almost do not enter into chemical reactions. In the free form, argon and neon are gases that have monatomic molecules.

The atoms of lithium, sodium and potassium contain one electron at the outer level and have similar properties, so they are placed in the same group of the periodic system.

III. Findings.

1. The properties of chemical elements, arranged in ascending order of the charge of the nucleus, are periodically repeated, since the structure of the external energy levels of the atoms of the elements is periodically repeated.

2. A smooth change in the properties of chemical elements within one period can be explained by a gradual increase in the number of electrons at the external energy level.

3. The reason for the similarity of the properties of chemical elements belonging to the same family lies in the same structure of the external energy levels of their atoms.

Electronic configuration atom is a formula showing the arrangement of electrons in an atom by levels and sublevels. After studying the article, you will find out where and how electrons are located, get acquainted with quantum numbers and be able to build the electronic configuration of an atom by its number, at the end of the article there is a table of elements.

Why study the electronic configuration of elements?

Atoms are like a constructor: there are a certain number of parts, they differ from each other, but two parts of the same type are exactly the same. But this constructor is much more interesting than the plastic one, and here's why. The configuration changes depending on who is nearby. For example, oxygen next to hydrogen maybe turn into water, next to sodium into gas, and being next to iron completely turns it into rust. To answer the question why this happens and to predict the behavior of an atom next to another, it is necessary to study the electronic configuration, which will be discussed below.

How many electrons are in an atom?

An atom consists of a nucleus and electrons revolving around it, the nucleus consists of protons and neutrons. In the neutral state, each atom has the same number of electrons as the number of protons in its nucleus. The number of protons was designated by the serial number of the element, for example, sulfur has 16 protons - the 16th element of the periodic system. Gold has 79 protons - the 79th element of the periodic table. Accordingly, there are 16 electrons in sulfur in the neutral state, and 79 electrons in gold.

Where to look for an electron?

Observing the behavior of an electron, certain patterns were derived, they are described by quantum numbers, there are four of them in total:

• Principal quantum number
• Orbital quantum number
• Magnetic quantum number
• Spin quantum number

Orbital

Further, instead of the word orbit, we will use the term "orbital", the orbital is the wave function of the electron, roughly - this is the area in which the electron spends 90% of the time.

N - level
L - shell
M l - orbital number
M s - the first or second electron in the orbital

Orbital quantum number l

As a result of the study of the electron cloud, it was found that depending on the level of energy, the cloud takes four main forms: a ball, dumbbells and the other two, more complex. In ascending order of energy, these forms are called s-, p-, d- and f-shells. Each of these shells can have 1 (on s), 3 (on p), 5 (on d) and 7 (on f) orbitals. The orbital quantum number is the shell on which the orbitals are located. The orbital quantum number for s, p, d and f orbitals, respectively, takes the values ​​0,1,2 or 3.

On the s-shell one orbital (L=0) - two electrons
There are three orbitals on the p-shell (L=1) - six electrons
There are five orbitals on the d-shell (L=2) - ten electrons
There are seven orbitals (L=3) on the f-shell - fourteen electrons

Magnetic quantum number m l

There are three orbitals on the p-shell, they are denoted by numbers from -L to +L, that is, for the p-shell (L=1) there are orbitals "-1", "0" and "1". The magnetic quantum number is denoted by the letter m l .

Inside the shell, it is easier for electrons to be located in different orbitals, so the first electrons fill one for each orbital, and then its pair is added to each.

Consider a d-shell:
The d-shell corresponds to the value L=2, that is, five orbitals (-2,-1,0,1 and 2), the first five electrons fill the shell, taking the values ​​M l =-2,M l =-1,M l =0 , M l =1,M l =2.

Spin quantum number m s

Spin is the direction of rotation of an electron around its axis, there are two directions, so the spin quantum number has two values: +1/2 and -1/2. Only two electrons with opposite spins can be on the same energy sublevel. The spin quantum number is denoted m s

Principal quantum number n

The main quantum number is the energy level, at the moment seven energy levels are known, each is denoted by an Arabic numeral: 1,2,3,...7. The number of shells at each level is equal to the level number: there is one shell on the first level, two on the second, and so on.

Electron number

So, any electron can be described by four quantum numbers, the combination of these numbers is unique for each position of the electron, let's take the first electron, the lowest energy level is N=1, one shell is located on the first level, the first shell at any level has the shape of a ball (s -shell), i.e. L=0, the magnetic quantum number can take only one value, M l =0 and the spin will be equal to +1/2. If we take the fifth electron (in whatever atom it is), then the main quantum numbers for it will be: N=2, L=1, M=-1, spin 1/2.

Electronic configurations of atoms

The total number of electrons in an atom is determined by the charge of its nucleus, i.e., the proton number. It is equal to the atomic number of the element. Electrons, depending on their energy, are distributed in the atom over energy levels and sublevels, each of which consists of a certain number of orbitals.

The distribution of electrons is expressed using the electronic formulas (or electronic configurations) of the atom. For example, hydrogen, an element with atomic number 1, has an electronic formula: 1H 1s1. In this formula, the number of the energy level is written as a number, followed by a letter indicating the type of sublevel, and finally, the number at the top right indicates the number of electrons in this sublevel.

Schematically, the electronic structure of an atom is depicted using an electronic graphic diagram in which orbitals are represented as cells, and electrons as arrows.

The electronic graphic scheme of the hydrogen atom is written as follows:

For the correct representation of electronic formulas, several basic rules must be observed.

1st rule: The distribution of electrons in an atom, which is in the ground (most stable) state, is determined by the principle of minimum energy: the ground state of the atom corresponds to the lowest possible energy levels and sublevels.

Therefore, electrons (for atoms of elements of the first three periods) fill the orbitals in order of increasing their energy:

1s→2s→2p→3s→3p

2nd rule: No more than two electrons can be in each orbital, and with opposite spins.

Thus, helium 2He following hydrogen has the electronic formula:

2Not 1s2 ,

Since only two electrons can be on the first electron layer, this layer in the helium atom is complete and, therefore, very stable.

For atoms of elements of the second period, the second energy level is filled, on which no more than 8 electrons can be located. First, electrons fill the 2s orbital (for lithium and beryllium atoms):

Since the 2s orbital is filled, the fifth electron at the boron atom B occupies one of the three 2p orbitals. The electronic formula of the boron atom:

and the electronic graphic scheme:

Please note that the 2p sublevel is depicted close to the 2s sublevel, but slightly higher. This emphasizes his belonging to the same level (second) and at the same time a greater supply of energy.

3rd rule. Sets the order in which the orbitals of one sublevel are filled. Electrons of one sublevel first fill the orbitals one by one (i.e., all empty), and if the number of electrons is greater than the number of orbitals, then two each. Therefore, the electronic formulas of carbon and nitrogen atoms are:

6C 1s22s22p2 and 7N 1s22s22p3

and electronic graphic schemes:

At atoms of oxygen, fluorine and neon, the number of electrons increases, and they are forced to be placed on p-orbitals of the second energy level by two:

6O 1s22s22p4; 6F 1s22s22p5; 6Ne 1s22s22p6

Electronic graphic schemes of atoms of these elements:

The electronic configuration of the outer layer 2s22p6 corresponds to its complete filling and is therefore stable.

In the atoms of the elements of the third period, the third electron layer begins to form. First, the s-sublevel of sodium and magnesium is filled with electrons:

11Na 1s22s22p63s1 12Mg 1s22s22p63s2

and then the p-sublevel of aluminum, silicon, chlorine and argon:

18Ar 1s22s22p63s23p6

Electronic graphic scheme for the argon atom:

In an argon atom, there are 8 electrons on the outer electron layer. Therefore, it is completed, since in the atom of any element at the external energy level, there can be no more than 8 electrons at most.

The building up of the third electron layer is not exhausted by this. In accordance with the formula 2n2, 18 electrons can be on it: 8 on the s- and p-sublevels and 10 on the d-sublevel. This sublevel will be formed for the elements of the fourth period. But first, the first two elements of the fourth period - potassium and calcium - have a fourth electron layer, which opens with an s-sublevel (the energy of the 4s sublevel is slightly less than that of the 3d sublevel:

19K 1s22s22p63s23p64s1 and 19Ca 1s22s22p63s23p64s2

Only after this will the d-sublevel of the third, now preexternal, energy level begin to be filled with electrons. The electronic configuration of the scandium atom:

21Sc 1s22s22p63s23p64s23d1,

titanium atom:

21Ti 1s22s22p63s23p64s23d2,

and so on, up to zinc. The electronic configuration of its atom is:

21Zn 1s22s22p63s23p64s23d10,

and the electronic graphic scheme:

Since for elements of the fourth period only the orbitals of the third and fourth energy levels are filled with electrons, the completely filled levels (in this case, the first and second) are usually not indicated on the electronic graphic diagrams. Instead of them, in electronic formulas, they write the symbol of the nearest element of the VIII A-group with completely filled energy s- and p-sublevels: for example, the electronic formula of chlorine is 3s23p5, zinc is 3d104s2, and antimony is 51Sb -4d105s25p3

In addition to electronic formulas and electronic graphic diagrams, electronic diagrams of atoms are sometimes used, in which only the number of electrons at each energy level (electronic layer) is indicated:

The electronic structure of an atom is determined by the charge of its nucleus, which is equal to the atomic number of the element in the periodic system.

The distribution of electrons over energy levels, sublevels and orbitals is displayed using electronic formulas and electronic graphic diagrams, as well as electronic diagrams of atoms.

An atom of any element can have no more than 8 electrons on the outer electron layer. 3.2. Types of chemical bonds

covalent bond- the most general type of chemical bond that arises due to the socialization of an electron pair through exchange mechanism, when each of the interacting atoms supplies one electron, or donor-acceptor mechanism, if the electron pair is transferred to the common use of one atom (donor) to another atom (acceptor) (Fig. 3.2).

A classic example of a non-polar covalent bond (the difference in electronegativity is zero) is observed in homonuclear molecules: H–H, F–F. The energy of a two-electron two-center bond lies in the range of 200–2000 kJ⋅mol–1.

When a heteroatomic covalent bond is formed, the electron pair is shifted to a more electronegative atom, which makes such a bond polar. The ionicity of the polar bond in percent is calculated by the empirical relation 16(χ A - χ B) + 3.5(χ A - χ B) 2, where χ A and χ B are the electronegativity of atoms A and B of the AB molecule. Except polarizability covalent bond has the property satiety- the ability of an atom to form as many covalent bonds as it has energetically available atomic orbitals. About the third property of a covalent bond - focus will be discussed below (cf. valence bond method).

Ionic bond- a special case of covalent, when the formed electron pair completely belongs to a more electronegative atom, which becomes an anion. The basis for separating this bond into a separate type is the fact that compounds with such a bond can be described in the electrostatic approximation, considering the ionic bond due to the attraction of positive and negative ions. The interaction of ions of the opposite sign does not depend on the direction, and the Coulomb forces do not have the property of saturation. Therefore, each ion in an ionic compound attracts such a number of ions of the opposite sign that an ionic-type crystal lattice is formed. There are no molecules in an ionic crystal. Each ion is surrounded by a certain number of ions of a different sign (coordination number of the ion). Ion pairs can exist in the gaseous state as polar molecules. In the gaseous state, NaCl has a dipole moment of ~3∙10 –29 C∙m, which corresponds to a shift of 0.8 electron charge per bond length of 0.236 nm from Na to Cl, i.e. Na 0.8+ Cl 0.8– .

The metallic bond arises as a result of partial delocalization of valence electrons, which move quite freely in the metal lattice, electrostatically interacting with positively charged ions. Bonding forces are not localized and not directed, and delocalized electrons cause high thermal and electrical conductivity.

hydrogen bond. Its formation is due to the fact that as a result of a strong displacement of an electron pair to an electronegative atom, a hydrogen atom with an effective positive charge can interact with another electronegative atom (F, O, N, less often Cl, Br, S). The energy of such an electrostatic interaction is 20–100 kJ∙mol–1. Hydrogen bonds can be intra- and intermolecular. An intramolecular hydrogen bond is formed, for example, in acetylacetone and is accompanied by cycle closure (Fig. 3.3).

molecules carboxylic acids in non-polar solvents, they dimerize due to two intermolecular hydrogen bonds (Fig. 3.4).

The hydrogen bond plays an extremely important role in biological macromolecules, such inorganic compounds as H 2 O, H 2 F 2, NH 3. Due to hydrogen bonds, water is characterized by such high melting and boiling points compared to H 2 E (E = S, Se, Te). If there were no hydrogen bonds, then water would melt at –100°C and boil at –80°C.

Van der Waals (intermolecular) bond- the most universal type of intermolecular bond, due to dispersion forces(induced dipole - induced dipole), induction interaction (permanent dipole - induced dipole) and orientation interaction (permanent dipole - permanent dipole). The energy of the van der Waals bond is less than the hydrogen bond and is 2–20 kJ∙mol–1.

Chemical bond in solids. The properties of solids are determined by the nature of the particles occupying the nodes of the crystal lattice and the type of interaction between them.

Solid argon and methane form atomic and molecular crystals, respectively. Since the forces between atoms and molecules in these lattices are of the weak van der Waals type, such substances melt at fairly low temperatures. Most of the substances that are liquid and gaseous at room temperature form molecular crystals at low temperatures.

The melting points of ionic crystals are higher than those of atomic and molecular crystals because the electrostatic forces acting between the ions far exceed the weak van der Waals forces. Ionic compounds are harder and more brittle. Such crystals are formed by elements with very different electronegativity (for example, alkali metal halides). Ionic crystals containing polyatomic ions have lower melting points; so for NaCl t pl. = 801 °C, and for NaNO 3 t pl = 306.5 °C.

In covalent crystals, the lattice is built of atoms connected by a covalent bond, so these crystals have high hardness, melting point and low thermal and electrical conductivity.

Crystal lattices formed by metals are called metallic. The nodes of such lattices contain positive metal ions, and the interstitials contain valence electrons (electron gas).

Among metals, d-elements have the highest melting point, which is explained by the presence in the crystals of these elements of a covalent bond formed by unpaired d-electrons, in addition to the metallic bond formed by s-electrons.

Valence bond method(MVS) otherwise called the theory of localized electron pairs, since the method is based on the assumption that the chemical bond between two atoms is carried out using one or more electron pairs, which are localized mainly between them. In contrast to MMO, in which the simplest chemical bond can be either two- or multicenter, in MVS it is always two-electron and necessarily two-center. The number of elementary chemical bonds that an atom or ion can form is equal to its valence. Just as in MMO, valence electrons take part in the formation of a chemical bond. The wave function that describes the state of the electrons that form a bond is called a localized orbital (LO).

Note that the electrons described by LO, in accordance with the principle pauli must have oppositely directed spins, that is, in the MVS all spins are paired, and all molecules must be diamagnetic. Consequently, the MVS fundamentally cannot explain the magnetic properties of molecules.

Nevertheless, the principle of localized connections has a number of important advantages, one of which is its extreme visibility. The MHS is quite good, for example, at predicting the valence possibilities of atoms and the geometry of the resulting molecule. The latter circumstance is associated with the so-called AO hybridization. It was introduced to explain the fact that two-electron two-center chemical bonds formed by AO in different energy states have the same energy. So, Be * (2s 1 1p 1), B * (2s 1 2p 2), C * (2s 1 2p 3) form two, three and four bonds, respectively, due to s- and p-orbitals, and therefore one of them should be stronger than others. However, experience shows that in BeH 2 , BCl 3 , CH 4 all bonds are equivalent. For BeH 2, the bond angle is 180°, for BCl 3 it is 120°, and for CH 4 it is 109°28".

According to the concept of hybridization, chemical bonds are formed by mixed - hybrid orbitals (GO), which are a linear combination of AO of a given atom (s- and p-AO Be, B, C), have the same energy and shape, a certain orientation in space (symmetry ). Thus, s- and p-orbitals give two sp-GOs located at an angle of 180° relative to each other.

In the CH 4 molecule, hybrid orbitals of four carbon AOs (one s and three p) are called sp 3 orbitals; they are completely equivalent energetically and are spatially directed to the tetrahedron vertices.

Thus, when one atom forms several bonds, and its valence electrons belong to different orbitals (s and p; s, p and d), to explain the geometry of molecules in the MVS, it is necessary to involve the theory of hybridization of atomic orbitals. The main provisions of the theory are as follows:

The introduction of hybrid orbitals serves to describe directed localized bonds. Hybrid orbitals provide the maximum overlap of AO in the direction of localized σ-bonds.

The number of hybrid orbitals is equal to the number of AO involved in hybridization.

Valence AOs close in energy are hybridized, regardless of whether they are completely filled, half-filled, or empty in the atom.

AOs that have common signs of symmetry participate in hybridization.

According to Table. 3.3 hybrid orbitals give molecules with angles of 180°, 120°, 109°28", 90°. These are regular geometric figures. Such molecules are formed when all peripheral atoms in a multi-electron molecule (or ion) are the same and their number coincides with the number of hybrid orbitals However, if the number of hybrid orbitals is greater than the number of bonded atoms, then a part of the hybrid orbitals is populated by electron pairs that do not participate in bond formation, - non-binding or undivided electronic pairs.

H–Be–H, HC≡CH

H 2 C \u003d CH 2, C 6 H 6, BCl 3

CH 4 , CCl 4 , H 3 C–CH 3

d 2 sp 3 or sp 3 d 2

As an example, consider NH 3 and H 2 O molecules. Nitrogen and oxygen atoms are prone to sp 3 hybridization. Nitrogen on sp 3 -GO, in addition to the three bonding pairs of electrons that form a bond with three hydrogen atoms, has one non-bonding pair. It is she who, occupying one sp 3 -GO, distorts the angle of the H–N–H bond to 107.3°. There are two such non-bonding pairs in the H 2 O molecule, and the H–O–H angle is 104.5° (Fig. 3.17).

The electrons of bonding and non-bonding pairs interact differently. The stronger the interelectronic repulsion, the larger the conditional surface on the sphere occupied by the electron pair. For a qualitative explanation of experimental facts, it is usually assumed that non-bonding pairs occupy a larger volume than bonding pairs, and the volume of bonding pairs is the smaller, the greater the electronegativity of peripheral atoms (method Gillespie).

Physical properties of metals.

Density. This is one of the most important characteristics of metals and alloys. By density, metals are divided into the following groups:

lungs(density not more than 5 g / cm 3) - magnesium, aluminum, titanium, etc .:

heavy- (density from 5 to 10 g / cm 3) - iron, nickel, copper, zinc, tin, etc. (this is the most extensive group);

very heavy(density more than 10 g / cm 3) - molybdenum, tungsten, gold, lead, etc.

Table 2 shows the density values ​​of metals. (This and subsequent tables characterize the properties of those metals that form the basis of alloys for artistic casting).

Table 2. Density of metal.

Melting temperature. Depending on the melting temperature, the metal is divided into the following groups:

fusible(melting point does not exceed 600 o C) - zinc, tin, lead, bismuth, etc.;

medium melting(from 600 o C to 1600 o C) - these include almost half of the metals, including magnesium, aluminum, iron, nickel, copper, gold;

refractory(more than 1600 o C) - tungsten, molybdenum, titanium, chromium, etc.

Mercury is a liquid.

In the manufacture of art castings, the melting temperature of the metal or alloy determines the choice of melting unit and refractory molding material. When additives are introduced into the metal, the melting temperature, as a rule, decreases.

Table 3. Melting and boiling points of metals.

Specific heat. This is the amount of energy required to raise the temperature of a unit mass by one degree. The specific heat capacity decreases with an increase in the element's serial number in the periodic table. The dependence of the specific heat of an element in the solid state on the atomic mass is described approximately by the Dulong and Petit law:

m a c m = 6.

where, m a- atomic mass; c m- specific heat capacity (J / kg * o C).

Table 4 shows the values ​​of the specific heat capacity of some metals.

Table 4. Specific heat capacity of metals.

Algorithm for compiling the electronic formula of an element:

1. Determine the number of electrons in an atom using the Periodic Table of Chemical Elements D.I. Mendeleev.

2. By the number of the period in which the element is located, determine the number of energy levels; the number of electrons in the last electronic level corresponds to the group number.

3. Divide the levels into sublevels and orbitals and fill them with electrons in accordance with the rules for filling orbitals:

It must be remembered that the first level has a maximum of 2 electrons. 1s2, on the second - a maximum of 8 (two s and six R: 2s 2 2p 6), on the third - a maximum of 18 (two s, six p, and ten d: 3s 2 3p 6 3d 10).

• Principal quantum number n should be minimal.
• Filled in first s- sublevel, then p-, d-b f- sublevels.
• Electrons fill orbitals in ascending order of orbital energy (Klechkovsky's rule).
• Within the sublevel, electrons first occupy free orbitals one at a time, and only after that they form pairs (Hund's rule).
• There cannot be more than two electrons in one orbital (Pauli principle).

Examples.

1. Compose the electronic formula of nitrogen. Nitrogen is number 7 on the periodic table.

2. Compose the electronic formula of argon. In the periodic table, argon is at number 18.

1s 2 2s 2 2p 6 3s 2 3p 6.

3. Compose the electronic formula of chromium. In the periodic table, chromium is number 24.

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

Energy diagram of zinc.

4. Compose the electronic formula of zinc. In the periodic table, zinc is number 30.

1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10

Note that part of the electronic formula, namely 1s 2 2s 2 2p 6 3s 2 3p 6 is the electronic formula of argon.

The electronic formula of zinc can be represented as.

To correctly depict the electronic configurations of atoms, you need to answer the following questions: 1. How to determine the total number of electrons in an atom? 2. What is the maximum number of electrons at the levels, sublevels? 3. What is the order of filling sublevels and orbitals? 3

Electronic configurations (on the example of a hydrogen atom) 1. Scheme of the electronic structure The scheme of the electronic structure of atoms shows the distribution of electrons over energy levels 2. Electronic formula 1s 1, where s is the designation of the sublevel; 1 - number of electrons Electron formulas of atoms show the distribution of electrons over energy sublevels 3. Electron graphic formula Electron graphic formulas of atoms show the distribution of electrons in orbitals and electron spins 4

2. According to the sample, make up the electronic formula of aluminum. The order in which the energy levels in the atom are filled. 1s 2, 2s 2, 2p 6, 3s 2, 3p 1 6 Aluminum has 13 electrons The 1s sublevel is the first to be filled in the atom. It can have a maximum of 2 electrons, mark them and subtract them from the total number of electrons. It remains to place 11 electrons. The next 2s sublevel is filled, it can have 2 electrons. It remains to place 9 electrons. The next 2p sublevel is filled, it can have 6 electrons. Next, fill in the 3s sublevel. We reached the 3p sublevel, it can have a maximum of 6 electrons, but only 1 remains, and we place it. 1s = Al s2s2s 2p2p 3p - 2 = - 6 = - 2 = 9 3 1

3. Determine: Are the energy levels in order. If the levels are in order, then leave them like that. If the levels are not in order, then rewrite them, arranging them in ascending order. No. 4s and 3d sublevels are out of order. It is necessary to rewrite and arrange them as they increase. 7 Cr 24 1s 2 2p62p6 3s 2 4s 2 3p 6 3d 4 2s22s2 1s 2 2p62p6 3s 2 4s 2 3p 6 3d 4 2s22s2

Rules for drawing up an electronic graphic diagram Each sublevel has a certain number of orbitals Each orbital can contain no more than two electrons. If there are two electrons in an orbital, then they must have a different spin (the arrows look in different directions). 8 s p d f
5. Geographical journey Determine in which groups of the periodic system are the chemical elements whose electronic formulas of atoms are given in the first column of the table. The letters corresponding to the correct answers will give the name of the country. 10 JAMAICA Electronic formulas of Group IIIIIIIVVVIVII 1s 2 2s 1 YAGLRKAO 1s 2 2s 2 2p 6 3s 2 3p 5 VISNPDM 1s 2 2s 2 2p 6 3s 2 3p 4 EFTZYAO 1s 2 2s 2 2p 4 GRISYK 1s 3p 2 s 2s 2 4s 1 CUERMIP 1s 2 2s 2 2p 6 3s 1 ANLO