Electronic configuration of an atom. Electronic formulas of atoms and diagrams Atomic structure chemistry online

Since during chemical reactions the nuclei of the reacting atoms remain unchanged (with the exception of radioactive transformations), the chemical properties of atoms depend on the structure of their electronic shells. Theory electronic structure of the atom built on the basis of the apparatus of quantum mechanics. Thus, the structure of atomic energy levels can be obtained on the basis of quantum mechanical calculations of the probabilities of finding electrons in the space around the atomic nucleus ( rice. 4.5).

Rice. 4.5. Scheme of dividing energy levels into sublevels

The fundamentals of the theory of the electronic structure of an atom are reduced to the following provisions: the state of each electron in an atom is characterized by four quantum numbers: the principal quantum number n = 1, 2, 3,; orbital (azimuthal) l=0,1,2, n–1; magnetic m l = –l, –1,0,1,  l; spin m s = -1/2, 1/2 .

According to Pauli principle, in the same atom there cannot be two electrons having the same set of four quantum numbers n, l, m l , m s; collections of electrons with the same principal quantum numbers n form electron layers, or energy levels of the atom, numbered from the nucleus and denoted as K, L, M, N, O, P, Q, and in the energy layer with a given value n can be no more than 2n 2 electrons. Collections of electrons with the same quantum numbers n And l, form sublevels, designated as they move away from the core as s, p, d, f.

The probabilistic determination of the position of the electron in space around the atomic nucleus corresponds to the Heisenberg uncertainty principle. According to quantum mechanical concepts, an electron in an atom does not have a specific trajectory of motion and can be located in any part of the space around the nucleus, and its various positions are considered as an electron cloud with a certain negative charge density. The space around the nucleus in which an electron is most likely to be found is called orbital. It contains about 90% of the electron cloud. Each sublevel 1s, 2s, 2p etc. corresponds to a certain number of orbitals of a certain shape. For example, 1s- And 2s- orbitals are spherical and 2p-orbitals ( 2p x , 2p y , 2p z-orbitals) are oriented in mutually perpendicular directions and have the shape of a dumbbell ( rice. 4.6).

Rice. 4.6. Shape and orientation of electron orbitals.

During chemical reactions, the atomic nucleus does not undergo changes; only the electronic shells of the atoms change, the structure of which explains many of the properties of chemical elements. Based on the theory of the electronic structure of the atom, the deep physical meaning of Mendeleev’s periodic law of chemical elements was established and the theory of chemical bonding was created.

The theoretical justification of the periodic system of chemical elements includes data on the structure of the atom, confirming the existence of a connection between the periodicity of changes in the properties of chemical elements and the periodic repetition of similar types of electronic configurations of their atoms.

In the light of the doctrine of the structure of the atom, Mendeleev’s division of all elements into seven periods becomes justified: the number of the period corresponds to the number of energy levels of atoms filled with electrons. In small periods, with an increase in the positive charge of atomic nuclei, the number of electrons at the external level increases (from 1 to 2 in the first period, and from 1 to 8 in the second and third periods), which explains the change in the properties of elements: at the beginning of the period (except for the first) there is alkali metal, then a gradual weakening of metallic properties and strengthening of non-metallic properties is observed. This pattern can be traced for elements of the second period in table 4.2.

Table 4.2.

In large periods, as the charge of the nuclei increases, the filling of levels with electrons is more difficult, which explains the more complex change in the properties of elements compared to elements of small periods.

The identical nature of the properties of chemical elements in subgroups is explained by the similar structure of the external energy level, as shown in table 4.3, illustrating the sequence of filling energy levels with electrons for subgroups of alkali metals.

Table 4.3.

The group number usually indicates the number of electrons in an atom that can participate in the formation of chemical bonds. This is the physical meaning of the group number. In four places of the periodic table, the elements are not arranged in order of increasing atomic mass: Ar And K,Co And Ni,Te And I,Th And Pa. These deviations were considered shortcomings of the periodic table of chemical elements. The doctrine of the structure of the atom explained these deviations. Experimental determination of nuclear charges showed that the arrangement of these elements corresponds to an increase in the charges of their nuclei. In addition, the experimental determination of the charges of atomic nuclei made it possible to determine the number of elements between hydrogen and uranium, as well as the number of lanthanides. Now all places in the periodic table are filled in the interval from Z=1 before Z=114, however, the periodic system is not complete, the discovery of new transuranium elements is possible.

Electronic configuration an atom is a numerical representation of its 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. Electronic configuration helps 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. The atomic number is a positive integer starting at 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 are talking 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 an even number of electrons. There are the following sublevels:

   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 be as follows: 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 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 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, which represent the total number of 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 record 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 that relate 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 for 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.

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. Using 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: