How to make an electronic formula of an element in chemistry. Electronic and electron-graphic formulas of atoms of elements

DEFINITION

Electronic formula(configuration) of an atom of a chemical element shows the arrangement of electrons on electron shells (levels and sublevels) in an atom or molecule.

Most often, electronic formulas are written for atoms in the ground or excited state and for ions.

There are several rules that must be considered when compiling the electronic formula of an atom of a chemical element. These are the Pauli principle, Klechkovsky's rules or Hund's rule.

When compiling an electronic formula, it should be taken into account that the number of the period of a chemical element determines the number of energy levels (shells) in the atom, and its serial number determines the number of electrons.

According to Klechkovsky's rule, the energy levels are filled in the ascending order of the sum of the main and orbital quantum numbers (n + l), and for equal values ​​of this sum, in the ascending order of n:

1s< 2s < 2p < 3s < 3p < 4s ≈ 3d < 4p < 5s ≈ 4d < 5p < 6s ≈ 5d ≈ 4f < 6p и т.д.

Thus, the value n + l = 5 corresponds to the energy sublevels 3d (n = 3, l=2), 4d (n=4, l=1) and 5s (n=5, l=0). The first of these sublevels is filled with the one with the lower value of the main quantum number.

The behavior of electrons in atoms is subject to the exclusion principle formulated by the Swiss scientist W. Pauli: there cannot be two electrons in an atom that would have the same all four quantum numbers. According to Pauli principle, on one orbital, characterized by certain values ​​of three quantum numbers (principal, orbital and magnetic), there can be only two electrons that differ in the value of the spin quantum number. It follows from the Pauli principle consequence: the maximum possible number of electrons in each energy level is twice the square of the principal quantum number.

Electronic formula of the atom

The electronic formula of an atom is depicted as follows: each energy level corresponds to a certain main quantum number n, denoted by an Arabic numeral; each number is followed by a letter corresponding to the energy sublevel and denoting the orbital quantum number. The superscript next to the letter shows the number of electrons in the sublevel. For example, the electronic formula of the sodium atom is:

11 N 1s 2 2s 2 2p 6 3s 1 .

When filling energy sublevels with electrons, it is also necessary to observe hund rule: in this sublevel, electrons tend to occupy energy states in such a way that the total spin is maximum (this is most clearly reflected in the preparation of electron-graphic formulas).

Examples of problem solving

EXAMPLE 1

Many metals are common in nature, not only in the composition of various rocks or minerals, but also in free - native form. Examples include gold, silver and copper. However, active metallic elements such as sodium, whose electron-graphic formula we will study, do not occur as a simple substance. The reason is their high reactivity, which leads to the rapid oxidation of the substance by atmospheric oxygen. That is why in the laboratory the metal is stored under a layer of kerosene or technical oil. The chemical activity of all alkali metal elements can be explained by the structural features of their atoms. Let's consider the electron-graphic formula of sodium and find out how its characteristics affect the physical properties and features of interaction with other substances.

sodium atom

The position of an element in the main subgroup of the first group of the periodic system affects the structure of its electrically neutral particle. This diagram illustrates the arrangement of electrons around the nucleus of an atom and determines the number of energy levels in it:

The number of protons, neutrons, electrons in a sodium atom will be equal to 11, 12, 11, respectively. The proton number and the number of electrons are determined by the element's serial number, and the number of neutral nuclear particles will be equal to the difference between the nucleon number (atomic mass) and the proton number (serial number ). To record the distribution of negatively charged particles in an atom, you can use the following electronic formula: 1s 2 2s 2 2p 6 3s 1.

The relationship between the structure of the atom and the properties of matter

The properties of sodium as an alkali metal can be explained by the fact that it belongs to the s-elements, its valency is 1, and the oxidation state is +1. One unpaired electron on the third, last, layer determines its reduction characteristics. In reactions with other atoms, sodium always donates its own negative particle to more electronegative elements. For example, being oxidized by atmospheric oxygen, Na atoms become positively charged particles - cations that are part of the basic oxide Na 2 O molecule. This reaction has the following form:

4Na + O 2 \u003d 2Na 2 O.

Physical properties

The electronic graphic formula of sodium and its crystal lattice determine such element parameters as the state of aggregation, melting and boiling points, as well as the ability to conduct heat and electric current. Sodium is a light (density 0.97 g/cm3) and very soft silvery metal. The presence of freely moving electrons in the crystal lattice causes high thermal and electrical conductivity. It occurs naturally in minerals such as common salt NaCl and sylvinite NaCl × KCl. Sodium is very common not only in inanimate nature, for example, in the composition of rock salt deposits or sea water of the seas and oceans. He, along with chlorine, sulfur, calcium, phosphorus and other elements, is among the ten most important organogenic chemical elements that form living biological systems.

Features of chemical properties

The electron-graphic formula of sodium clearly shows that the only s-electron rotating on the last, third energy layer of the Na atom is weakly bound to the positively charged nucleus. It easily leaves the limits of the atom, so sodium in reactions with oxygen, water, hydrogen and nitrogen behaves like a strong reducing agent. Here are examples of reaction equations typical for alkali metals:

2Na + H 2 \u003d 2NaH;

6Na + N 2 \u003d 2Na 3 N;

2Na + 2H 2 O \u003d 2NaOH + H 2.

The reaction with water ends with the formation of chemically aggressive compounds - alkalis. Sodium hydroxide, also called, exhibits the properties of active bases and in the solid state has found application as a gas desiccant. In industry, sodium metal is obtained by electrolysis of a salt melt - sodium chloride or the corresponding hydroxide, while a layer of sodium metal is formed on the cathode.

In our article, we examined the electronic graphic formula of sodium, and also studied its properties and production in industry.

In order to learn how to compose electron-graphic formulas, it is significant to realize the theory of the structure of the nuclear nucleus. The nucleus of an atom is made up of protons and neutrons. Electrons are located in electron orbitals around the nucleus of an atom.

You will need

• - pen;
• - note paper;
• - the periodic system of elements (Mendeleev's table).

Instruction

1. Electrons in an atom occupy vacant orbitals in a sequence called the energy scale: 1s/2s, 2p/3s, 3p/4s, 3d, 4p/5s, 4d, 5p/6s, 4d, 5d, 6p/7s, 5f, 6d, 7p . Two electrons with opposite spins - directions of rotation can be located on one orbital.

2. The design of electron shells is expressed with the support of graphic electronic formulas. Use a matrix to write a formula. One cell can contain one or two electrons with opposite spins. Electrons are represented by arrows. The matrix clearly shows that two electrons can be located in the s-orbital, 6 in the p-orbital, 10 in the d-orbital, and 14 in the f-orbital.

3. Consider the rule for compiling an electronic graphic formula using manganese as an example. Find manganese in the periodic table. Its serial number is 25, which means there are 25 electrons in the atom, this is an element of the fourth period.

4. Write down the serial number and symbol of the element next to the matrix. In accordance with the energy scale, fill in the 1s, 2s, 2p, 3s, 3p, 4s tiers step by step, entering two electrons per cell. You get 2+2+6+2+6+2=20 electrons. These tiers are completely filled.

5. You have five more electrons left and an empty 3d tier. Arrange the electrons in the cells of the d-sublevel, starting from the left. Place the electrons with identical spins in the cells first one by one. If all cells are filled, starting from the left, add a second electron with the opposite spin. Manganese has five d-electrons, located one at a time in the entire cell.

6. Electron graphic formulas clearly show the number of unpaired electrons that determine the valence.

When creating theoretical and factual works in mathematics, physics, chemistry, a student or schoolchild is faced with the need to insert special symbols and difficult formulas. Having the Word application from the Microsoft office suite, it is allowed to type an electronic formula every difficulty.

Instruction

1. Open the newest document in Microsoft Word. Give it a name and save it in the same folder where your work is, so that you don’t look for it in the future.

2. Go to the "Insert" tab. On the right, find the symbol ?, and next to it is the inscription "Formula". Click on the arrow. A window will appear where you can prefer the built-in formula, say the formula quadratic equation.

3. Click on the arrow and a variety of symbols will appear on the top panel that you may need when writing this particular formula. By changing it the way you want it, you can save it. From now on, it will drop out in the list of built-in formulas.

4. If you need to transfer the formula to text, the one that later needs to be placed on the site, then click on the energetic field with it with the right mouse button and select not the highly professional, but the linear method of writing. In particular, the formula of the same quadratic equation in this case takes the form: x=(-b±?(b^2-4ac))/2a.

5. Another option for writing an electronic formula in Word is through the constructor. Hold down the Alt and = keys at the same time. You will immediately have a field for writing a formula, and a constructor will open in the top panel. Here you can prefer all the signs that may be required to write an equation and solve any problem.

6. Some linear notation symbols may be obscure to a reader unfamiliar with computer symbols. In this case, it makes sense to save the most difficult formulas or equations in graphical form. To do this, open the easiest graphic editor Paint: "Start" - "Programs" - "Paint". After that, zoom in on the formula document so that it takes up every screen. This is necessary so that the saved image has the highest resolution. Press PrtScr on your keyboard, go to Paint and press Ctrl+V.

7. Trim off any excess. As a result, you will get a solid image with the necessary formula.

Related videos

Note!
Remember that chemistry is a science of exceptions. The atoms of the secondary subgroups of the Periodic system have an electron "breakthrough". For example, in chromium with atomic number 24, one of the electrons from the 4s-tier goes to the d-tier cell. Molybdenum, niobium, etc. have a similar result. In addition, there is a representation of the excited state of the atom, when paired electrons are unpaired and transferred to neighboring orbitals. Therefore, when compiling electronic graphic formulas for the elements of the fifth and subsequent periods of the secondary subgroup, refer to the reference book.

Problem 56.
Write an electron-graphic formula for the elements of the 4th period, determine their valence electrons and characterize them using quantum numbers.
Decision:
Electronic formulas display the distribution of electrons in an atom by energy levels, sublevels (atomic orbitals). Electronic configuration denoted by character groups nl x , where n is the principal quantum number, l - orbital quantum number (instead of it indicate the corresponding letter designation - s, p, d, f ), x is the number of electrons in a given sublevel (orbitals). In this case, it should be taken into account that the electron occupies the energy sublevel at which it has the lowest energy - a smaller sum n+1 (Klechkovsky's rule ). The sequence of filling energy levels and sublevels is as follows:

1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s (5d1) 4f 5d 6p 7s (6d1-2) 5f 6d 7p

a) Element No. 19
Since the number of electrons in an atom of an element is equal to its serial number in the table of D.I. Mendeleev, then for the 19th element - potassium (K - serial number 19), the electronic formula looks like:

Valence electron potassium 4s 1 - are located on 4s - sublevel There is 1 electron in the valence orbital of the K atom. Therefore, the element is placed in the first group of the periodic system of D.I. Mendeleev.

b) Element No. 20
For element No. 20 - calcium (Ca - serial No. 20), the electronic formula is:

Valence electrons calcium 4s 2 - are located on 4s -sublevel There are 2 electrons in the valence orbital of the Ca atom. Therefore, the element is placed in the second group of the periodic system of D.I. Mendeleev.

c) Element No. 21
For element No. 21 - scandium (Ca - serial No. 21), the electronic formula is:

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

Valence electrons scandium 4s 2 3d 1 - are located on 4s - and 3d -sublevels. There are 3 electrons in the valence orbitals of the Sc atom. Therefore, the element is placed in the third group of the periodic system of D.I. Mendeleev.

d) Element No. 22
For element No. 22 - titanium (Ti - serial No. 22), the electronic formula is:

1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d2

Valence electrons of scandium 4s 2 3d 2 - are located on 4s- and 3d- sublevels. There are 4 electrons in the valence orbitals of the Ti atom. Therefore, the element is placed in the fourth group of the periodic system of D.I. Mendeleev.

e) Element No. 23
For element No. 23 - vanadium (V - serial No. 23), the electronic formula is:

1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d3

Valence electrons of scandium 4s 2 3d 3 - are located on 4s- and 3d- sublevels. There are 5 electrons in the valence orbitals of the V atom. Therefore, the element is placed in the fifth group of the periodic system of D.I. Mendeleev.

f) Element No. 24
For the chromium element (Cr - serial number 24), the electronic formula is:

1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d5

Valence electrons chrome 4s 1 3d 5 - are located on 4s- and 3- sublevels. There are 6 electrons in the valence orbitals of the Cr atom. Therefore, the element is placed in the sixth group of the periodic system of D.I. Mendeleev.
At the chromium atom, one electron from the 4s sublevel passes to the 3d sublevel, and in this case the chromium atom acquires a more stable state 4s 1 3d 5 than 4s 2 3d 4 . This is explained by the fact that it is energetically more favorable for the chromium atom when there are not 4 but 5 electrons on the 3d sublevel - all cells are filled with one electron each. Thus, the 4s 1 3d 5 valence electron configuration is energetically more favorable for the chromium atom than 4s 2 3d 4 .

g) Element No. 25 - manganese (Mn - serial No. 25) the electronic formula is:

1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d5

Valence electrons of manganese 4s 2 3d 5 - are located on 4s- and 3d- sublevels. There are 7 electrons in the valence orbitals of the Mn atom. Therefore, the element is placed in the seventh group of the periodic system of D.I. Mendeleev.

h) Element No. 26 - iron (Fe - serial No. 26) the electronic formula is:

1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d6

Valence electrons gland 4s 2 3d 6 - are located on 4s- and 3d -sublevels. There are 8 electrons in the valence orbitals of the Fe atom. Therefore, the element is placed in the eighth group of the periodic system of D.I. Mendeleev.

j) Element No. 27 - sobalt (Co - serial No. 27) the electronic formula is:

1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d7

Valence electrons sobalt 4s 2 3d 7 - are located on 4s- and 3d- sublevels. There are 9 electrons in the valence orbitals of the Co atom. Therefore, the element is placed in the ninth group of the periodic system of D.I. Mendeleev.

k) Element No. 28 - nickel (Ni - serial No. 28) the electronic formula is:

1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 8

Nickel valence electrons 4s 2 3d 8 - are located on 4s- and 3d- sublevels. There are 10 electrons in the valence orbitals of the Ni atom. Therefore, the element is placed in the tenth group of the periodic system of D.I. Mendeleev.

l) Element No. 29 - copper (Cu - serial No. 29) the electronic formula is:

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

Valence electrons copper 4s 1 3d 10 - are located on 4s- and 3d- sublevels. There are 11 electrons in the valence orbitals of the Cu atom. Therefore, the element is placed in the eleventh group of the periodic system of D.I. Mendeleev.
The copper atom has a slip ( "failure"): one electron of the 4s sublevel goes to the 3d sublevel. This is explained by the fact that the state of the atom is considered to be more energetically favorable if the d-sublevel contains not 9, but 10 electrons. Because it is energetically more favorable for the copper atom when all five d-cells on the 3d-sublevel are filled, but not when four d-cells are filled, but on the fifth one there is only one electron. To fill the fifth d-cell of the 3d-sublevel, one electron of the 4s-sublevel goes to the 3d-sublevel, as if " fails". Thus, the valence electron configuration 4s 1 3d 10, rather than 4s 2 3d 9, is energetically more favorable for the copper atom.

m) Element No. 30 - zinc (Zn - serial No. 30) the electronic formula is:

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

Valence electrons of zinc 4s 2 3d 10 - are located on 4s- and 3d- sublevels. There are 12 electrons in the valence orbitals of the Zn atom. Therefore, the element is placed in the twelfth group of the periodic system of D.I. Mendeleev.

o) Element No. 31 - gallium (Ga - serial No. 31) the electronic formula is:

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

The valence electrons of gallium 4s 2 3d 10 4p 1 - are located on 4s-, 3d- and 4p- sublevels. There are 13 electrons in the valence orbitals of the Ga atom. Therefore, the element is placed in the thirteenth group of the periodic system of D.I. Mendeleev.

o) Element No. 32 - germanium (Ge - serial No. 32) the electronic formula is:

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

Valence electrons germanium 4s 2 3d 10 4r 2 - are located on 4s-, 3d- and 4p- sublevels. There are 14 electrons in the valence orbitals of the Ge atom. Therefore, the element is placed in the fourteenth group of the periodic system of D.I. Mendeleev.

p) Element No. 33 - arsenic (As - serial No. 33) the electronic formula is:

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

Valence electrons arsenic 4s 2 3d 10 4p 3 - are located on 4s-, 3d- and 4p- sublevels. There are 15 electrons in the valence orbitals of the As atom. Therefore, the element is placed in the fifteenth group of the periodic system of D.I. Mendeleev.

c) Element No. 34 - selenium (Se - serial No. 34) the electronic formula is:

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

Valence electrons Selene 4s 2 3d 10 4p 4 - are located on 4s-, 3d- and 4p- sublevels. There are 16 electrons in the valence orbitals of the Se atom. Therefore, the element is placed in the sixteenth group of the periodic system of D.I. Mendeleev.

c) Element No. 35 - bromine (Br - serial No. 35) the electronic formula is:

1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4r 5

Valence electrons bromine 4s 2 3d 10 4 r 5 - are located on 4s-, 3d- and 4p -sublevels. There are 17 electrons in the valence orbitals of the Br atom. Therefore, the element is placed in the seventeenth group of the periodic system of D.I. Mendeleev.

r) Element No. 36 - krypton (Kr - serial No. 36) the electronic formula looks like:

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

Valence electrons krypton 4s 2 3d 10 4p 6 - are located on 4s-, 3d- and 4p- sublevels. There are 18 electrons in the valence orbitals of the Kr atom. Therefore, the element is placed in the eighteenth group of the periodic system of D.I. Mendeleev.

The structure of the electron shells of atoms of the elements of the first four periods: $s-$, $p-$ and $d-$elements. The electronic configuration of the atom. Ground and excited states of atoms

The concept of an atom arose in the ancient world to designate the particles of matter. In Greek, atom means "indivisible".

Electrons

The Irish physicist Stoney, on the basis of experiments, came to the conclusion that electricity is carried by the smallest particles that exist in the atoms of all chemical elements. In $1891$, Stoney proposed to call these particles electrons, which in Greek means "amber".

A few years after the electron got its name, English physicist Joseph Thomson and French physicist Jean Perrin proved that electrons carry a negative charge. This is the smallest negative charge, which in chemistry is taken as the 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 the 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 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 applied, for example, to a TV screen, cause a glow.

The conclusion was made: electrons escape from the atoms of the material from which the cathode is made.

Free electrons or their flux can also be obtained in other ways, for example, by heating a metal wire or by falling light on metals formed by elements of the main subgroup of group I of the periodic table (for example, cesium).

The state of electrons in an atom

The state of an electron in an atom is understood as a set of information about energy specific 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. can only talk about probabilities finding it in the space around the nucleus. It can be located in any part of this space surrounding the nucleus, and the totality of its various positions is considered as an electron cloud with a certain negative charge density. Figuratively, this can be imagined as follows: if it were possible to photograph the position of an electron in an atom in hundredths or millionths of a second, as in a photo finish, then the electron in such photographs would be represented as a point. Overlaying countless such photographs would result in a picture of an electron cloud with the highest density where there are most of these points.

The figure shows a "cut" of such an electron density in a hydrogen atom passing through the nucleus, and a sphere is bounded by a dashed line, inside which the probability of finding an electron is $90%$. The contour closest to the nucleus covers the region of space in which the probability of finding an electron is $10%$, the probability of finding an electron inside the second contour from the nucleus is $20%$, inside the third one - $≈30%$, 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 determine simultaneously and exactly the energy and location of the electron. The more accurately 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 electron detection probability region has no clear boundaries. However, it is possible to single out the space where the probability of finding an electron is maximum.

The space around the atomic nucleus, in which the electron is most likely to be found, is called the 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. According to the form, $4$ of currently known types of orbitals are distinguished, which are denoted by the Latin letters $s, p, d$ and $f$. A graphic representation of some forms of electronic orbitals is shown in the figure.

The most important characteristic of the motion of an electron in a certain orbit is the energy of its connection with the nucleus. Electrons with similar energy values ​​form a single electronic layer, or energy level. Energy levels are numbered starting from the nucleus: $1, 2, 3, 4, 5, 6$ and $7$.

An integer $n$ denoting the number of the energy level is called the principal quantum number.

It characterizes the energy of electrons occupying a given energy level. The electrons of the first energy level, closest to the nucleus, have the lowest energy. Compared with the electrons of the first level, the electrons of the subsequent levels are characterized by large stock energy. Consequently, the electrons of the outer level are the least strongly bound to the nucleus of the atom.

The number of energy levels (electronic layers) in an atom is equal to the number of the period in the system of D. I. Mendeleev, to which the chemical element belongs: the atoms of the elements of the first period have one energy level; the second period - two; seventh period - seven.

The largest number of electrons in the 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: the first energy level closest to the nucleus can contain no more than two electrons; on the second - no more than $8$; on the third - no more than $18$; on the fourth - no more than $32$. And how, in turn, are the energy levels (electronic layers) arranged?

Starting from the second energy level $(n = 2)$, each of the levels is subdivided into sublevels (sublayers), which differ somewhat from each other in the binding energy with the nucleus.

The number of sublevels is equal to the value of the main quantum number: the first energy level has one sub level; the second - two; third - three; the fourth is four. Sublevels, in turn, are formed by orbitals.

Each value of $n$ corresponds to the number of orbitals equal to $n^2$. According to the data presented in the table, it is possible to trace the relationship between the principal quantum number $n$ and the number of sublevels, the type and number of orbitals, and the maximum number of electrons per sublevel and level.

Principal quantum number, types and number of orbitals, maximum number of electrons at sublevels and levels.

It is customary to designate sublevels in Latin letters, as well as the shape of the orbitals of which they consist: $s, p, d, f$. 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 for 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.

atom nucleus

But not only electrons are part of atoms. Physicist Henri Becquerel discovered that a natural mineral containing uranium salt also emits unknown radiation, illuminating photographic films that are closed from light. This phenomenon has been called radioactivity.

There are three types of radioactive rays:

1. $α$-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;
2. $β$-rays are a stream of electrons;
3. $γ$-rays are electromagnetic waves with a negligible mass that do not carry an electric charge.

Consequently, the atom has a complex structure - it consists of a positively charged nucleus and electrons.

How is the atom arranged?

In 1910 in Cambridge, near London, Ernest Rutherford with 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, confirming, it would seem, 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 running into some kind of obstacle.

By placing the 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, that area in which there are electrons that have a negative charge. If we apply a figurative comparison, then the entire volume of the atom can be likened to the stadium in Luzhniki, and the nucleus - soccer ball located in the center of the field.

An atom of any chemical element is comparable to a tiny solar system. Therefore, such a 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 particles of two types - protons and neutrons.

Protons have a charge equal to the charge of electrons, but opposite in sign $(+1)$, and a mass equal to the mass of a hydrogen atom (it is accepted in chemistry as a unit). Protons are denoted by $↙(1)↖(1)p$ (or $р+$). 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 denoted by $↙(0)↖(1)n$ (or $n^0$).

Protons and neutrons are collectively 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:

Since the mass of the electron, which is negligible, can be neglected, it is obvious that the entire mass of the atom is concentrated in the nucleus. Electrons are denoted 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. And how to determine the number of neutrons?

As you know, the mass of an atom is the sum of the mass of protons and neutrons. Knowing the ordinal number of the element $(Z)$, i.e. the number of protons, and the mass number $(A)$, equal to the sum of the numbers of protons and neutrons, you can find the number of neutrons $(N)$ using the formula:

For example, the number of neutrons in an iron atom is:

$56 – 26 = 30$.

The table shows 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- the same and topos- place, means "occupying one place" (cell) in the Periodic system of elements.

Chemical elements found in nature are a mixture of isotopes. Thus, carbon has three isotopes with a mass of $12, 13, 14$; oxygen - three isotopes with a mass of $16, 17, 18$, etc.

Usually given in the Periodic system, the relative atomic mass of a chemical element is the average value of the atomic masses of a natural mixture of isotopes of a given element, taking into account their relative 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 $25%$); 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 the isotopes of most chemical elements, such as potassium, argon:

$↖(39)↙(19)(K)$ and $↖(40)↙(19)(K)$, $↖(39)↙(18)(Ar)$ and $↖(40)↙(18 )(Ar)$

However, hydrogen isotopes differ greatly in properties due to the dramatic fold increase in their relative atomic mass; they were even given individual names and chemical signs: 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 chemical element.

A chemical element is a collection of atoms with the same nuclear charge.

The structure of the electron shells of atoms of the elements of the first four periods

Consider the mapping of the electronic configurations of the atoms of the elements by the periods of the system of D. I. Mendeleev.

Elements of the first period.

Schemes of the electronic structure of atoms show the distribution of electrons over electronic layers (energy levels).

The electronic formulas of atoms show the distribution of electrons over energy levels and sublevels.

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

In a helium atom, the first electron layer is complete - it has $2$ electrons.

Hydrogen and helium are $s$-elements, these atoms have $s$-orbitals filled with electrons.

Elements of the second period.

For all elements of the second period, the first electron layer is filled, and the electrons fill the $s-$ and $p$ orbitals of the second electron layer in accordance with the principle of least energy (first $s$, and then $p$) and the rules of Pauli and Hund.

In the neon atom, the second electron layer is complete - it has $8$ 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 3s-, 3p- and 3d-sublevels.

The structure of the electron shells of atoms of the elements of the third period.

A $3.5$-electron orbital is completed at the magnesium atom. $Na$ and $Mg$ are $s$-elements.

For aluminum and subsequent elements, the $3d$ sublevel is filled with electrons.

$↙(18)(Ar)$ Argon

$1s^2(2)s^2(2)p^6(3)s^2(3)p^6$

In an argon atom, the outer layer (the third electron layer) has $8$ electrons. 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 $3d$-orbitals left unfilled.

All elements from $Al$ to $Ar$ - $p$ -elements.

$s-$ and $r$ -elements form main subgroups in the Periodic system.

Elements of the fourth period.

Potassium and calcium atoms have a fourth electron layer, the $4s$-sublevel is filled, because it has less energy than the $3d$-sublevel. To simplify the graphical electronic formulas of the atoms of the elements of the fourth period:

1. we denote conditionally the graphic electronic formula of argon as follows: $Ar$;
2. we will not depict the sublevels that are not filled for 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 pre-external electron layer is filled, they are referred to transition elements.

Pay attention to the structure of the electron shells of chromium and copper atoms. A "failure" of one electron from the $4s-$ to the $3d$ sublevel occurs in them, 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	Diagram of the electronic structure	Electronic formula	Graphic 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 $1s^2(2)s^2(2)p ^6(3)p^6(3)d^5(4)s^1$
$↙(29)(Сu)$ Chromium		$1s^2(2)s^2(2)p^6(3)p^6(4)s^1(3)d^(10)$ 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 the $3s, 3p$ and $3d$ sublevels are filled in it, in total there are $18$ of electrons on them.

In the elements following zinc, the fourth electron layer, the $4p$-sublevel, continues to be filled. Elements from $Ga$ to $Kr$ - $r$ -elements.

The outer (fourth) layer of a krypton atom is completed, it has $8$ of electrons. But just in the fourth electron layer, as you know, there can be $32$ of electrons; the krypton atom still has $4d-$ and $4f$-sublevels unfilled.

The elements of the fifth period is coming filling in the sublevels in the following order: $5s → 4d → 5p$. And there are also exceptions related to the "failure" of electrons, for $↙(41)Nb$, $↙(42)Mo$, $↙(44)Ru$, $↙(45)Rh$, $↙(46) Pd$, $↙(47)Ag$. $f$ appear in the sixth and seventh periods -elements, i.e. elements whose $4f-$ and $5f$-sublevels of the third outside electronic layer are being filled, respectively.

$4f$ -elements called lanthanides.

$5f$ -elements called actinides.

The order of filling of electronic sublevels in the atoms of elements of the sixth period: $↙(55)Cs$ and $↙(56)Ba$ - $6s$-elements; $↙(57)La ... 6s^(2)5d^(1)$ - $5d$-element; $↙(58)Ce$ – $↙(71)Lu - 4f$-elements; $↙(72)Hf$ – $↙(80)Hg - 5d$-elements; $↙(81)Т1$ – $↙(86)Rn - 6d$-elements. But even here there are elements in which the order of filling of electron 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 electronic families, or blocks:

1. $s$ -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;
2. $r$ -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;
3. $d$ -elements; the $d$-sublevel of the preexternal level of the atom is filled with electrons; $d$-elements include elements of secondary subgroups of groups I–VIII, i.e. elements of intercalated decades of large periods located between $s-$ and $p-$elements. They are also called transition elements;
4. $f$ -elements;$f-$sublevel of the third level of the atom outside is filled with electrons; these include lanthanides and actinides.

The electronic configuration of the atom. Ground and excited states of atoms

The Swiss physicist W. Pauli in $1925$ established that An atom can have at most two electrons in one orbital. having opposite (antiparallel) spins (translated from English as a spindle), i.e. possessing such properties that can be conditionally imagined as the rotation of an electron around its imaginary axis clockwise or counterclockwise. This principle is called the Pauli principle.

If there is one electron in an orbital, then it is called unpaired, if two, then this paired electrons, i.e. electrons with opposite spins.

The figure shows a diagram of the division of energy levels into sublevels.

$s-$ Orbital, as you already know, has a spherical shape. The hydrogen atom electron $(n = 1)$ is located on this orbital and is unpaired. According to this his 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 (orbital type), and the number that is written to the right of the letter (as an exponent) shows the number of electrons in the sublevel.

For a helium atom He, which has two paired electrons in the same $s-$orbital, this formula is: $1s^2$. The electron shell of the helium atom is complete and very stable. Helium is a noble gas. The second energy level $(n = 2)$ has four orbitals, one $s$ and three $p$. Second-level $s$-orbital electrons ($2s$-orbitals) have a 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 amount of electron energy on it and, therefore, with a corresponding diameter, growing as the value of $n$.$s-$Orbital increases, as you already know , has a spherical shape. The hydrogen atom electron $(n = 1)$ is located on 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 (orbital type), and the number that is written to the right of the letter (as an exponent) shows the number of electrons in the sublevel.

For a helium atom $He$, which has two paired electrons in the same $s-$orbital, this formula is: $1s^2$. The electron shell of the helium atom is complete and very stable. Helium is a noble gas. The second energy level $(n = 2)$ has four orbitals, one $s$ and three $p$. Electrons of $s-$orbitals of the second level ($2s$-orbitals) have a 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 amount of electron energy on it and, therefore, with a corresponding diameter, growing as the value of $n$ increases.

$r-$ Orbital It has the shape of a dumbbell, or volume eight. All three $p$-orbitals are located in the atom mutually perpendicularly along the spatial coordinates drawn through the nucleus of the atom. It should be emphasized again that each energy level (electronic layer), starting from $n= 2$, has three $p$-orbitals. As the value of $n$ increases, the 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 weaker bound to the atomic nucleus, so a lithium atom can easily give it away (as you probably remember, this process is called oxidation), turning into a lithium ion $Li^+$.

In the beryllium atom Be, the fourth electron is also placed 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.

The fifth electron of the boron atom occupies the $2p$-orbital: $1s^(2)2s^(2)2p^(1)$. Next, the $2p$-orbitals of the $C, N, O, F$ atoms are filled, which ends with the neon noble gas: $1s^(2)2s^(2)2p^(6)$.

For elements of the third period, $3s-$ 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 above full electronic formulas, 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 respectively $4s-$ and $5s$-orbitals: $↙(19)K 2, 8, 8, 1;$ $↙(38)Sr 2, 8, 18, 8, 2$. Starting from the third element of each large period, the next ten electrons will go to the previous $3d-$ and $4d-$orbitals, respectively (for elements of secondary 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 (respectively $4p-$ and $5p-$) $p-$sublevel will start 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, as follows: 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 $14$ of electrons will enter the third energy level from the outside, the $4f$ and $5f$ orbitals of the lantonides and actinides, respectively: $↙(64)Gd 2, 8, 18, 25, 9, 2;$ $↙(92 )U 2, 8, 18, 32, 21, 9, 2$.

Then the second energy level from the outside ($d$-sublevel) will begin to build up again for the elements of side subgroups: $↙(73)Ta 2, 8, 18, 32, 11, 2;$ $↙(104)Rf 2, 8, 18 , 32, 32, 10, 2$. And, finally, only after the $d$-sublevel is completely filled with ten electrons, the $p$-sublevel will be filled again: $↙(86)Rn 2, 8, 18, 32, 18, 8$.

Very often, the structure of the electron shells of atoms is depicted using energy or quantum cells - they write down the so-called graphic electronic formulas. For this record, the following notation is used: each quantum cell is denoted by a cell that corresponds to one orbital; each electron is indicated by an arrow corresponding to the direction of the spin. When writing a graphical electronic formula, two rules should be remembered: Pauli principle, according to which a cell (orbital) can have no more than two electrons, 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 already be oppositely directed.