Electronic configuration an atom is a numerical representation of its electron orbitals. Electron orbitals are areas various shapes, located around the atomic nucleus, in which the electron is mathematically probable. The electronic configuration helps to quickly and easily tell the reader how many electron orbitals an atom has, as well as determine the number of electrons in each orbital. After reading this article, you will master the method of compiling electronic configurations.

Steps

Distribution of electrons using the periodic system of D. I. Mendeleev

Find the atomic number of your atom. Each atom has a certain number of electrons associated with it. Find the symbol for your atom in the periodic table. The atomic number is a positive integer starting from 1 (for hydrogen) and increasing by one for each subsequent atom. The atomic number is the number of protons in an atom, and therefore it is also the number of electrons in an atom with zero charge.

Determine the charge of an atom. Neutral atoms will have the same number of electrons as shown in the periodic table. However, charged atoms will have more or fewer electrons, depending on the magnitude of their charge. If you are working with a charged atom, add or subtract electrons as follows: add one electron for every negative charge and subtract one for every positive charge.

• For example, a sodium atom with a charge of -1 will have an extra electron in addition to its base atomic number of 11. In other words, an atom will have 12 electrons in total.
• If we are talking about a sodium atom with a charge of +1, one electron must be subtracted from the base atomic number 11. So the atom will have 10 electrons.

1. Memorize the basic list of orbitals. As the number of electrons increases in an atom, they fill the various sublevels of the electron shell of the atom according to a certain sequence. Each sublevel of the electron shell, when filled, contains an even number of electrons. There are the following sublevels:

   Understand the electronic configuration record. Electronic configurations are written down in order to clearly reflect the number of electrons in each orbital. Orbitals are written sequentially, with the number of atoms in each orbital written as a superscript to the right of the orbital name. The completed electronic configuration has the form of a sequence of sublevel designations and superscripts.

   • Here, for example, is the simplest electronic configuration: 1s 2 2s 2 2p 6 . This configuration shows that there are two electrons in the 1s sublevel, two electrons in the 2s sublevel, and six electrons in the 2p sublevel. 2 + 2 + 6 = 10 electrons in total. This is the electronic configuration of the neutral neon atom (neon atomic number is 10).

2. Remember the order of the orbitals. Keep in mind that electron orbitals are numbered in ascending order of electron shell number, but arranged in ascending energy order. For example, a filled 4s 2 orbital has less energy (or less mobility) than a partially filled or filled 3d 10, so the 4s orbital is written first. Once you know the order of the orbitals, you can easily fill them in according to the number of electrons in the atom. The order in which the orbitals are filled is as follows: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p.

   • The electronic configuration of an atom in which all orbitals are filled will have the following form: 10 7p 6
   • Note that the above notation, when all orbits are filled, is the electron configuration of the element Uuo (ununoctium) 118, the highest numbered atom in the Periodic Table. Therefore, this electronic configuration contains all currently known electronic sublevels of a neutrally charged atom.

3. Fill in the orbitals according to the number of electrons in your atom. For example, if we want to write down the electronic configuration of a neutral calcium atom, we must start by looking up its atomic number in the periodic table. Its atomic number is 20, so we will write the configuration of an atom with 20 electrons according to the above order.

   • Fill in the orbitals in the above order until you reach the twentieth electron. The first 1s orbital will have two electrons, the 2s orbital will also have two, the 2p orbital will have six, the 3s orbital will have two, the 3p orbital will have 6, and the 4s orbital will have 2 (2 + 2 + 6 +2 +6 + 2 = 20 .) In other words, the electronic configuration of calcium has the form: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 .
   • Note that the orbitals are in ascending order of energy. For example, when you are ready to move to the 4th energy level, then first write down the 4s orbital, and then 3d. After the fourth energy level, you move on to the fifth, where the same order is repeated. This happens only after the third energy level.

4. Use the periodic table as a visual cue. You have probably already noticed that the shape of the periodic table corresponds to the order of electronic sublevels in electronic configurations. For example, atoms in the second column from the left always end in "s 2 ", while atoms on the right edge of the thin middle section always end in "d 10 ", and so on. Use the periodic table as a visual guide to writing configurations - as the order in which you add to the orbitals corresponds to your position in the table. See below:

   • In particular, the two leftmost columns contain atoms whose electronic configurations end in s-orbitals, the right-hand block of the table contains atoms whose configurations end in p-orbitals, and at the bottom of the atoms end in f-orbitals.
   • For example, when you write down the electronic configuration of chlorine, think like this: "This atom is located in the third row (or "period") of the periodic table. It is also located in the fifth group of the orbital block p of the periodic table. Therefore, its electronic configuration will end with. ..3p 5
   • Note that the elements in the d and f orbital regions of the table have energy levels that do not correspond to the period in which they are located. For example, the first row of a block of elements with d-orbitals corresponds to 3d orbitals, although it is located in the 4th period, and the first row of elements with f-orbitals corresponds to the 4f orbital, despite the fact that it is located in the 6th period.

5. Learn the abbreviations for writing long electronic configurations. The atoms on the right side of the periodic table are called noble gases. These elements are chemically very stable. To shorten the process of writing long electronic configurations, simply write in square brackets the chemical symbol for the nearest noble gas with fewer electrons than your atom, and then continue to write the electronic configuration of subsequent orbital levels. See below:

   • To understand this concept, it will be helpful to write an example configuration. Let's write the configuration of zinc (atomic number 30) using the noble gas abbreviation. The complete zinc configuration looks like this: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 . However, we see that 1s 2 2s 2 2p 6 3s 2 3p 6 is the electronic configuration of argon, a noble gas. Simply replace the electronic configuration part of zinc with the chemical symbol for argon in square brackets (.)
   • So, the electronic configuration of zinc, written in abbreviated form, is: 4s 2 3d 10 .
   • Note that if you are writing the electronic configuration of a noble gas, say argon, you cannot write! One must use the abbreviation of the noble gas in front of this element; for argon it will be neon ().

   Using ADOMAH Periodic Table

   1. Master the ADOMAH periodic table. This method of recording the electronic configuration does not require memorization, however, it requires a modified periodic table, since in the traditional periodic table, starting from the fourth period, the period number does not correspond to the electron shell. Find the ADOMAH periodic table, a special type of periodic table designed by scientist Valery Zimmerman. It is easy to find with a short internet search.

      • In the ADOMAH periodic table, the horizontal rows represent groups of elements such as halogens, noble gases, alkali metals, alkaline earth metals etc. Vertical columns correspond to electronic levels, and the so-called "cascades" (diagonal lines connecting blocks s,p,d and f) correspond to periods.
      • Helium is moved to hydrogen, since both of these elements are characterized by a 1s orbital. The period blocks (s,p,d and f) are shown on the right side and the level numbers are given at the bottom. Elements are represented in boxes numbered from 1 to 120. These numbers are the usual atomic numbers, which represent the total number of electrons in a neutral atom.

   2. Find your atom in the ADOMAH table. To write down the electronic configuration of an element, find its symbol in the ADOMAH periodic table and cross out all elements with a higher atomic number. For example, if you need to write down the electronic configuration of erbium (68), cross out all the elements from 69 to 120.

      • Pay attention to the numbers from 1 to 8 at the base of the table. These are the electronic level numbers, or column numbers. Ignore columns that contain only crossed out items. For erbium, columns with numbers 1,2,3,4,5 and 6 remain.

   3. Count the orbital sublevels up to your element. Looking at the block symbols shown to the right of the table (s, p, d, and f) and the column numbers shown at the bottom, ignore the diagonal lines between the blocks and break the columns into block-columns, listing them in order from bottom to top. And again, ignore the blocks in which all the elements are crossed out. Write the column blocks starting from the column number followed by the block symbol, thus: 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 6s (for erbium).

      • Please note: The above electronic configuration Er is written in ascending order of the electronic sublevel number. It can also be written in the order in which the orbitals are filled. To do this, follow the cascades from bottom to top, not columns, when you write column blocks: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 12 .

   4. Count the electrons for each electronic sublevel. Count the elements in each column block that have not been crossed out by attaching one electron from each element, and write their number next to the block symbol for each column block as follows: 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 4d 10 4f 12 5s 2 5p 6 6s 2 . In our example, this is the electronic configuration of erbium.

   5. Be aware of incorrect electronic configurations. There are eighteen typical exceptions related to the electronic configurations of atoms in the lowest energy state, also called the ground energy state. They do not obey the general rule only in the last two or three positions occupied by electrons. In this case, the actual electronic configuration assumes that the electrons are in a state of lower energy compared to the standard configuration of the atom. Exception atoms include:

      • Cr(..., 3d5, 4s1); Cu(..., 3d10, 4s1); Nb(..., 4d4, 5s1); Mo(..., 4d5, 5s1); Ru(..., 4d7, 5s1); Rh(..., 4d8, 5s1); Pd(..., 4d10, 5s0); Ag(..., 4d10, 5s1); La(..., 5d1, 6s2); Ce(..., 4f1, 5d1, 6s2); Gd(..., 4f7, 5d1, 6s2); Au(..., 5d10, 6s1); AC(..., 6d1, 7s2); Th(..., 6d2, 7s2); Pa(..., 5f2, 6d1, 7s2); U(..., 5f3, 6d1, 7s2); Np(..., 5f4, 6d1, 7s2) and cm(..., 5f7, 6d1, 7s2).
   • To find the atomic number of an atom when it is written in electronic form, simply add up all the numbers that follow the letters (s, p, d, and f). This only works for neutral atoms, if you're dealing with an ion it won't work - you'll have to add or subtract the number of extra or lost electrons.
   • The number following the letter is a superscript, do not make a mistake in the control.
   • The "stability of a half-filled" sublevel does not exist. This is a simplification. Any stability that pertains to "half-full" sublevels is due to the fact that each orbital is occupied by one electron, so repulsion between electrons is minimized.
   • Each atom tends to a stable state, and the most stable configurations have filled sublevels s and p (s2 and p6). Noble gases have this configuration, so they rarely react and are located on the right in the periodic table. Therefore, if a configuration ends in 3p 4 , then it needs two electrons to reach a stable state (it takes more energy to lose six, including s-level electrons, so four is easier to lose). And if the configuration ends in 4d 3 , then it needs to lose three electrons to reach a stable state. In addition, half-filled sublevels (s1, p3, d5..) are more stable than, for example, p4 or p2; however, s2 and p6 will be even more stable.
   • When you're dealing with an ion, that means the number of protons is not the same as the number of electrons. The charge of the atom in this case will be shown at the top right (usually) of the chemical symbol. Therefore, an antimony atom with a charge of +2 has the electronic configuration 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 1 . Note that 5p 3 has changed to 5p 1 . Be careful when the configuration of a neutral atom ends at sublevels other than s and p. When you take electrons, you can only take them from valence orbitals (s and p orbitals). Therefore, if the configuration ends with 4s 2 3d 7 and the atom gets +2 charge, then the configuration will end with 4s 0 3d 7 . Please note that 3d 7 Not changes, instead electrons of the s-orbital are lost.
   • There are conditions when an electron is forced to "move to a higher energy level." When a sublevel lacks one electron to be half or full, take one electron from the nearest s or p sublevel and move it to the sublevel that needs an electron.
   • There are two options for writing an electronic configuration. They can be written in ascending order of the numbers of energy levels or in the order in which the electron orbitals are filled, as was shown above for erbium.
   • You can also write the electronic configuration of an element by writing only the valence configuration, which is the last s and p sublevel. Thus, the valence configuration of antimony will be 5s 2 5p 3 .
   • Ions are not the same. It's much more difficult with them. Skip two levels and follow the same pattern depending on where you started and how high the number of electrons is.

The location of electrons on energy shells or levels is written using electronic formulas chemical elements. Electronic formulas or configurations help to represent the structure of an element's atom.

The structure of the atom

The atoms of all elements consist of a positively charged nucleus and negatively charged electrons that are located around the nucleus.

The electrons are at different energy levels. The farther an electron is from the nucleus, the more energy it has. The size of the energy level is determined by the size of the atomic orbit or orbital cloud. This is the space in which the electron moves.

Rice. 1. General structure atom.

Orbitals can have different geometric configurations:

• s-orbitals- spherical;
• p-, d and f-orbitals- dumbbell-shaped, lying in different planes.

At the first energy level of any atom, there is always an s-orbital with two electrons (an exception is hydrogen). Starting from the second level, the s- and p-orbitals are at the same level.

Rice. 2. s-, p-, d and f-orbitals.

Orbitals exist regardless of the location of electrons on them and can be filled or vacant.

Formula entry

Electronic configurations of atoms of chemical elements are written according to the following principles:

• each energy level corresponds to a serial number, denoted by an Arabic numeral;
• the number is followed by a letter denoting the orbital;
• a superscript is written above the letter, corresponding to the number of electrons in the orbital.

Recording examples:

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% electron 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 ordinal number of the chemical element assigned to it in 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 content in nature. Chemical properties The isotopes of most chemical elements are exactly the same. However, hydrogen isotopes differ greatly in properties due to a sharp fold increase in their relative atomic mass; they have even been given individual names and chemical symbols.

Elements of the first period

Scheme electronic structure 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 outer 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 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 serial numbers of chemical elements), atomic orbitals are populated in such a way that the appearance of electrons in orbitals with more high energy depends only on the principal quantum number n and does not depend on all other quantum numbers, including 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

An atom is the smallest particle of matter, consisting of a nucleus and electrons. The structure of the electron shells of atoms is determined by the position of the element in the Periodic system of chemical elements of D. I. Mendeleev.

Electron and electron shell of an atom

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

Rice. 1. Atom.

An electron is a completely unique particle that has a dual nature: it has both the properties of a wave and a particle. They are constantly moving around the nucleus.

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

Rice. 2. Shapes of electronic orbitals.

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

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

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

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

The structure of the electron shells of atoms

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

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

s-Elements elements are called, in the atoms of which the last electron enters the s-sublevel. The p-elements,d-elements andf-elements.

The beginning of each period corresponds to the opening of a new electronic layer. The period number is equal to the number of the opened electron layer. Each period, except the first, ends with the filling of the p-sublevel of the layer opened at the beginning of this period. The first period contains only s-elements (two). In the fourth and fifth periods between the s- (two) and p-elements (six) are d-elements (ten). In the sixth and seventh, after a pair of s-elements, there is (in violation of Klechkovsky's rules) one d-element, then fourteen f-elements (they are placed separate lines at the bottom of the table - lanthanides and actinides), then nine d-elements and, as always, the periods end with six p-elements.

Vertically, the table is divided into 8 groups, each group - into the main and secondary subgroups. In the main subgroups are s- and p-elements, in the secondary - d-elements. The main subgroup is easy to determine - it contains elements of 1-3 periods. Strictly below them are the remaining elements of the main subgroup. The elements of the secondary subgroup are located to the side (left or right).

Valence of atoms

In the classical view, valence is determined by the number of unpaired electrons in the ground or excited state of atoms. Basic state- electronic state of an atom in which its energy is minimal. excited state- the electronic state of an atom, corresponding to the transition of one or more electrons from an orbital with a lower energy to a free orbital with a higher energy. For s- and p-elements, the transition of electrons is possible only within the outer electronic layer. For d-elements, transitions are possible within the d-sublevel of the pre-outer layer and the s- and p-sublevels of the outer layer. For f-elements, transitions are possible within (n-2)f-, (n-1)d-, ns- and np-sublevels, where n is the number of the outer electronic layer. Valence electrons called electrons that determine the valency of an atom in its ground or excited state. Valence electron layer- the layer on which the valence electrons are located.

Describe using quantum numbers the electrons of the outer layer of the sulfur atom and the valence electrons of iron (ground state). Indicate the possible valencies and oxidation states of the atoms of these elements.

1). Sulfur atom.

Sulfur has serial number 16. It is in the third period, the sixth group, the main subgroup. Therefore, this is a p-element, the outer electronic layer is the third, and it is valence. It has six electrons. The electronic structure of the valence layer has the form

   

For all electrons n=3, since they are located on the third layer. Let's consider them in order:

 n=3, L=0 (the electron is located in the s-orbital), m l =0 (for the s-orbital, only such a value of the magnetic quantum number is possible), m s =+1/2 (rotation around its own axis occurs clockwise) ;

 n=3, L=0, m l \u003d 0 (these three quantum numbers are the same as those of the first electron, since both electrons are in the same orbital), m s \u003d -1/2 (only here the difference appears, required by the Pauli principle);

 n=3, L=1 (this is a p-electron), m l \u003d +1 (out of three possible values ​​m l \u003d 1, 0 for the first p-orbital, we choose the maximum, this is the p x-orbital), m s \u003d +1 / 2;

 n=3, L=1, m l = +1, m s =-1/2;

 n=3, L=1, m l \u003d 0 (this is r y-orbital), m s \u003d +1/2;

 n=3, L=1, m l \u003d -1 (this is the p z-orbital), m s \u003d +1/2.

Consider the valencies and oxidation states of sulfur. On the valence layer in the ground state of the atom, there are two electron pairs, two unpaired electrons and five free orbitals. Therefore, the valency of sulfur in this state is II. Sulfur is a non-metal. Before the completion of the layer, it lacks two electrons, therefore, in compounds with atoms of less electronegative elements, for example, with metals, it can exhibit a minimum oxidation state of -2. Depairing of electron pairs is possible, because there are free orbitals on this layer. Therefore, in the first excited state (S*)

In compounds with atoms of more electronegative elements, such as oxygen, all six valence electrons can be displaced from sulfur atoms, so its maximum oxidation state is +6.

2). Iron.

The serial number of iron is 26. It is located in the fourth period, in the eighth group, a side subgroup. This is a d-element, the sixth in a series of d-elements of the fourth period. The valence electrons of iron (eight) are located on the 3d-sublevel (six, in accordance with the position in the row of d-elements) and on the 4s-sublevel (two):

    

Let's consider them in order:

 n=3, L=2, m l = +2, m s = +1/2;

 n=3, L=2, m l = +2, m s = -1/2;

 n=3, L=2, m l = +1, m s = +1/2;

 n=3, L=2, m l = 0, m s = +1/2;

 n=3, L=2, m l = -1, m s = +1/2;

 n=3, L=2, m l = -2, m s = +1/2;

 n=4, L=0, m l = 0, m s = +1/2;

 n=4, L=0, m l = 0, m s = -1/2.

Valence

There are no unpaired electrons on the outer layer, therefore, the minimum valency of iron (II) appears in the excited state of the atom:

After the electrons of the outer layer are used, 4 unpaired electrons of the 3d sublevel can be involved in the formation of chemical bonds. Therefore, the maximum valency of iron is VI.

Oxidation state

Iron is a metal, therefore, it is characterized by positive oxidation states from +2 (4s-sublevel electrons are involved) to +6 (4s- and all unpaired 3d-electrons are involved).