The structure of the electronic shells of atoms. Electronic structure of the atom Electronic layers of elements
Chemicals are what the world around us is made of.
The properties of each chemical substance are divided into two types: chemical, which characterize its ability to form other substances, and physical, which are objectively observed and can be considered in isolation from chemical transformations. For example, the physical properties of a substance are its state of aggregation (solid, liquid or gaseous), thermal conductivity, heat capacity, solubility in various media (water, alcohol, etc.), density, color, taste, etc.
The transformation of some chemical substances into other substances is called chemical phenomena or chemical reactions. It should be noted that there are also physical phenomena that are obviously accompanied by a change in any physical properties of a substance without its transformation into other substances. Physical phenomena, for example, include the melting of ice, freezing or evaporation of water, etc.
The fact that a chemical phenomenon is taking place during a process can be concluded by observing characteristic signs of chemical reactions, such as color changes, the formation of precipitates, the release of gas, the release of heat and (or) light.
For example, a conclusion about the occurrence of chemical reactions can be made by observing:
Formation of sediment when boiling water, called scale in everyday life;
The release of heat and light when a fire burns;
Change in color of a cut of a fresh apple in air;
Formation of gas bubbles during dough fermentation, etc.
The smallest particles of a substance that undergo virtually no changes during chemical reactions, but only connect with each other in a new way, are called atoms.
The very idea of the existence of such units of matter arose in ancient Greece in the minds of ancient philosophers, which actually explains the origin of the term “atom,” since “atomos” literally translated from Greek means “indivisible.”
However, contrary to the idea of ancient Greek philosophers, atoms are not the absolute minimum of matter, i.e. they themselves have a complex structure.
Each atom consists of so-called subatomic particles - protons, neutrons and electrons, designated respectively by the symbols p +, n o and e -. The superscript in the notation used indicates that the proton has a unit positive charge, the electron has a unit negative charge, and the neutron has no charge.
As for the qualitative structure of an atom, in each atom all protons and neutrons are concentrated in the so-called nucleus, around which the electrons form an electron shell.
The proton and neutron have almost the same masses, i.e. m p ≈ m n, and the mass of the electron is almost 2000 times less than the mass of each of them, i.e. m p /m e ≈ m n /m e ≈ 2000.
Since the fundamental property of an atom is its electrical neutrality, and the charge of one electron is equal to the charge of one proton, from this we can conclude that the number of electrons in any atom is equal to the number of protons.
For example, the table below shows the possible composition of atoms:
Type of atoms with the same nuclear charge, i.e. with the same number of protons in their nuclei is called a chemical element. Thus, from the table above we can conclude that atom1 and atom2 belong to one chemical element, and atom3 and atom4 belong to another chemical element.
Each chemical element has its own name and individual symbol, which is read in a certain way. So, for example, the simplest chemical element, the atoms of which contain only one proton in the nucleus, is called “hydrogen” and is denoted by the symbol “H”, which is read as “ash”, and a chemical element with a nuclear charge of +7 (i.e. containing 7 protons) - “nitrogen”, has the symbol “N”, which is read as “en”.
As you can see from the table above, atoms of one chemical element can differ in the number of neutrons in their nuclei.
Atoms that belong to the same chemical element, but have a different number of neutrons and, as a result, mass, are called isotopes.
For example, the chemical element hydrogen has three isotopes - 1 H, 2 H and 3 H. The indices 1, 2 and 3 above the symbol H mean the total number of neutrons and protons. Those. Knowing that hydrogen is a chemical element, which is characterized by the fact that there is one proton in the nuclei of its atoms, we can conclude that in the 1 H isotope there are no neutrons at all (1-1 = 0), in the 2 H isotope - 1 neutron (2-1=1) and in the 3 H isotope – two neutrons (3-1=2). Since, as already mentioned, the neutron and proton have the same masses, and the mass of the electron is negligibly small in comparison with them, this means that the 2 H isotope is almost twice as heavy as the 1 H isotope, and the 3 H isotope is even three times heavier . Due to such a large scatter in the masses of hydrogen isotopes, the isotopes 2 H and 3 H were even assigned separate individual names and symbols, which is not typical for any other chemical element. The 2H isotope was named deuterium and given the symbol D, and the 3H isotope was given the name tritium and given the symbol T.
If we take the mass of the proton and neutron as one, and neglect the mass of the electron, in fact, the upper left index, in addition to the total number of protons and neutrons in the atom, can be considered its mass, and therefore this index is called the mass number and is designated by the symbol A. Since the charge of the nucleus of any Protons correspond to the atom, and the charge of each proton is conventionally considered equal to +1, the number of protons in the nucleus is called the charge number (Z). By denoting the number of neutrons in an atom as N, the relationship between mass number, charge number, and number of neutrons can be expressed mathematically as:
According to modern concepts, the electron has a dual (particle-wave) nature. It has the properties of both a particle and a wave. Like a particle, an electron has mass and charge, but at the same time, the flow of electrons, like a wave, is characterized by the ability to diffraction.
To describe the state of an electron in an atom, the concepts of quantum mechanics are used, according to which the electron does not have a specific trajectory of motion and can be located at any point in space, but with different probabilities.
The region of space around the nucleus where an electron is most likely to be found is called an atomic orbital.
An atomic orbital can have different shapes, sizes, and orientations. An atomic orbital is also called an electron cloud.
Graphically, one atomic orbital is usually denoted as a square cell:
Quantum mechanics has an extremely complex mathematical apparatus, therefore, in the framework of a school chemistry course, only the consequences of quantum mechanical theory are considered.
According to these consequences, any atomic orbital and the electron located in it are completely characterized by 4 quantum numbers.
- The principal quantum number, n, determines the total energy of an electron in a given orbital. The range of values of the main quantum number is all natural numbers, i.e. n = 1,2,3,4, 5, etc.
- The orbital quantum number - l - characterizes the shape of the atomic orbital and can take any integer value from 0 to n-1, where n, recall, is the main quantum number.
Orbitals with l = 0 are called s-orbitals. s-Orbitals are spherical in shape and have no directionality in space:
Orbitals with l = 1 are called p-orbitals. These orbitals have the shape of a three-dimensional figure eight, i.e. a shape obtained by rotating a figure eight around an axis of symmetry, and outwardly resemble a dumbbell:
Orbitals with l = 2 are called d-orbitals, and with l = 3 – f-orbitals. Their structure is much more complex.
3) Magnetic quantum number – m l – determines the spatial orientation of a specific atomic orbital and expresses the projection of the orbital angular momentum onto the direction of the magnetic field. The magnetic quantum number m l corresponds to the orientation of the orbital relative to the direction of the external magnetic field strength vector and can take any integer values from –l to +l, including 0, i.e. the total number of possible values is (2l+1). So, for example, for l = 0 m l = 0 (one value), for l = 1 m l = -1, 0, +1 (three values), for l = 2 m l = -2, -1, 0, +1 , +2 (five values of magnetic quantum number), etc.
So, for example, p-orbitals, i.e. orbitals with an orbital quantum number l = 1, having the shape of a “three-dimensional figure of eight,” correspond to three values of the magnetic quantum number (-1, 0, +1), which, in turn, correspond to three directions perpendicular to each other in space.
4) The spin quantum number (or simply spin) - m s - can conventionally be considered responsible for the direction of rotation of the electron in the atom; it can take on values. Electrons with different spins are indicated by vertical arrows directed in different directions: ↓ and .
The set of all orbitals in an atom that have the same principal quantum number is called the energy level or electron shell. Any arbitrary energy level with some number n consists of n 2 orbitals.
A set of orbitals with the same values of the principal quantum number and orbital quantum number represents an energy sublevel.
Each energy level, which corresponds to the principal quantum number n, contains n sublevels. In turn, each energy sublevel with orbital quantum number l consists of (2l+1) orbitals. Thus, the s sublevel consists of one s orbital, the p sublevel consists of three p orbitals, the d sublevel consists of five d orbitals, and the f sublevel consists of seven f orbitals. Since, as already mentioned, one atomic orbital is often denoted by one square cell, the s-, p-, d- and f-sublevels can be graphically represented as follows:
Each orbital corresponds to an individual strictly defined set of three quantum numbers n, l and m l.
The distribution of electrons among orbitals is called the electron configuration.
The filling of atomic orbitals with electrons occurs in accordance with three conditions:
- Minimum energy principle: Electrons fill orbitals starting from the lowest energy sublevel. The sequence of sublevels in increasing order of their energies is as follows: 1s<2s<2p<3s<3p<4s≤3d<4p<5s≤4d<5p<6s…;
To make it easier to remember this sequence of filling out electronic sublevels, the following graphic illustration is very convenient:
- Pauli principle: Each orbital can contain no more than two electrons.
If there is one electron in an orbital, then it is called unpaired, and if there are two, then they are called an electron pair.
- Hund's rule: the most stable state of an atom is one in which, within one sublevel, the atom has the maximum possible number of unpaired electrons. This most stable state of the atom is called the ground state.
In fact, the above means that, for example, the placement of 1st, 2nd, 3rd and 4th electrons in three orbitals of the p-sublevel will be carried out as follows:
The filling of atomic orbitals from hydrogen, which has a charge number of 1, to krypton (Kr), with a charge number of 36, will be carried out as follows:
Such a representation of the order of filling of atomic orbitals is called an energy diagram. Based on the electronic diagrams of individual elements, it is possible to write down their so-called electronic formulas (configurations). So, for example, an element with 15 protons and, as a consequence, 15 electrons, i.e. phosphorus (P) will have the following energy diagram:
When converted into an electronic formula, the phosphorus atom will take the form:
15 P = 1s 2 2s 2 2p 6 3s 2 3p 3
The normal size numbers to the left of the sublevel symbol show the energy level number, and the superscripts to the right of the sublevel symbol show the number of electrons in the corresponding sublevel.
Below are the electronic formulas of the first 36 elements of the periodic table by D.I. Mendeleev.
| period | Item no. | symbol | Name | electronic formula |
| I | 1 | H | hydrogen | 1s 1 |
| 2 | He | helium | 1s 2 | |
| II | 3 | Li | lithium | 1s 2 2s 1 |
| 4 | Be | beryllium | 1s 2 2s 2 | |
| 5 | B | boron | 1s 2 2s 2 2p 1 | |
| 6 | C | carbon | 1s 2 2s 2 2p 2 | |
| 7 | N | nitrogen | 1s 2 2s 2 2p 3 | |
| 8 | O | oxygen | 1s 2 2s 2 2p 4 | |
| 9 | F | fluorine | 1s 2 2s 2 2p 5 | |
| 10 | Ne | neon | 1s 2 2s 2 2p 6 | |
| III | 11 | Na | sodium | 1s 2 2s 2 2p 6 3s 1 |
| 12 | Mg | magnesium | 1s 2 2s 2 2p 6 3s 2 | |
| 13 | Al | aluminum | 1s 2 2s 2 2p 6 3s 2 3p 1 | |
| 14 | Si | silicon | 1s 2 2s 2 2p 6 3s 2 3p 2 | |
| 15 | P | phosphorus | 1s 2 2s 2 2p 6 3s 2 3p 3 | |
| 16 | S | sulfur | 1s 2 2s 2 2p 6 3s 2 3p 4 | |
| 17 | Cl | chlorine | 1s 2 2s 2 2p 6 3s 2 3p 5 | |
| 18 | Ar | argon | 1s 2 2s 2 2p 6 3s 2 3p 6 | |
| IV | 19 | K | potassium | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 |
| 20 | Ca | calcium | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 | |
| 21 | Sc | scandium | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 1 | |
| 22 | Ti | titanium | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 2 | |
| 23 | V | vanadium | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 3 | |
| 24 | Cr | chromium | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 5 here we observe the jump of one electron with s on d sublevel | |
| 25 | Mn | manganese | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 5 | |
| 26 | Fe | iron | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 6 | |
| 27 | Co | cobalt | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 7 | |
| 28 | Ni | nickel | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 8 | |
| 29 | Cu | copper | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 10 here we observe the jump of one electron with s on d sublevel | |
| 30 | Zn | zinc | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 | |
| 31 | Ga | gallium | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 1 | |
| 32 | Ge | germanium | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 2 | |
| 33 | As | arsenic | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 3 | |
| 34 | Se | selenium | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 4 | |
| 35 | Br | bromine | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 5 | |
| 36 | Kr | krypton | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 |
As already mentioned, in their ground state, electrons in atomic orbitals are located according to the principle of least energy. However, in the presence of empty p-orbitals in the ground state of the atom, often, by imparting excess energy to it, the atom can be transferred to the so-called excited state. For example, a boron atom in its ground state has an electronic configuration and an energy diagram of the following form:
And in an excited state (*), i.e. When some energy is imparted to a boron atom, its electron configuration and energy diagram will look like this:
Depending on which sublevel in the atom is filled last, chemical elements are divided into s, p, d or f.
Finding s, p, d and f elements in the table D.I. Mendeleev:
- The s-elements have the last s-sublevel to be filled. These elements include elements of the main (on the left in the table cell) subgroups of groups I and II.
- For p-elements, the p-sublevel is filled. The p-elements include the last six elements of each period, except the first and seventh, as well as elements of the main subgroups of groups III-VIII.
- d-elements are located between s- and p-elements in large periods.
- f-Elements are called lanthanides and actinides. They are listed at the bottom of the D.I. table. Mendeleev.
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
- 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.
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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. 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).
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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.
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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.
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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.
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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
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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.
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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.
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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 .
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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.
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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.
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.
Electronic configuration of an atom is a formula showing the arrangement of electrons in an atom by levels and sublevels. After studying the article, you will learn where and how electrons are located, get acquainted with quantum numbers and be able to construct the electronic configuration of an atom by its number; at the end of the article there is a table of elements.
Why study the electronic configuration of elements?
Atoms are like a construction set: there is a certain number of parts, they differ from each other, but two parts of the same type are absolutely the same. But this construction set is much more interesting than the plastic one and here’s why. The configuration changes depending on who is nearby. For example, oxygen next to hydrogen Maybe turn into water, when near sodium it turns into gas, and when near iron it completely turns it into rust. To answer the question of why this happens and predict the behavior of an atom next to another, it is necessary to study the electronic configuration, which will be discussed below.
How many electrons are in an atom?
An atom consists of a nucleus and electrons rotating around it; the nucleus consists of protons and neutrons. In the neutral state, each atom has the number of electrons equal to the number of protons in its nucleus. The number of protons is designated by the atomic number of the element, for example, sulfur has 16 protons - the 16th element of the periodic table. Gold has 79 protons - the 79th element of the periodic table. Accordingly, sulfur has 16 electrons in the neutral state, and gold has 79 electrons.
Where to look for an electron?
By observing the behavior of the electron, certain patterns were derived; they are described by quantum numbers, there are four in total:
- Principal quantum number
- Orbital quantum number
- Magnetic quantum number
- Spin quantum number
Orbital
Further, instead of the word orbit, we will use the term “orbital”; an orbital is the wave function of an electron; roughly, it is the region in which the electron spends 90% of its time.
N - level
L - shell
M l - orbital number
M s - first or second electron in the orbital
Orbital quantum number l
As a result of studying the electron cloud, they found that depending on the energy level, the cloud takes four main forms: a ball, dumbbells and two other, more complex ones. In order of increasing energy, these forms are called the s-, p-, d- and f-shell. Each of these shells can have 1 (on s), 3 (on p), 5 (on d) and 7 (on f) orbitals. The orbital quantum number is the shell in which the orbitals are located. The orbital quantum number for the s,p,d and f orbitals takes the values 0,1,2 or 3, respectively.
There is one orbital on the s-shell (L=0) - two electrons
There are three orbitals on the p-shell (L=1) - six electrons
There are five orbitals on the d-shell (L=2) - ten electrons
There are seven orbitals on the f-shell (L=3) - fourteen electrons
Magnetic quantum number m l
There are three orbitals on the p-shell, they are designated by numbers from -L to +L, that is, for the p-shell (L=1) there are orbitals “-1”, “0” and “1”. The magnetic quantum number is denoted by the letter m l.
Inside the shell, it is easier for electrons to be located in different orbitals, so the first electrons fill one in each orbital, and then a pair of electrons is added to each one.
Consider the d-shell:
The d-shell corresponds to the value L=2, that is, five orbitals (-2,-1,0,1 and 2), the first five electrons fill the shell taking the values M l =-2, M l =-1, M l =0 , M l =1,M l =2.
Spin quantum number m s
Spin is the direction of rotation of an electron around its axis, there are two directions, so the spin quantum number has two values: +1/2 and -1/2. One energy sublevel can only contain two electrons with opposite spins. The spin quantum number is denoted m s
Principal quantum number n
The main quantum number is the energy level; currently seven energy levels are known, each indicated by an Arabic numeral: 1,2,3,...7. The number of shells at each level is equal to the level number: there is one shell on the first level, two on the second, etc.
Electron number
So, any electron can be described by four quantum numbers, the combination of these numbers is unique for each position of the electron, take the first electron, the lowest energy level is N = 1, at the first level there is one shell, the first shell at any level has the shape of a ball (s -shell), i.e. L=0, the magnetic quantum number can take only one value, M l =0 and the spin will be equal to +1/2. If we take the fifth electron (in whatever atom it is), then the main quantum numbers for it will be: N=2, L=1, M=-1, spin 1/2.
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.
The structure of the electronic shells of atoms of elements of the first four periods: $s-$, $p-$ and $d-$elements. Electronic configuration of an atom. Ground and excited states of atoms
The concept of atom arose in the ancient world to denote particles of matter. Translated from Greek, atom means “indivisible.”
Electrons
The Irish physicist Stoney, based on experiments, came to the conclusion that electricity is carried by the smallest particles existing in the atoms of all chemical elements. In $1891, Mr. Stoney proposed to call 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 a unit $(–1)$. Thomson even managed to determine the speed of the electron (it is equal to the speed of light - $300,000 km/s) and the mass of the electron (it is $1836$ times less than the mass of a hydrogen atom).
Thomson and Perrin connected the poles of a current source with two metal plates - a cathode and an anode, soldered into a glass tube from which the air was evacuated. When a voltage of about 10 thousand volts was applied to the electrode plates, a luminous discharge flashed in the tube, and particles flew from the cathode (negative pole) to the anode (positive pole), which scientists first called cathode rays, and then found out that it was a stream of electrons. Electrons hitting special substances, such as those on a TV screen, cause a glow.
The conclusion was drawn: electrons escape from the atoms of the material from which the cathode is made.
Free electrons or their flow can be obtained in other ways, for example, by heating a metal wire or by shining light on metals formed by elements of the main subgroup of group I of the periodic table (for example, cesium).
State of electrons in an atom
The state of an electron in an atom is understood as the totality of information about energy certain electron in space, in which it is located. We already know that an electron in an atom does not have a trajectory of motion, i.e. we can only talk about probabilities its location in the space around the nucleus. It can be located in any part of this space surrounding the nucleus, and the set of different positions is 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 a point. If countless such photographs were superimposed, the picture would be of an electron cloud with the greatest density where there are the most of these points.
The figure shows a “cut” of such an electron density in a hydrogen atom passing through the nucleus, and the dashed line limits the sphere within which the probability of detecting an electron is $90%$. The contour closest to the nucleus covers a region of space in which the probability of detecting an electron is $10%$, the probability of detecting an electron inside the second contour from the nucleus is $20%$, inside the third is $≈30%$, etc. There is some uncertainty in the state of the electron. To characterize this special state, the German physicist W. Heisenberg introduced the concept of uncertainty principle, i.e. showed that it is impossible to simultaneously and accurately determine the energy and location of an electron. The more precisely the energy of an electron is determined, the more uncertain its position, and vice versa, having determined the position, it is impossible to determine the energy of the electron. The probability range for detecting an electron does not have clear boundaries. However, it is possible to select a space where the probability of finding an electron is maximum.
The space around the atomic nucleus in which an electron is most likely to be found is called an orbital.
It contains approximately $90%$ of the electron cloud, which means that about $90%$ of the time the electron is in this part of space. Based on their shape, there are four known types of orbitals, which are designated by the 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 the energy of its binding 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 $n$ denoting 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 are characterized by a large amount of energy. Consequently, the electrons of the outer level are least tightly bound to the atomic nucleus.
The number of energy levels (electronic layers) in an atom is equal to the number of the period in the D.I. Mendeleev system to which the chemical element belongs: atoms of elements of the first period have one energy level; second period - two; seventh period - seven.
The largest number of electrons at an energy level is determined by the formula:
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 $32$. And how, in turn, are energy levels (electronic layers) arranged?
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 sub level; the second - two; third - three; fourth - four. Sublevels, in turn, are formed by orbitals.
Each value of $n$ corresponds to a number of orbitals equal to $n^2$. According to the data presented in the table, one can trace the connection between the principal quantum number $n$ and the number of sublevels, the type and number of orbitals, and the maximum number of electrons at the sublevel and level.
Main quantum number, types and number of orbitals, maximum number of electrons in sublevels and levels.
| Energy level $(n)$ | Number of sublevels equal to $n$ | Orbital type | Number of orbitals | Maximum number of electrons | ||
| in the sublevel | in level equal to $n^2$ | in the sublevel | at a level equal to $n^2$ | |||
| $K(n=1)$ | $1$ | $1s$ | $1$ | $1$ | $2$ | $2$ |
| $L(n=2)$ | $2$ | $2s$ | $1$ | $4$ | $2$ | $8$ |
| $2p$ | $3$ | $6$ | ||||
| $M(n=3)$ | $3$ | $3s$ | $1$ | $9$ | $2$ | $18$ |
| $3p$ | $3$ | $6$ | ||||
| $3d$ | $5$ | $10$ | ||||
| $N(n=4)$ | $4$ | $4s$ | $1$ | $16$ | $2$ | $32$ |
| $4p$ | $3$ | $6$ | ||||
| $4d$ | $5$ | $10$ | ||||
| $4f$ | $7$ | $14$ | ||||
Sublevels are usually denoted by Latin letters, as well as the shape of the orbitals of which they consist: $s, p, d, f$. So:
- $s$-sublevel - the first sublevel of each energy level closest to the atomic nucleus, consists of one $s$-orbital;
- $p$-sublevel - the second sublevel of each, except the first, energy level, consists of three $p$-orbitals;
- $d$-sublevel - the third sublevel of each, starting from the third, energy level, consists of five $d$-orbitals;
- The $f$-sublevel of each, starting from the fourth energy level, consists of seven $f$-orbitals.
Atomic nucleus
But not only electrons are part of atoms. Physicist Henri Becquerel discovered that a natural mineral containing a uranium salt also emits unknown radiation, exposing photographic films shielded from light. This phenomenon was called radioactivity.
There are three types of radioactive rays:
- $α$-rays, which consist of $α$-particles having a charge $2$ times greater than the charge of an electron, but with a positive sign, and a mass $4$ times greater than the mass of a hydrogen atom;
- $β$-rays represent a flow of electrons;
- $γ$-rays are electromagnetic waves with negligible mass that do not carry an electrical charge.
Consequently, the atom has a complex structure - it consists of a positively charged nucleus and electrons.
How is an atom structured?
In 1910, in Cambridge, near London, Ernest Rutherford and his students and colleagues studied the scattering of $α$ particles passing through thin gold foil and falling on a screen. Alpha particles usually deviated from the original direction by only one degree, seemingly confirming the uniformity and uniformity of the properties of gold atoms. And suddenly the researchers noticed that some $α$ particles abruptly changed the direction of their path, as if encountering some kind of obstacle.
By placing a screen in front of the foil, Rutherford was able to detect even those rare cases when $α$ particles, reflected from gold atoms, flew in the opposite direction.
Calculations showed that the observed phenomena could occur if the entire mass of the atom and all its positive charge were concentrated in a tiny central nucleus. The radius of the nucleus, as it turned out, is 100,000 times smaller than the radius of the entire atom, the region in which electrons with a negative charge are located. If we apply a figurative comparison, then the entire volume of an atom can be likened to the stadium in Luzhniki, and the nucleus can be likened to a soccer ball located in the center of the field.
An atom of any chemical element is comparable to a tiny solar system. Therefore, this model of the atom, proposed by Rutherford, is called planetary.
Protons and Neutrons
It turns out that the tiny atomic nucleus, in which the entire mass of the atom is concentrated, consists of two types of particles - protons and neutrons.
Protons have a charge equal to the charge of the electrons, but opposite in sign $(+1)$, and a mass equal to the mass of the hydrogen atom (it is taken as unity in chemistry). Protons are designated by the sign $↙(1)↖(1)p$ (or $p+$). Neutrons do not carry a charge, they are neutral and have a mass equal to the mass of a proton, i.e. $1$. Neutrons are designated by the sign $↙(0)↖(1)n$ (or $n^0$).
Protons and neutrons together are called nucleons(from lat. nucleus- core).
The sum of the number of protons and neutrons in an atom is called mass number. For example, the mass number of an aluminum atom is:
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 as follows: $e↖(-)$.
Since the atom is electrically neutral, it is also obvious that that the number of protons and electrons in an atom is the same. It is equal to the atomic number of the chemical element, assigned to it in the Periodic Table. For example, the nucleus of an iron atom contains $26$ protons, and $26$ electrons revolve around the nucleus. How to determine the number of neutrons?
As is known, the mass of an atom consists of the mass of protons and neutrons. Knowing the serial number of the element $(Z)$, i.e. the number of protons, and the mass number $(A)$, equal to the sum of the numbers of protons and neutrons, the number of neutrons $(N)$ can be found using the formula:
For example, the number of neutrons in an iron atom is:
$56 – 26 = 30$.
The table presents the main characteristics of elementary particles.
Basic characteristics of elementary particles.
Isotopes
Varieties of atoms of the same element that have the same nuclear charge but different mass numbers are called isotopes.
Word isotope consists of two Greek words: isos- identical and topos- place, means “occupying one place” (cell) in the Periodic Table of Elements.
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.
Usually, the relative atomic mass of a chemical element 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, therefore the values of atomic masses are quite often fractional. For example, natural chlorine atoms are a mixture of two isotopes - $35$ (there are $75%$ in nature) and $37$ (there are $25%$ in nature); therefore, the relative atomic mass of chlorine is $35.5$. Isotopes of chlorine are written as follows:
$↖(35)↙(17)(Cl)$ and $↖(37)↙(17)(Cl)$
The chemical properties of chlorine isotopes are exactly the same, as are the isotopes of most chemical elements, for example potassium, argon:
$↖(39)↙(19)(K)$ and $↖(40)↙(19)(K)$, $↖(39)↙(18)(Ar)$ and $↖(40)↙(18 )(Ar)$
However, hydrogen isotopes vary greatly in properties due to the dramatic multiple increase in their relative atomic mass; they were even given individual names and chemical symbols: protium - $↖(1)↙(1)(H)$; deuterium - $↖(2)↙(1)(H)$, or $↖(2)↙(1)(D)$; tritium - $↖(3)↙(1)(H)$, or $↖(3)↙(1)(T)$.
Now we can give a modern, more rigorous and scientific definition of a chemical element.
A chemical element is a collection of atoms with the same nuclear charge.
The structure of the electronic shells of atoms of elements of the first four periods
Let's consider the display of electronic configurations of atoms of elements according to the periods of the D.I. Mendeleev system.
Elements of the first period.
Diagrams of the electronic structure of atoms show the distribution of electrons across electronic layers (energy levels).
Electronic formulas of atoms show the distribution of electrons across energy levels and sublevels.
Graphic electronic formulas of atoms show the distribution of electrons not only across levels and sublevels, but also across orbitals.
In a helium atom, the first electron layer is complete - it contains $2$ electrons.
Hydrogen and helium are $s$ elements; the $s$ orbital of these atoms is filled with electrons.
Elements of the second period.
For all second-period elements, the first electron 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 contains $8$ electrons.
Elements of the third period.
For atoms of elements of the third period, the first and second electron layers are completed, so the third electron layer is filled, in which electrons can occupy the 3s-, 3p- and 3d-sub levels.
The structure of the electronic shells of atoms of elements of the third period.
The magnesium atom completes its $3.5$ electron orbital. $Na$ and $Mg$ are $s$-elements.
In aluminum and subsequent elements, the $3d$ sublevel is filled with electrons.
| $↙(18)(Ar)$ Argon |  |
$1s^2(2)s^2(2)p^6(3)s^2(3)p^6$ |  |
An argon atom has $8$ electrons in its outer layer (third electron layer). As the outer layer is completed, but in total in the third electron layer, as you already know, there can be 18 electrons, which means that the elements of the third period have unfilled $3d$-orbitals.
All elements from $Al$ to $Ar$ are $р$ -elements.
$s-$ and $p$ -elements form main subgroups in the Periodic Table.
Elements of the fourth period.
Potassium and calcium atoms have a fourth electron layer and the $4s$ sublevel is filled, because it has lower energy than the $3d$ sublevel. To simplify the graphical electronic formulas of atoms of elements of the fourth period:
- Let us denote the conventional graphical electronic formula of argon as follows: $Ar$;
- We will not depict sublevels that are not filled in these atoms.
$K, Ca$ - $s$ -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 side subgroups, their outer electron layer is filled, they are classified as transitional 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 $3d^5$ and $3d^(10)$ electronic configurations:
$↙(24)(Cr)$ $1s^(2)2s^(2)2p^(6)3s^(2)3p^(6)3d^(4) 4s^(2)…$ 
$↙(29)(Cu)$ $1s^(2)2s^(2)2p^(6)3s^(2)3p^(6)3d^(9)4s^(2)…$ 
| Element symbol, serial number, name | Electronic structure diagram | Electronic formula | Graphical electronic formula |
| $↙(19)(K)$ Potassium |  |
$1s^2(2)s^2(2)p^6(3)p^6(4)s^1$ | |
| $↙(20)(C)$ Calcium |  |
$1s^2(2)s^2(2)p^6(3)p^6(4)s^2$ | |
| $↙(21)(Sc)$ Scandium |  |
$1s^2(2)s^2(2)p^6(3)p^6(4)s^1(3)d^1$ or $1s^2(2)s^2(2)p ^6(3)p^6(3)d^1(4)s^1$ |  |
| $↙(22)(Ti)$ Titanium |  |
$1s^2(2)s^2(2)p^6(3)p^6(4)s^2(3)d^2$ or $1s^2(2)s^2(2)p ^6(3)p^6(3)d^2(4)s^2$ |  |
| $↙(23)(V)$ Vanadium |  |
$1s^2(2)s^2(2)p^6(3)p^6(4)s^2(3)d^3$ or $1s^2(2)s^2(2)p ^6(3)p^6(3)d^3(4)s^2$ |  |
| $↙(24)(Cr)$ Chrome |  |
$1s^2(2)s^2(2)p^6(3)p^6(4)s^1(3)d^5$ or $1s^2(2)s^2(2)p ^6(3)p^6(3)d^5(4)s^1$ |  |
| $↙(29)(Cu)$ Chrome |  |
$1s^2(2)s^2(2)p^6(3)p^6(4)s^1(3)d^(10)$ or $1s^2(2)s^2(2 )p^6(3)p^6(3)d^(10)(4)s^1$ |  |
| $↙(30)(Zn)$ Zinc |  |
$1s^2(2)s^2(2)p^6(3)p^6(4)s^2(3)d^(10)$ or $1s^2(2)s^2(2 )p^6(3)p^6(3)d^(10)(4)s^2$ |  |
| $↙(31)(Ga)$ Gallium |  |
$1s^2(2)s^2(2)p^6(3)p^6(4)s^2(3)d^(10)4p^(1)$ or $1s^2(2) s^2(2)p^6(3)p^6(3)d^(10)(4)s^(2)4p^(1)$ |  |
| $↙(36)(Kr)$ Krypton |  |
$1s^2(2)s^2(2)p^6(3)p^6(4)s^2(3)d^(10)4p^6$ or $1s^2(2)s^ 2(2)p^6(3)p^6(3)d^(10)(4)s^(2)4p^6$ |  |
In the zinc atom, the third electron layer is complete - all $3s, 3p$ and $3d$ sublevels 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 $Кr$ - $р$ -elements.
The outer (fourth) layer of the krypton atom is complete and has $8$ electrons. But in total in the fourth electron layer, as you know, there can be $32$ electrons; the krypton atom still has unfilled $4d-$ and $4f$ sublevels.
For elements of the fifth period, sublevels are filled in in the following order: $5s → 4d → 5p$. And there are also exceptions associated with the “failure” of electrons in $↙(41)Nb$, $↙(42)Mo$, $↙(44)Ru$, $↙(45)Rh$, $↙(46) Pd$, $↙(47)Ag$. $f$ appears in the sixth and seventh periods -elements, i.e. elements for which the $4f-$ and $5f$ sublevels of the third outside electronic layer are filled, respectively.
$4f$ -elements called lanthanides.
$5f$ -elements called actinides.
The order of filling electronic sublevels in atoms of elements of the sixth period: $↙(55)Cs$ and $↙(56)Ba$ - $6s$ elements; $↙(57)La ... 6s^(2)5d^(1)$ - $5d$-element; $↙(58)Се$ – $↙(71)Lu - 4f$-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 of electronic orbitals is violated, which, for example, is associated with greater energy stability of half and completely filled $f$-sublevels, i.e. $nf^7$ and $nf^(14)$.
Depending on which sublevel of the atom is filled with electrons last, all elements, as you already understood, 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 intercalary decades of large periods located between $s-$ and $p-$elements. They are also called transition elements;
- $f$ -elements; electrons fill the $f-$sublevel of the third outer level of the atom; these include lanthanides and actinides.
Electronic configuration of an atom. Ground and excited states of atoms
Swiss physicist W. Pauli in $1925 found that an atom can have no more than two electrons in one orbital, having opposite (antiparallel) backs (translated from English as a spindle), i.e. possessing properties that can be conventionally 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 an orbital, it is called unpaired, if two, then this paired electrons, i.e. electrons with opposite spins.
The figure shows a diagram of dividing energy levels into sublevels.
$s-$ Orbital, as you already know, has a spherical shape. The electron of the hydrogen atom $(n = 1)$ is located in this orbital and is unpaired. For this reason it electronic formula, or electronic configuration, is written like this: $1s^1$. In electronic formulas, the number of the energy level is indicated by the number in front of the letter $(1...)$, the Latin letter denotes the sublevel (type of orbital), and the number written to the right above the letter (as an exponent) shows the number of electrons in the sublevel.
For a helium atom He, which has two paired electrons in one $s-$orbital, this formula is: $1s^2$. The electron shell of the helium atom is complete and very stable. Helium is a noble gas. At the second energy level $(n = 2)$ there are four orbitals, one $s$ and three $p$. Electrons of the $s$-orbital of the second level ($2s$-orbital) have higher energy, because are at a greater distance from the nucleus than the electrons of the $1s$ orbital $(n = 2)$. In general, for each value of $n$ there is one $s-$orbital, but with a corresponding supply of electron energy on it and, therefore, with a corresponding diameter, growing as the value of $n$ increases. The $s-$Orbital, as you already know , has a spherical shape. The electron of the hydrogen atom $(n = 1)$ is located in this orbital and is unpaired. Therefore, its electronic formula, or electronic configuration, is written as follows: $1s^1$. In electronic formulas, the number of the energy level is indicated by the number in front of the letter $(1...)$, the Latin letter denotes the sublevel (type of orbital), and the number written to the right above the letter (as an exponent) shows the number of electrons in the sublevel.
For a helium atom $He$, which has two paired electrons in one $s-$orbital, this formula is: $1s^2$. The electron shell of the helium atom is complete and very stable. Helium is a noble gas. At the second energy level $(n = 2)$ there are four orbitals, one $s$ and three $p$. Electrons of $s-$orbitals of the second level ($2s$-orbitals) have higher energy, because are at a greater distance from the nucleus than the electrons of the $1s$ orbital $(n = 2)$. In general, for each value of $n$ there is one $s-$orbital, but with a corresponding supply of electron energy on it and, therefore, with a corresponding diameter, growing as the value of $n$ increases. 
$p-$ Orbital has the shape of a dumbbell, or a voluminous figure eight. All three $p$-orbitals are located in the atom mutually perpendicular along the spatial coordinates drawn through the nucleus of the atom. It should be emphasized once again that each energy level (electronic layer), starting from $n= 2$, has three $p$-orbitals. As the value of $n$ increases, electrons occupy $p$-orbitals located at large distances from the nucleus and directed along the $x, y, z$ axes.
For elements of the second period $(n = 2)$, first one $s$-orbital is filled, and then three $p$-orbitals; electronic formula $Li: 1s^(2)2s^(1)$. The $2s^1$ electron is more weakly bound to the nucleus of the atom, so the lithium atom can easily give it up (as you obviously remember, this process is called oxidation), turning into a lithium ion $Li^+$.
In the beryllium Be atom, the fourth electron is also located in the $2s$ orbital: $1s^(2)2s^(2)$. The two outer electrons of the beryllium atom are easily detached - $B^0$ is oxidized into the $Be^(2+)$ cation.
In the boron atom, the fifth electron occupies the $2p$ orbital: $1s^(2)2s^(2)2p^(1)$. Next, the $C, N, O, F$ atoms are filled with $2p$-orbitals, which ends with the noble gas neon: $1s^(2)2s^(2)2p^(6)$.
For elements of the third period, the $3s-$ and $3p$ orbitals are filled, respectively. Five $d$-orbitals of the third level remain free:
$↙(11)Na 1s^(2)2s^(2)2p^(6)3s^(1)$,
$↙(17)Cl 1s^(2)2s^(2)2p^(6)3s^(2)3p^(5)$,
$↙(18)Ar 1s^(2)2s^(2)2p^(6)3s^(2)3p^(6)$.
Sometimes in diagrams depicting the distribution of electrons in atoms, only the number of electrons at each energy level is indicated, i.e. write abbreviated electronic formulas of atoms of chemical elements, in contrast to the full electronic formulas given above, for example:
$↙(11)Na 2, 8, 1;$ $↙(17)Cl 2, 8, 7;$ $↙(18)Ar 2, 8, 8$.
For elements of large periods (fourth and fifth), the first two electrons occupy $4s-$ and $5s$ orbitals, respectively: $↙(19)K 2, 8, 8, 1;$ $↙(38)Sr 2, 8, 18, 8, 2$. Starting from the third element of each major period, the next ten electrons will go to the previous $3d-$ and $4d-$orbitals, respectively (for elements of side subgroups): $↙(23)V 2, 8, 11, 2;$ $↙( 26)Fr 2, 8, 14, 2;$ $↙(40)Zr 2, 8, 18, 10, 2;$ $↙(43)Tc 2, 8, 18, 13, 2$. As a rule, when the previous $d$-sublevel is filled, the outer ($4р-$ and $5р-$, respectively) $р-$sublevel will begin to be filled: $↙(33)As 2, 8, 18, 5;$ $ ↙(52)Te 2, 8, 18, 18, 6$.
For elements of large periods - the sixth and incomplete seventh - electronic levels and sublevels are filled with electrons, as a rule, like this: the first two electrons enter the outer $s-$sublevel: $↙(56)Ba 2, 8, 18, 18, 8, 2;$ $↙(87)Fr 2, 8, 18, 32, 18, 8, 1$; the next one electron (for $La$ and $Ca$) to the previous $d$-sublevel: $↙(57)La 2, 8, 18, 18, 9, 2$ and $↙(89)Ac 2, 8, 18, 32, 18, 9, 2$.
Then the next $14$ electrons will go to the third outer energy level, to the $4f$ and $5f$ orbitals of lanthanides and actinides, respectively: $↙(64)Gd 2, 8, 18, 25, 9, 2;$ $↙(92 )U 2, 8, 18, 32, 21, 9, 2$.
Then the second external energy level ($d$-sublevel) of elements of side subgroups will begin to build up again: $↙(73)Ta 2, 8, 18, 32, 11, 2;$ $↙(104)Rf 2, 8, 18 , 32, 32, 10, 2$. And finally, only after the $d$-sublevel is completely filled with ten electrons will the $p$-sublevel be filled again: $↙(86)Rn 2, 8, 18, 32, 18, 8$.
Very often the structure of the electronic shells of atoms is depicted using energy or quantum cells - the so-called graphic electronic formulas. 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 principle, according to which there can be no more than two electrons in a cell (orbital), but with antiparallel spins, 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 be in opposite directions.
