What is an electronic level in chemistry. Energy levels and sublevels. Distribution of electrons using the periodic system of D. I. Mendeleev

More strictly speaking, the relative arrangement of sublevels is determined not so much by their greater or lesser energy, but by the requirement of a minimum of the total energy of the atom.

The distribution of electrons in atomic orbitals occurs, starting from the orbital with the lowest energy (principle of minimum energy), those. The electron enters the nearest orbital to the nucleus. This means that first those sublevels are filled with electrons for which the sum of the values ​​of quantum numbers ( n+l) was minimal. Thus, the energy of an electron at the 4s sublevel is less than the energy of an electron located at the 3d sublevel. Consequently, the filling of sublevels with electrons occurs in the following order: 1s< 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s < 5d ~ 4f < 6p < 7s < 6d ~ 5f < 7p.

Based on this requirement, the minimum energy is reached for most atoms when their sublevels are filled in the sequence shown above. But there are exceptions that you can find in the tables "Electronic configurations of elements", however, these exceptions rarely need to be taken into account when considering chemical properties elements.

Atom chrome has an electronic configuration not 4s 2 3d 4 , but 4s 1 3d 5 . This is an example of how the stabilization of states with parallel spins of electrons prevails over the insignificant difference in the energy states of the 3d and 4s sublevels (Hund's rules), that is, the energetically favorable states for the d-sublevel are d5 And d10. Energy diagrams of the valence sublevels of chromium and copper atoms are shown in Fig. 2.1.1.

A similar transition of one electron from the s-sublevel to the d-sublevel occurs in 8 more elements: Cu, Nb, Mo, Ru, Ag, Pt, Au. At the atom Pd there is a transition of two s-electrons to the d-sublevel: Pd 5s 0 4d 10 .

Fig.2.1.1. Energy diagrams of valence sublevels of chromium and copper atoms

Filling rules electron shells:

1. First, find out how many electrons the atom of the element of interest to us contains. To do this, it is enough to know the charge of its nucleus, which is always equal to the serial number of the element in the Periodic Table of D.I. Mendeleev. The serial number (the number of protons in the nucleus) is exactly equal to the number of electrons in the entire atom.

2. Successively fill the orbitals, starting from the 1s orbital, with the available electrons, taking into account the principle of minimum energy. In this case, it is impossible to place more than two electrons with oppositely directed spins on each orbital (Pauli's rule).

3. We write down the electronic formula of the element.

An atom is a complex, dynamically stable microsystem of interacting particles: protons p +, neutrons n 0 and electrons e -.

Fig.2.1.2. Filling of energy levels with electrons of the element phosphorus

The electronic structure of the hydrogen atom (z = 1) can be depicted as follows:

+1 H 1s 1 , n = 1 , where the quantum cell (atomic orbital) is denoted as a line or square, and electrons as arrows.

Each atom of the subsequent chemical element in the periodic system is a multi-electron atom.

The lithium atom, like the hydrogen and helium atom, has the electronic structure of an s-element, because. the last electron of the lithium atom "sits down" on the s-sublevel:

+3 Li 1s 2 2s 1 2p 0

The first electron in the p-state appears in the boron atom:

+5 V 1s 2 2s 2 2p 1

Writing an electronic formula is easier to show on specific example. Let's say we need to find out the electronic formula of an element with serial number 7. There should be 7 electrons in an atom of such an element. Let's fill the orbitals with seven electrons, starting from the bottom 1s orbital.

So, 2 electrons will be placed in 1s orbitals, 2 more electrons in 2s orbitals, and the remaining 3 electrons can be placed in three 2p orbitals.

The electronic formula of the element with serial number 7 (this is the element nitrogen, having the symbol “N”) looks like this:

+7 N 1s 2 2s 2 2p 3

Consider the action of Hund's rule on the example of a nitrogen atom: N 1s 2 2s 2 2p 3. At the 2nd electronic level, there are three identical p-orbitals: 2px, 2py, 2pz. Electrons will populate them so that each of these p-orbitals will have one electron. This is explained by the fact that in neighboring cells, electrons repel each other less, as similarly charged particles. received by us electronic formula nitrogen carries very important information: the 2nd (external) electronic level of nitrogen is not completely filled with electrons (it has 2 + 3 = 5 valence electrons) and three electrons are missing until it is completely filled.

The outer level of an atom is the level farthest from the nucleus that contains valence electrons. It is this shell that comes into contact when it collides with the outer levels of other atoms in chemical reactions. When interacting with other atoms, nitrogen is able to accept 3 additional electrons to its outer level. In this case, the nitrogen atom will receive a completed, that is, the most filled external electronic level, on which 8 electrons will be located.

A completed level is more energetically advantageous than an incomplete one, so the nitrogen atom should easily react with any other atom that can give it 3 extra electrons to complete its outer level.

(1887-1961) to describe the state of an electron in a hydrogen atom. He combined mathematical expressions for oscillatory processes and the de Broglie equation and obtained the following linear differential homogeneous equation:

where ψ is the wave function (analogous to the amplitude for wave motion in classical mechanics), which characterizes the motion of an electron in space as a wave-like perturbation; x, y, z- coordinates, m is the rest mass of the electron, h is Planck's constant, E is the total energy of the electron, E p is the potential energy of the electron.

The solutions of the Schrödinger equation are wave functions. For a one-electron system (hydrogen atom), the expression for the potential energy of an electron has a simple form:

E p = − e 2 / r,

where e is the charge of an electron, r is the distance from the electron to the nucleus. In this case, the Schrödinger equation has an exact solution.

To solve a wave equation, we must separate its variables. For this, replace Cartesian coordinates x, y, z into spherical r, θ, φ. Then the wave function can be represented as a product of three functions, each of which contains only one variable:

ψ( x,y,z) = R(r) Θ(θ) Φ(φ)

Function R(r) is called the radial component of the wave function, and Θ(θ) Φ(φ) - its angular components.

In the course of solving the wave equation, integers are introduced - the so-called quantum numbers(the main thing n, orbital l and magnetic m l). Function R(r) depends on n And l, the function Θ(θ) - from l And m l, the function Φ(φ) - from m l .

The geometric image of the one-electron wave function is atomic orbital. It is a region of space around the nucleus of an atom, in which the probability of finding an electron is high (usually a probability value of 90-95% is chosen). This word comes from the Latin orbit"(path, track), but has a different meaning, which does not coincide with the concept of the trajectory (path) of an electron around an atom, proposed by N. Bohr for the planetary model of the atom. The contours of the atomic orbital are a graphical representation of the wave function obtained by solving the wave equation for one electron.

quantum numbers

Quantum numbers that arise when solving the wave equation serve to describe the states of a quantum chemical system. Each atomic orbital is characterized by a set of three quantum numbers: the main n, orbital l and magnetic m l .

Principal quantum number n characterizes the energy of the atomic orbital. It can take any positive integer value. The greater the value n, the higher the energy and the larger the size of the orbital. The solution of the Schrödinger equation for the hydrogen atom gives the following expression for the electron energy:

E= −2π 2 me 4 / n 2 h 2 = −1312,1 / n 2 (kJ/mol)

Thus, each value of the principal quantum number corresponds to a certain value of the electron energy. Energy levels with specific values n sometimes spelled out K, L, M, N... (for n = 1, 2, 3, 4...).

Orbital quantum number l characterizes the energy sublevel. Atomic orbitals with different orbital quantum numbers differ in energy and shape. For each n integer values ​​allowed l from 0 to ( n−1). Values l= 0, 1, 2, 3... correspond to energy sublevels s, p, d, f.

The form s-orbitals spherical, p Orbitals are like dumbbells d- And f-orbitals have a more complex shape.

Magnetic quantum number m l responsible for the orientation of atomic orbitals in space. For every value l magnetic quantum number m l can take integer values ​​from −l to +l (total 2 l+ 1 values). For example, R-orbitals ( l= 1) can be oriented in three ways ( m l = -1, 0, +1).

An electron occupying a certain orbital is characterized by three quantum numbers describing this orbital and a fourth quantum number ( spin) m s, which characterizes the electron spin - one of the properties (along with mass and charge) of this elementary particle. Spin- own magnetic moment momentum of an elementary particle. Although this word in English means " rotation", the spin is not associated with any movement of the particle, but has a quantum nature. The electron spin is characterized by the spin quantum number m s, which can be equal to +1/2 and −1/2.

Quantum numbers for an electron in an atom:

Energy levels and sublevels

The set of states of an electron in an atom with the same value n called energy level. The number of levels at which the electrons are in the ground state of the atom coincides with the number of the period in which the element is located. The numbers of these levels are indicated by numbers: 1, 2, 3, ... (less often - by letters K, L, M, ...).

Energy sublevel- a set of energy states of an electron in an atom, characterized by the same values ​​of quantum numbers n And l. Sublevels are denoted by letters: s, p, d, f... The first energy level has one sublevel, the second - two sublevels, the third - three sublevels and so on.

If the orbitals are designated in the diagram as cells (square frames), and the electrons as arrows (or ↓), then you can see that the main quantum number characterizes the energy level (EU), the combination of the main and orbital quantum numbers - the energy sublevel (EPL ), a set of principal, orbital and magnetic quantum numbers - atomic orbital, and all four quantum numbers are an electron.

Each orbital corresponds to a certain energy. The designation of the orbital includes the number of the energy level and the letter corresponding to the corresponding sublevel: 1 s, 3p, 4d etc. For each energy level, starting from the second, the existence of three equal in energy p orbitals located in three mutually perpendicular directions. At each energy level, starting from the third, there are five d-orbitals with a more complex four-leaf shape. Starting from the fourth energy level, even more complex shapes appear. f-orbitals; There are seven on each level. An atomic orbital with an electron charge distributed over it is often called an electron cloud.

electron density

The spatial distribution of the electron charge is called the electron density. Based on the fact that the probability of finding an electron in an elementary volume d V equals |ψ| 2d V, we can calculate the radial distribution function of the electron density.

If we take the volume of a spherical layer of thickness d as an elementary volume r on distance r from the nucleus of an atom

d V= 4π r 2d r,

and the function of the radial distribution of the probability of finding an electron in an atom (probability of electron density) is equal to

W r= 4π r 2 |ψ| 2d r

It represents the probability of finding an electron in a spherical layer of thickness d r at a certain distance of the layer from the nucleus of the atom.

For 1 s-orbitals, the probability of detecting an electron is maximum in the layer located at a distance of 52.9 nm from the nucleus. As you move away from the nucleus of an atom, the probability of finding an electron approaches zero. In case 2 s-orbitals, two maxima and a nodal point appear on the curve, where the probability of finding an electron is zero. In general, for an orbital characterized by quantum numbers n And l, the number of nodes on the graph of the radial probability distribution function is ( n − l − 1).

To describe the state of an electron in an atom, in addition to quantum numbers, they use:

• atomic energy level diagrams;
• electronic formulas or configurations.

Energy level chart

Rice. Energy levels and sublevels of the atom.

The figure shows the level diagram of an atom, which can be used to describe the electrons of any atom.

Energy levels atoms (electron clouds forming electron atomic layers) are indicated by numbers 1, 2, 3, 4 ...

Energy sublevels atom (energy levels characterizing the binding energy of an electron with an atomic nucleus) are denoted by the letters s, p, d, f.

Energy sublevels can be displayed as quantum cells (figure on the right): free (empty cell); partially filled (one vertical arrow pointing up or down, indicating an unpaired electron); completely filled (two vertical differently directed arrows denoting paired electrons).

Electronic formula of the atom

Everything on the energy level diagrams is quite clear and visual, but cumbersome. Using the electronic configuration, the diagram can be expressed in one short line.

Consider a carbon atom that has two energy levels with only 6 electrons (2 on the inside and 4 on the outside):

The figures below show examples of the electronic formulas of carbon and sodium atoms (model of the electron shell) and their graphic representation:

Rice. Electronic formula of carbon.

Rice. Electronic formula of sodium.

In the electronic configuration, the name of the energy level orbital is indicated in the superscript of which is the number of electrons located in this orbital.

The electron shell of an atom is formed according to the following principles:

• minimum energy principle- orbitals with the lowest energy (closest to the atomic nucleus) are filled first:
  1s; 2s; 2p; 3s; 3p; 4s (3d); 4p; 5s (4d); 5p; 6s(4f)(5d); 6p; 7s;
• Pauli principle- on one atomic orbital there can be no more than 2 electrons with opposite spins (paired electrons);
• hund rule- atomic orbitals are filled in such a way that the sum of their spins is maximum.

For example, the electronic formula for chlorine is: 1s 2 2s 2 2p 6 3s 2 3p 5 .

The serial number of chlorine in the table is 17. This means that the chlorine atom contains 17 protons and 17 electrons. That is, we need to place 17 electrons on the diagram (according to the rules).

As mentioned above, schematically, the electron is displayed as an arrow. If there are two electrons in the orbital, then they are displayed as two differently directed arrows (electrons with different spins).

• First, we fill in the lowest energy level: the 1s orbital. It has 2 electrons on it.
• The next 2 electrons occupy the 2s orbital.
• Next energy level: 2p orbital - 6 electrons.
• The next 2 electrons are the 3s orbital.
• The remaining 5 electrons are located in the 3p orbital, forming two spin pairs (the last electron does not have a pair).

So the energy level diagram for chlorine would look like this:

The attentive reader, most likely, noticed that the order of filling the electronic energy sublevels in atoms is somewhat disturbed, for example, the 4s sublevel is filled first, and only then 3d. This violation is explained Klechkovsky's rule, which says that electrons fill atomic levels (sublevels) in ascending order of the sum (n + l), if the sums of the main and orbital quantum numbers are equal, filling occurs in the order of increasing n(see Quantum-mechanical model of the structure of the atom).

• For sublevel 4s: n+l = 4+0 = 4;
• For the 3d sublevel; n+l = 3+2 = 5.

Sublevels 3d, 4p, 5s have equal sums n+l=5, so the filling goes in ascending order of the main quantum number: 3d→4p→5s.

Klechkovsky's rule has a number of exceptions, when sublevels close to each other differ slightly in energy, in which case the electron tends to occupy a sublevel with lower energy, even if it is "overlying", while the "lower" level remains unfilled. For example 5d 1 is filled before 4f.

• s-elements(14): electrons fill the s-sublevel of the outer level - hydrogen, helium + the first 2 elements of each period;
• p-elements(30): electrons fill the p-sublevel of the outer level - the last 6 elements of each period;
• d-elements(32): electrons fill the d-sublevel of the second level from the outside - elements of intercalary decades of large periods, which are between s- and p-elements;
• f-elements(28): electrons fill the f-sublevel of the third level from the outside - lanthanides and actinides.

Valence electrons

Earlier we said that an atom is a neutrally charged particle, since the number of electrons and protons in it is the same. However, the electrons in the outermost orbitals are weakly attracted by the positive protons in the nucleus of the atom. Therefore, the atoms of the elements are able to give and attach electrons.

NEED TO KNOW! The valence electrons include the outer electrons, plus those pre-outer electrons whose energy is greater than that of the outer ones.

Electronic configuration an atom is a numerical representation of it electron orbitals. Electron orbitals are regions of various shapes located around the atomic nucleus, in which it is mathematically probable that an electron will be found. The electronic configuration helps to quickly and easily tell the reader how many electron orbitals an atom has, as well as to 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 even number 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 orbitals are filled, is the electron configuration of the element Uuo (ununoctium) 118, the atom periodic system with the highest number. 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, the atoms in the second column from the left always end in "s 2 ", while the 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 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 in. ..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 electron 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 don't obey 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 are dealing with an ion, then nothing will work - you will 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). This configuration is noble gases, 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 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.

Energy sublevels - section Chemistry, Basics inorganic chemistry Orbital Quantum Number L Forma...

According to the limits of changes in the orbital quantum number from 0 to (n-1), a strictly limited number of sublevels is possible in each energy level, namely: the number of sublevels is equal to the level number.

The combination of the principal (n) and orbital (l) quantum numbers completely characterizes the energy of an electron. The energy reserve of an electron is reflected by the sum (n+l).

So, for example, the electrons of the 3d sublevel have a higher energy than the electrons of the 4s sublevel:

The order in which levels and sublevels in an atom are filled with electrons is determined by rule V.M. Klechkovsky: the filling of the electronic levels of the atom occurs sequentially in the order of increasing sum (n + 1).

In accordance with this, the real energy scale of sublevels is determined, according to which the electron shells of all atoms are built:

1s ï 2s2p ï 3s3p ï 4s3d4p ï 5s4d5p ï 6s4f5d6p ï 7s5f6d…

3. Magnetic quantum number (m l) characterizes the direction of the electron cloud (orbital) in space.

The more complex the shape of the electron cloud (i.e., the higher the value of l), the more variations in the orientation of this cloud in space and the more individual energy states of the electron exist, characterized by a certain value of the magnetic quantum number.

Mathematically m l takes integer values ​​from -1 to +1, including 0, i.e. total (21+1) values.

Let us designate each individual atomic orbital in space as an energy cell ð, then the number of such cells in sublevels will be:

Poduro-ven	Possible values ​​m l	The number of individual energy states (orbitals, cells) in the sublevel
s (l=0)		one
p (l=1)	-1, 0, +1	three
d (l=2)	-2, -1, 0, +1, +2	five
f (l=3)	-3, -2, -1, 0, +1, +2, +3	seven

For example, a spherical s-orbital is uniquely directed in space. Dumbbell-shaped orbitals of each p-sublevel are oriented along three coordinate axes

4. Spin quantum number m s characterizes the electron's own rotation around its axis and takes only two values:

p- sublevel + 1 / 2 and - 1 / 2, depending on the direction of rotation in one direction or another. According to the Pauli principle, no more than 2 electrons with oppositely directed (antiparallel) spins can be located in one orbital:

Such electrons are called paired. An unpaired electron is schematically represented by a single arrow:.

Knowing the capacity of one orbital (2 electrons) and the number of energy states in the sublevel (m s), we can determine the number of electrons in the sublevels:

You can write the result differently: s 2 p 6 d 10 f 14 .

These numbers must be well remembered for the correct writing of the electronic formulas of the atom.

So, four quantum numbers - n, l, m l , m s - completely determine the state of each electron in an atom. All electrons in an atom with the same value of n make up an energy level, with the same values n and l - energy sublevel, with the same values ​​of n, l and m l- a separate atomic orbital (quantum cell). Electrons in the same orbital have different spins.

Taking into account the values ​​of all four quantum numbers, we determine the maximum number of electrons in the energy levels (electronic layers):

Large numbers of electrons (18.32) are contained only in the deep-lying electron layers of atoms, the outer electron layer can contain from 1 (for hydrogen and alkali metals) to 8 electrons (inert gases).

It is important to remember that the filling of electron shells with electrons occurs according to principle of least energy: The sublevels with the lowest energy value are filled first, then those with higher values. This sequence corresponds to the energy scale of V.M. Klechkovsky.

The electronic structure of an atom is displayed by electronic formulas, which indicate energy levels, sublevels and the number of electrons in sublevels.

For example, the hydrogen atom 1 H has only 1 electron, which is located in the first layer from the nucleus at the s-sublevel; the electronic formula of the hydrogen atom is 1s 1.

The lithium atom 3 Li has only 3 electrons, 2 of which are in the s-sublevel of the first layer, and 1 is placed in the second layer, which also begins with the s-sublevel. The electronic formula of the lithium atom is 1s 2 2s 1.

The phosphorus atom 15 P has 15 electrons located in three electron layers. Remembering that the s-sublevel contains no more than 2 electrons, and the p-sublevel contains no more than 6, we gradually place all the electrons into sublevels and draw up the electronic formula of the phosphorus atom: 1s 2 2s 2 2p 6 3s 2 3p 3.

When compiling the electronic formula of the manganese atom 25 Mn, it is necessary to take into account the sequence of increasing sublevel energy: 1s2s2p3s3p4s3d…

We gradually distribute all 25 Mn electrons: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 5 .

The final electronic formula of the manganese atom (taking into account the distance of electrons from the nucleus) looks like this:

1s2

2s 2 2p 6

3s 2 3p 6 3d 5

4s 2

The electronic formula of manganese fully corresponds to its position in the periodic system: the number of electronic layers (energy levels) - 4 is equal to the period number; there are 2 electrons in the outer layer, the penultimate layer is not completed, which is typical for metals of secondary subgroups; the total number of mobile, valence electrons (3d 5 4s 2) - 7 is equal to the group number.

Depending on which of the energy sublevels in the atom -s-, p-, d- or f- is built up last, all chemical elements are divided into electronic families: s-elements(H, He, alkali metals, metals of the main subgroup of the 2nd group of the periodic system); p-elements(elements of the main subgroups 3, 4, 5, 6, 7, 8th groups of the periodic system); d-elements(all metals of secondary subgroups); f-elements(lanthanides and actinides).

The electronic structures of atoms are a deep theoretical justification for the structure of the periodic system, the length of periods (i.e., the number of elements in periods) follows directly from the capacitance of the electronic layers and the sequence of increasing energy of sublevels:

Each period begins with an s-element with the structure of the outer layer s 1 ( alkali metal) and ends with a p-element with the structure of the outer layer …s 2 p 6 (inert gas). I-th period contains only two s-elements (H and He), the II-nd and III-th small periods contain two s-elements and six p-elements each. In the IV and V-th large periods between s- and p-elements "wedged" in 10 d-elements - transition metals, allocated to secondary subgroups. In periods VI and VII, 14 more f-elements are added to the analogous structure, which are similar in properties to lanthanum and actinium, respectively, and isolated as subgroups of lanthanides and actinides.

When studying the electronic structures of atoms, pay attention to their graphic representation, for example:

13 Al 1s 2 2s 2 2p 6 3s 2 3p 1

both versions of the image are used: a) and b):

For the correct arrangement of electrons in orbitals, it is necessary to know Gund's rule: the electrons in the sublevel are arranged so that their total spin is maximum. In other words, the electrons first occupy all free cells of the given sublevel one by one.

For example, if it is necessary to place three p-electrons (p 3) in the p-sublevel, which always has three orbitals, then out of two options Hund's rule corresponds to the first option:

As an example, consider the graphical electronic circuit of a carbon atom:

6 C 1s 2 2s 2 2p 2

The number of unpaired electrons in an atom is a very important characteristic. According to the theory of covalent bonding, only unpaired electrons can form chemical bonds and determine the valence capabilities of an atom.

If there are free energy states (unoccupied orbitals) in the sublevel, the atom, upon excitation, “steams”, separates the paired electrons, and its valence capabilities increase:

6 C 1s 2 2s 2 2p 3

Carbon in the normal state is 2-valent, in the excited state it is 4-valent. The fluorine atom has no opportunities for excitation (because all the orbitals of the outer electron layer are occupied), therefore fluorine in its compounds is monovalent.

Example 1 What are quantum numbers? What values ​​can they take?

Solution. The motion of an electron in an atom has a probabilistic character. The circumnuclear space, in which an electron can be located with the highest probability (0.9-0.95), is called the atomic orbital (AO). Atomic orbital, like any geometric figure, is characterized by three parameters (coordinates), called quantum numbers (n, l, m l). Quantum numbers do not take any, but certain, discrete (discontinuous) values. Neighboring values ​​of quantum numbers differ by one. Quantum numbers determine the size (n), shape (l) and orientation (m l) of an atomic orbital in space. Occupying one or another atomic orbital, an electron forms an electron cloud, which for electrons of the same atom can have different shape(Fig. 1). The forms of electron clouds are similar to AO. They are also called electronic or atomic orbitals. The electron cloud is characterized by four numbers (n, l, m 1 and m 5).

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All topics in this section:

Basic laws and concepts of chemistry
branch of chemistry dealing with quantitative composition substances and quantitative ratios (mass, volume) between the reacting substances, is called stoichiometry. According to this,

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