The structure of the electronic shell of ruthenium. Catalog of files on chemistry

Electronic configuration an atom is a numerical representation of its electron orbitals. Electron orbitals are regions various shapes, located around the atomic nucleus, in which it is mathematically probable that an electron will be found. Electron configuration helps you 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 drawing up 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 your atom's symbol on the periodic table. Atomic number is a whole positive number, starting from 1 (for hydrogen) and increasing by one for each subsequent atom. Atomic number is the number of protons in an atom, and therefore it is also the number of electrons of an atom with zero charge.

Determine the charge of an atom. Neutral atoms will have the same number of electrons as shown on the periodic table. However, charged atoms will have more or less 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 each negative charge and subtract one for each positive charge.

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

1. Remember the basic list of orbitals. As the number of electrons in an atom increases, they fill the various sublevels of the atom's electron shell according to a specific sequence. Each sublevel of the electron shell, when filled, contains even number electrons. The following sublevels are available:

   Understand electronic configuration notation. Electron configurations are written to clearly show 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 takes 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 a neutral neon atom (neon's atomic number is 10).

2. Remember the order of the orbitals. Keep in mind that electron orbitals are numbered in order of increasing electron shell number, but arranged in increasing order of energy. For example, a filled 4s 2 orbital has lower energy (or less mobility) than a partially filled or filled 3d 10 orbital, so the 4s orbital is written first. Once you know the order of the orbitals, you can easily fill them according to the number of electrons in the atom. The order of filling the orbitals 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 next view: 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 14 5d 10 6p 6 7s 2 5f 14 6d 10 7p 6
   • Note that the above entry, when all orbitals are filled, is the electron configuration of element Uuo (ununoctium) 118, the highest numbered atom in the periodic table. Therefore, this electronic configuration contains all the currently known electronic sublevels of a neutrally charged atom.

3. Fill the orbitals according to the number of electrons in your atom. For example, if we want to write down the electron 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 the orbitals according to the order above until you reach the twentieth electron. The first 1s orbital will have two electrons, the 2s orbital will also have two, the 2p will have six, the 3s will have two, the 3p will have 6, and the 4s 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 arranged in order of increasing energy. For example, when you are ready to move to the 4th energy level, first write down the 4s orbital, and then 3d. After the fourth energy level, you move 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've probably already noticed that the shape of the periodic table corresponds to the order of the electron sublevels in the electron configurations. For example, the atoms in the second column from the left always end in "s 2", and the atoms on the right edge of the thin middle part always end in "d 10", etc. Use the periodic table as a visual guide to writing configurations - how the order in which you add to the orbitals corresponds to your position in the table. See below:

   • Specifically, the leftmost two 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 the bottom half contains atoms that 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 p orbital block of the periodic table. Therefore, its electronic configuration will end with. ..3p 5
   • Note that elements in the d and f orbital region of the table are characterized by 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 a 4f orbital, despite being in the 6th period.

5. Learn abbreviations for writing long electron configurations. The atoms on the right edge 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 the chemical symbol of the nearest noble gas with fewer electrons than your atom in square brackets, and then continue writing the electron 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 abbreviation that includes the noble gas. The complete configuration of zinc 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 electron configuration of argon, a noble gas. Simply replace part of the electronic configuration for zinc with the chemical symbol for argon in square brackets (.)
   • So, the electronic configuration of zinc, written in abbreviated form, has the form: 4s 2 3d 10 .
   • Please note that if you are writing the electronic configuration of a noble gas, say argon, you cannot write it! One must use the abbreviation for the noble gas preceding this element; for argon it will be neon ().

   Using the periodic table ADOMAH

   1. Master the periodic table ADOMAH. This method of recording the electronic configuration does not require memorization, but 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 periodic table ADOMAH - a special type of periodic table developed by scientist Valery Zimmerman. It is easy to find with a short internet search.

      • IN periodic table ADOMAH 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 towards hydrogen because 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 1 to 120. These numbers are ordinary atomic numbers that represent total electrons in a neutral atom.

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

      • Note the numbers 1 through 8 at the bottom of the table. These are numbers of electronic levels, or numbers of columns. Ignore columns that contain only crossed out items. For erbium, columns numbered 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 base, ignore the diagonal lines between the blocks and break the columns into column blocks, listing them in order from bottom to top. Again, ignore blocks that have all the elements crossed out. Write 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 electron configuration of Er is written in ascending order of electron sublevel number. It can also be written in order of filling the orbitals. To do this, follow the cascades from bottom to top, rather than 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 electron sublevel. Count the elements in each column block that have not been crossed out, attaching one electron from each element, and write their number next to the block symbol for each column block thus: 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 state 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 with a 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 electron configuration 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 test.
   • There is no "half-full" sublevel stability. This is a simplification. Any stability that is attributed to "half-filled" sublevels is due to the fact that each orbital is occupied by one electron, thus minimizing repulsion between electrons.
   • Each atom tends to a stable state, and the most stable configurations have the s and p sublevels filled (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 (to lose six, including the s-sublevel electrons, requires more energy, so losing four is easier). And if the configuration ends in 4d 3, then to achieve a stable state it needs to lose three electrons. 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 are dealing with an ion, this means that the number of protons is not equal to the number of electrons. The charge of the atom in this case will be depicted at the top right (usually) of the chemical symbol. Therefore, an antimony atom with charge +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 neutral atom configuration ends in sublevels other than s and p. When you take away electrons, you can only take them from the valence orbitals (s and p orbitals). Therefore, if the configuration ends with 4s 2 3d 7 and the atom receives a charge of +2, then the configuration will end with 4s 0 3d 7. Please note that 3d 7 Not changes, electrons from the s orbital are lost instead.
   • There are conditions when an electron is forced to "move to a higher energy level." When a sublevel is one electron short of being half or full, take one electron from the nearest s or p sublevel and move it to the sublevel that needs the electron.
   • There are two options for recording the electronic configuration. They can be written in increasing order of energy level numbers or in the order of filling electron orbitals, as was shown above for erbium.
   • You can also write the electronic configuration of an element by writing only the valence configuration, which represents 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 large the number of electrons is.

Composition of the atom.

An atom is made up of atomic nucleus And electron shell.

The nucleus of an atom consists of protons ( p+) and neutrons ( n 0). Most hydrogen atoms have a nucleus consisting of one proton.

Number of protons N(p+) is equal to the nuclear charge ( Z) and the ordinal number of the element in the natural series of elements (and in the periodic table of elements).

N(p +) = Z

Sum of neutrons N(n 0), denoted simply by the letter N, and number of protons Z called mass number and is designated by the letter A.

A = Z + N

The electron shell of an atom consists of electrons moving around the nucleus ( e -).

Number of electrons N(e-) in the electron shell of a neutral atom is equal to the number of protons Z at its core.

The mass of a proton is approximately equal to the mass of a neutron and 1840 times the mass of an electron, so the mass of an atom is almost equal to the mass of the nucleus.

The shape of the atom is spherical. The radius of the nucleus is approximately 100,000 times smaller than the radius of the atom.

Chemical element- type of atoms (collection of atoms) with the same nuclear charge (with the same number of protons in the nucleus).

Isotope- a collection of atoms of the same element with the same number of neutrons in the nucleus (or a type of atom with the same number of protons and the same number of neutrons in the nucleus).

Different isotopes differ from each other in the number of neutrons in the nuclei of their atoms.

Designation of an individual atom or isotope: (E - element symbol), for example: .

Structure of the electron shell of an atom

Atomic orbital- state of an electron in an atom. The symbol for the orbital is . Each orbital has a corresponding electron cloud.

Orbitals of real atoms in the ground (unexcited) state are of four types: s, p, d And f.

Electronic cloud- the part of space in which an electron can be found with a probability of 90 (or more) percent.

Note: sometimes the concepts of “atomic orbital” and “electron cloud” are not distinguished, calling both “atomic orbital”.

The electron shell of an atom is layered. Electronic layer formed by electron clouds of the same size. The orbitals of one layer form electronic ("energy") level, their energies are the same for the hydrogen atom, but different for other atoms.

Orbitals of the same type are grouped into electronic (energy) sublevels:
s-sublevel (consists of one s-orbitals), symbol - .
p-sublevel (consists of three p
d-sublevel (consists of five d-orbitals), symbol - .
f-sublevel (consists of seven f-orbitals), symbol - .

The energies of orbitals of the same sublevel are the same.

When designating sublevels, the number of the layer (electronic level) is added to the sublevel symbol, for example: 2 s, 3p, 5d means s-sublevel of the second level, p-sublevel of the third level, d-sublevel of the fifth level.

The total number of sublevels at one level is equal to the level number n. The total number of orbitals at one level is equal to n 2. Accordingly, total number clouds in one layer is also equal n 2 .

Designations: - free orbital (without electrons), - orbital with an unpaired electron, - orbital with an electron pair (with two electrons).

The order in which electrons fill the orbitals of an atom is determined by three laws of nature (the formulations are given in simplified terms):

1. The principle of least energy - electrons fill the orbitals in order of increasing energy of the orbitals.

2. The Pauli principle - there cannot be more than two electrons in one orbital.

3. Hund's rule - within a sublevel, electrons first fill empty orbitals (one at a time), and only after that they form electron pairs.

The total number of electrons in the electronic level (or electron layer) is 2 n 2 .

The distribution of sublevels by energy is expressed as follows (in order of increasing energy):

1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p ...

This sequence is clearly expressed by an energy diagram:

The distribution of an atom's electrons across levels, sublevels, and orbitals (electronic configuration of an atom) can be depicted as an electron formula, an energy diagram, or, more simply, as a diagram of electron layers ("electron diagram").

Examples of the electronic structure of atoms:

Valence electrons- electrons of the atom that can take part in the formation chemical bonds. For any atom, these are all the outer electrons plus those pre-outer electrons whose energy is greater than that of the outer ones. For example: the Ca atom has 4 outer electrons s 2, they are also valence; the Fe atom has 4 outer electrons s 2 but he has 3 d 6, therefore the iron atom has 8 valence electrons. Valence electronic formula of the calcium atom is 4 s 2, and iron atoms - 4 s 2 3d 6 .

Periodic table chemical elements D. I. Mendeleev
(natural system of chemical elements)

Periodic law chemical elements(modern formulation): properties of chemical elements, as well as simple and complex substances, formed by them, are periodically dependent on the value of the charge from atomic nuclei.

Periodic table- graphic expression of the periodic law.

Natural series of chemical elements- a series of chemical elements arranged according to the increasing number of protons in the nuclei of their atoms, or, what is the same, according to the increasing charges of the nuclei of these atoms. The serial number of the element in this row equal to the number protons in the nucleus of any atom of that element.

The table of chemical elements is constructed by "cutting" the natural series of chemical elements into periods(horizontal rows of the table) and groupings (vertical columns of the table) of elements with a similar electronic structure of atoms.

Depending on the way you combine elements into groups, the table may be long-period(elements with the same number and type of valence electrons are collected in groups) and short period(elements with the same number of valence electrons are collected in groups).

The short-period table groups are divided into subgroups ( main And side), coinciding with the groups of the long-period table.

All atoms of elements of the same period have the same number of electron layers, equal to the period number.

Number of elements in periods: 2, 8, 8, 18, 18, 32, 32. Most of the elements of the eighth period were obtained artificially; the last elements of this period have not yet been synthesized. All periods except the first begin with an element forming alkali metal(Li, Na, K, etc.) and end with a noble gas forming element (He, Ne, Ar, Kr, etc.).

In the short-period table there are eight groups, each of which is divided into two subgroups (main and secondary), in the long-period table there are sixteen groups, which are numbered in Roman numerals with the letters A or B, for example: IA, IIIB, VIA, VIIB. Group IA of the long-period table corresponds to the main subgroup of the first group of the short-period table; group VIIB - secondary subgroup of the seventh group: the rest - similarly.

The characteristics of chemical elements naturally change in groups and periods.

In periods (with increasing serial number)

• nuclear charge increases
• the number of outer electrons increases,
• the radius of atoms decreases,
• the strength of the bond between electrons and the nucleus increases (ionization energy),
• electronegativity increases,
• oxidizing properties are enhanced simple substances("non-metallicity"),
• the reducing properties of simple substances weaken ("metallicity"),
• weakens the basic character of hydroxides and corresponding oxides,
• the acidic character of hydroxides and corresponding oxides increases.

In groups (with increasing serial number)

• nuclear charge increases
• the radius of atoms increases (only in A-groups),
• the strength of the bond between electrons and the nucleus decreases (ionization energy; only in A-groups),
• electronegativity decreases (only in A-groups),
• the oxidizing properties of simple substances weaken ("non-metallicity"; only in A-groups),
• the reducing properties of simple substances are enhanced ("metallicity"; only in A-groups),
• the basic character of hydroxides and corresponding oxides increases (only in A-groups),
• weakens the acidic character of hydroxides and corresponding oxides (only in A-groups),
• the stability of hydrogen compounds decreases (their reducing activity increases; only in A-groups).

Tasks and tests on the topic "Topic 9. "Structure of the atom. Periodic law and periodic system of chemical elements by D. I. Mendeleev (PSHE) "."

• Periodic law - Periodic law and structure of atoms grades 8–9
  You must know: the laws of filling orbitals with electrons (the principle of least energy, the Pauli principle, Hund's rule), the structure of the periodic table of elements.

  You must be able to: determine the composition of an atom by the position of the element in the periodic table, and, conversely, find an element in the periodic system, knowing its composition; depict the structure diagram, electronic configuration of an atom, ion, and, conversely, determine the position of a chemical element in the PSCE from the diagram and electronic configuration; characterize the element and the substances it forms according to its position in the PSCE; determine changes in the radius of atoms, properties of chemical elements and the substances they form within one period and one main subgroup of the periodic system.

  Example 1. Determine the number of orbitals in the third electron level. What are these orbitals?
  To determine the number of orbitals, we use the formula N orbitals = n 2 where n- level number. N orbitals = 3 2 = 9. One 3 s-, three 3 p- and five 3 d-orbitals.

  Example 2. Determine which element's atom has electronic formula 1 s 2 2s 2 2p 6 3s 2 3p 1 .
  In order to determine what element it is, you need to find out its atomic number, which is equal to the total number of electrons of the atom. In this case: 2 + 2 + 6 + 2 + 1 = 13. This is aluminum.

  After making sure that everything you need has been learned, proceed to completing the tasks. We wish you success.

  
  Recommended reading:
  • O. S. Gabrielyan and others. Chemistry 11th grade. M., Bustard, 2002;
  • G. E. Rudzitis, F. G. Feldman. Chemistry 11th grade. M., Education, 2001.

Let's find out how to create the electronic formula of a chemical element. This question is important and relevant, as it gives an idea not only of the structure, but also of the expected physical and chemical properties of the atom in question.

Compilation rules

In order to compose a graphical and electronic formula of a chemical element, it is necessary to have an understanding of the theory of atomic structure. To begin with, there are two main components of an atom: the nucleus and the negative electrons. The nucleus includes neutrons, which have no charge, as well as protons, which have a positive charge.

Discussing how to compose and determine the electronic formula of a chemical element, we note that to find the number of protons in the nucleus, the Mendeleev periodic system will be required.

The atomic number of an element corresponds to the number of protons present in its nucleus. The number of the period in which the atom is located characterizes the number of energy layers on which electrons are located.

To determine the number of neutrons devoid of electric charge, it is necessary to subtract its atomic number (number of protons) from the relative mass of an element’s atom.

Instructions

In order to understand how to compose the electronic formula of a chemical element, consider the filling rule negative particles sublevels, formulated by Klechkovsky.

Depending on how much free energy the free orbitals have, a series is compiled that characterizes the sequence of filling levels with electrons.

Each orbital contains only two electrons, which are arranged in antiparallel spins.

In order to express the structure of electronic shells, graphic formulas are used. What do the electronic formulas of atoms of chemical elements look like? How to create graphic options? These questions are included in school course chemistry, so let’s look at them in more detail.

There is a certain matrix (basis) that is used when compiling graphic formulas. The s-orbital is characterized by only one quantum cell, in which two electrons are located opposite each other. They are indicated graphically by arrows. For the p-orbital, three cells are depicted, each also containing two electrons, the d orbital contains ten electrons, and the f orbital is filled with fourteen electrons.

Examples of compiling electronic formulas

Let's continue the conversation about how to compose the electronic formula of a chemical element. For example, you need to create a graphical and electronic formula for the element manganese. First let's determine the position of this element in the periodic table. It has atomic number 25, therefore, there are 25 electrons in the atom. Manganese is a fourth period element and therefore has four energy levels.

How to write the electronic formula of a chemical element? We write down the sign of the element, as well as its serial number. Using Klechkovsky’s rule, we distribute electrons among energy levels and sublevels. We place them sequentially on the first, second, and third levels, placing two electrons in each cell.

Next, we sum them up, getting 20 pieces. Three levels are completely filled with electrons, and only five electrons remain on the fourth. Considering that each type of orbital has its own energy reserve, we distribute the remaining electrons into the 4s and 3d sublevels. As a result, the finished electronic graphic formula for the manganese atom has the following form:

1s2 / 2s2, 2p6 / 3s2, 3p6 / 4s2, 3d3

Practical significance

Using electron graphic formulas, you can clearly see the number of free (unpaired) electrons that determine the valence of a given chemical element.

We offer a generalized algorithm of actions with which you can create electron graphic formulas for any atoms located in the periodic table.

First of all, it is necessary to determine the number of electrons using the periodic table. The period number indicates the number of energy levels.

Belonging to a certain group is associated with the number of electrons located in the outer energy level. The levels are divided into sublevels and filled in taking into account the Klechkovsky rule.

Conclusion

In order to determine the valence possibilities of any chemical element located in the periodic table, it is necessary to compile an electronic graphic formula of its atom. The algorithm given above will allow you to cope with the task, determine possible chemical and physical properties atom.

Algorithm for composing 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. Based on 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 contains a maximum of 2 electrons 1s 2, 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.
• First to fill s- sublevel, then р-, d- b f- sublevels.
• Electrons fill the orbitals in order of increasing energy of the orbitals (Klechkovsky's rule).
• Within a sublevel, electrons first occupy free orbitals one by one, and only after that they form pairs (Hund’s rule).
• There cannot be more than two electrons in one orbital (Pauli principle).

Examples.

1. Let's create an electronic formula for nitrogen. Nitrogen is number 7 on the periodic table.

2. Let's create the electronic formula for argon. Argon is number 18 on the periodic table.

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

3. Let's create the electronic formula of chromium. Chromium is number 24 on the periodic table.

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

Energy diagram of zinc.

4. Let's create an electronic formula for zinc. Zinc is number 30 on the periodic table.

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

Please 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:

Electrons

The concept of atom arose in the ancient world to designate particles of matter. Translated from Greek, atom means “indivisible.”

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

State of electrons in an atom

The state of an electron in an atom is understood as 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. we can only talk about the probability of finding it in the space around the nucleus.

It can be located in any part of this space surrounding the core, and the totality of it various provisions considered as an electron cloud with a certain negative charge density. Figuratively, this can be imagined this way: if it were possible to photograph the position of an electron in an atom after hundredths or millionths of a second, as in a photo finish, then the electron in such photographs would be represented as dots. If countless such photographs were superimposed, the picture would be of an electron cloud with the greatest density where there would be the most of these points.

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

The most important characteristic of the motion of an electron in a certain orbital is 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.

The integer n, indicating the number of the energy level, is called the principal quantum number. It characterizes the energy of electrons occupying a given energy level. Electrons of the first energy level, closest to the nucleus, have the lowest energy. Compared to electrons of the first level, electrons of subsequent levels will be characterized by a large supply of energy. Consequently, the electrons of the outer level are least tightly bound to the atomic nucleus.

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

N = 2n 2 ,

where N is the maximum number of electrons; n is the level number, or the main quantum number. Consequently, at the first energy level closest to the nucleus there can be 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 divided into sublevels (sublayers), slightly different 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. The sublevels, in turn, are formed by orbitals. Each valuen corresponds to the number of orbitals equal to n.

Sublevels are usually designated with Latin letters, as well as the shape of the orbitals of which they are composed: s, p, d, f.

Protons and Neutrons

An atom of any chemical element is comparable to a tiny solar system. Therefore, this 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 taken as one in chemistry). Neutrons carry no charge, they are neutral and have a mass equal to the mass of a proton.

Protons and neutrons together are 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 is:

13 + 14 = 27

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

Since the mass of the electron, which is negligibly small, can be neglected, it is obvious that the entire mass of the atom is concentrated in the nucleus. Electrons are designated e - .

Since the atom electrically neutral, then it is also obvious that the number of protons and electrons in an atom is the same. It is equal to the serial number of the chemical element assigned to it in Periodic table. The mass of an atom consists of the mass of protons and neutrons. Knowing the atomic number of the element (Z), i.e. the number of protons, and the mass number (A), equal to the sum 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. Thus, carbon has three isotopes with masses 12, 13, 14; oxygen - three isotopes with masses 16, 17, 18, etc. The relative atomic mass of a chemical element usually given in the Periodic Table is the average value of the atomic masses of a natural mixture of isotopes of a given element, taking into account their relative abundance in nature. Chemical properties The isotopes of most chemical elements are exactly the same. However, hydrogen isotopes differ greatly in properties due to a sharp multiple increase in their relative atomic mass; they are even given individual names and chemical symbols.

Elements of the first period

Diagram of the electronic structure of the hydrogen atom:

Diagrams of the electronic structure of atoms show the distribution of electrons across electronic layers (energy levels).

Graphic electronic formula of the hydrogen atom (shows the distribution of electrons by energy levels and sublevels):

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

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

For all elements of the second period the first electronic layer is filled, and 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 Pauli and Hund rules.

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

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

The magnesium atom completes its 3s electron orbital. Na and Mg are s-elements.

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

Elements of the third period have unfilled 3d orbitals.

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

Elements of the fourth - seventh periods

A fourth electron layer appears in potassium and calcium atoms, and the 4s sublevel is filled, since it has lower 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 secondary subgroups, their outermost electronic layer is filled, and they are classified as transition elements.

Pay attention to the structure of the electronic shells of chromium and copper atoms. In them, one electron “fails” from the 4s to the 3d sublevel, 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 complete - all sublevels 3s, 3p and 3d are filled in it, with a total of 18 electrons. In the elements following zinc, the fourth electron layer, the 4p sublevel, continues to be filled.

Elements from Ga to Kr are p-elements.

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

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

4f elements are called lanthanides.

5f elements are called actinides.

The order of filling 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 here, too, there are elements in which the order of filling the electron orbitals is “violated,” which, for example, is associated with the greater energy stability of half and fully 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 electron 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 groups III-VIII.

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

• f-elements. The f-sublevel of the third outer level of the atom is filled with electrons; these include 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 as “spindle”), i.e., having such properties that conditionally can be 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, i.e. 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 electronic shells of atoms is depicted using energy or quantum cells - so-called graphical electronic formulas are written. For this notation, the following notation is used: each quantum cell is designated by a cell that corresponds to one orbital; Each electron is indicated by an arrow corresponding to the spin direction. When writing a graphical electronic formula, you should remember two rules: Pauli's principle and F. Hund's rule, according to which electrons occupy free cells first one at a time and have the same spin value, and only then pair, but the spins, according to the Pauli principle, will already be in opposite directions.

Hund's rule and Pauli's principle

Hund's rule- a rule of quantum chemistry that 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 a given sublayer must 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, one orbital contains two electrons with half-integer spins opposite sign, which pair (form a two-electron cloud) and, as a result, the total spin of the orbital becomes equal to zero.

Another wording: Lower 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 us analyze this rule using the example of filling p-sublevel orbitals 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 spin orientation).

Klechkovsky's rule

Klechkovsky's rule - as the total number of electrons in atoms increases (as the charges of their nuclei increase, or serial numbers chemical elements) atomic orbitals are populated in such a way that the appearance of electrons in an orbital with a higher energy depends only on the main 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 distance of the electron charge density from the nucleus and does not depend on the characteristics of its motion in the field of the nucleus.

The empirical Klechkovsky rule and the ordering scheme that follows from it are somewhat contradictory to the real energy sequence of atomic orbitals only in two similar cases: 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 is replaced by the d-sublevel of the previous layer, which leads to an energetically more stable state of the atom, namely: after filling orbital 6 with two electrons s