Electronic structure of atoms of chemical elements. Electronic configuration of an atom
Algorithm for compiling the electronic formula of an element:
1. Determine the number of electrons in an atom using the Periodic Table of Chemical Elements D.I. Mendeleev.
2. By the number of the period in which the element is located, determine the number of energy levels; the number of electrons in the last electronic level corresponds to the group number.
3. Divide the levels into sublevels and orbitals and fill them with electrons in accordance with the rules for filling orbitals:
It must be remembered that the first level has a maximum of 2 electrons. 1s2, on the second - a maximum of 8 (two s and six R: 2s 2 2p 6), on the third - a maximum of 18 (two s, six p, and ten d: 3s 2 3p 6 3d 10).
- Principal quantum number n should be minimal.
- Filled in first s- sublevel, then p-, d-b f- sublevels.
- Electrons fill orbitals in ascending order of orbital energy (Klechkovsky's rule).
- Within the sublevel, electrons first occupy free orbitals one at a time, and only after that they form pairs (Hund's rule).
- There cannot be more than two electrons in one orbital (Pauli principle).
Examples.
1. Compose the electronic formula of nitrogen. Nitrogen is number 7 on the periodic table.
2. Compose the electronic formula of argon. In the periodic table, argon is at number 18.
1s 2 2s 2 2p 6 3s 2 3p 6.
3. Compose the electronic formula of chromium. In the periodic table, chromium is number 24.
1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 5
Energy diagram of zinc.
4. Compose the electronic formula of zinc. In the periodic table, zinc is number 30.
1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10
Note that part of the electronic formula, namely 1s 2 2s 2 2p 6 3s 2 3p 6 is the electronic formula of argon.
The electronic formula of zinc can be represented as.
You need to enable JavaScript to run this app.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 find out where and how electrons are located, get acquainted with quantum numbers and be able to build 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 constructor: there are a certain number of parts, they differ from each other, but two parts of the same type are exactly the same. But this constructor 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, next to sodium into gas, and being next to iron completely turns it into rust. To answer the question why this happens and to 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 revolving around it, the nucleus consists of protons and neutrons. In the neutral state, each atom has the same number of electrons as the number of protons in its nucleus. The number of protons was indicated by the element's serial number, for example, sulfur has 16 protons - the 16th element of the periodic system. Gold has 79 protons - the 79th element of the periodic table. Accordingly, there are 16 electrons in sulfur in the neutral state, and 79 electrons in gold.
Where to look for an electron?
Observing the behavior of an electron, certain patterns were derived, they are described by quantum numbers, there are four of them 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", the orbital is the wave function of the electron, roughly - this is the area in which the electron spends 90% of the time.
N - level
L - shell
M l - orbital number
M s - the first or second electron in the orbital
Orbital quantum number l
As a result of the study of the electron cloud, it was found that depending on the level of energy, the cloud takes four main forms: a ball, dumbbells and the other two, more complex. In ascending order of energy, these forms are called s-, p-, d- and f-shells. 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 on which the orbitals are located. The orbital quantum number for s, p, d and f orbitals, respectively, takes the values 0,1,2 or 3.
On the s-shell one orbital (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 (L=3) on the f-shell - fourteen electrons
Magnetic quantum number m l
There are three orbitals on the p-shell, they are denoted 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 for each orbital, and then its pair is added to each.
Consider a 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. Only two electrons with opposite spins can be on the same energy sublevel. The spin quantum number is denoted m s
Principal quantum number n
The main quantum number is the energy level, at the moment seven energy levels are known, each is denoted 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, and so on.
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, let's take the first electron, the lowest energy level is N=1, one shell is located on the first level, 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.
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 next levels are characterized by a large amount of 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), slightly different from each other by 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.
| Energy level $(n)$ | Number of sublevels equal to $n$ | Orbital type | Number of orbitals | Maximum number of electrons | ||
| in sublevel | in level equal to $n^2$ | in sublevel | at a level equal to $n^2$ | |||
| $K(n=1)$ | $1$ | $1s$ | $1$ | $1$ | $2$ | $2$ |
| $L(n=2)$ | $2$ | $2s$ | $1$ | $4$ | $2$ | $8$ |
| $2p$ | $3$ | $6$ | ||||
| $M(n=3)$ | $3$ | $3s$ | $1$ | $9$ | $2$ | $18$ |
| $3p$ | $3$ | $6$ | ||||
| $3d$ | $5$ | $10$ | ||||
| $N(n=4)$ | $4$ | $4s$ | $1$ | $16$ | $2$ | $32$ |
| $4p$ | $3$ | $6$ | ||||
| $4d$ | $5$ | $10$ | ||||
| $4f$ | $7$ | $14$ | ||||
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:
- $α$-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 are a stream of electrons;
- $γ$-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 Luzhniki stadium, 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, 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 it is possible to 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 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 $s$, 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:
- we denote conditionally the graphic electronic formula of argon as follows: $Ar$;
- 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. In them, one electron "falls" 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 | 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 are filling the sublevels in the following order: $5s → 4d → 5р$. 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 here, too, 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:
- $s$ -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;
- $r$ -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;
- $d$ -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;
- $f$ -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 sublevel (orbital type) is denoted by the Latin letter, 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 sublevel (orbital type) is denoted by the Latin letter, 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 less 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 $14$ 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 at the same time have the same spin value, and only then pair, but the spins, according to the Pauli principle, will already be oppositely directed.
The concept of "atom" has been familiar to mankind since ancient Greece. According to the saying of the ancient philosophers, the atom is the smallest particle that is part of the substance.
The electronic structure of the atom
An atom consists of a positively charged nucleus containing protons and neutrons. Electrons move in orbits around the nucleus, each of which can be characterized by a set of four quantum numbers: principal (n), orbital (l), magnetic (m l) and spin (ms or s).
The principal quantum number determines the energy of the electron and the size of the electron clouds. The energy of an electron mainly depends on the distance of the electron from the nucleus: the closer the electron is to the nucleus, the lower its energy. In other words, the main quantum number determines the location of an electron on a particular energy level (quantum layer). The principal quantum number has the values of a series of integers from 1 to infinity.
The orbital quantum number characterizes the shape of the electron cloud. The different shape of electron clouds causes a change in the energy of electrons within the same energy level, i.e. splitting it into energy sublevels. The orbital quantum number can have values from zero to (n-1), in total n values. Energy sublevels are denoted by letters:
The magnetic quantum number shows the orientation of the orbital in space. It accepts any integer value from (+l) to (-l), including zero. The number of possible values of the magnetic quantum number is (2l+1).
An electron moving in the field of the nucleus of an atom, in addition to the orbital angular momentum, also has its own angular momentum, which characterizes its spindle-shaped rotation around its own axis. This property of an electron is called spin. The value and orientation of the spin is characterized by the spin quantum number, which can take the values (+1/2) and (-1/2). The positive and negative values of a spin are related to its direction.
Before all of the above became known and confirmed experimentally, there were several models of the structure of the atom. One of the first models of the structure of the atom was proposed by E. Rutherford, who, in experiments on the scattering of α-particles, showed that almost the entire mass of the atom is concentrated in a very small volume - a positively charged nucleus. According to his model, electrons move around the nucleus at a sufficiently large distance, and their number is such that, on the whole, the atom is electrically neutral.
Rutherford's model of the structure of the atom was developed by N. Bohr, who in his research also combined Einstein's teachings on light quanta and Planck's quantum theory of radiation. Louis de Broglie and Schrödinger completed what they started and presented to the world a modern model of the structure of the atom of a chemical element.
Examples of problem solving
EXAMPLE 1
| Exercise | Indicate the number of protons and neutrons that are contained in the nuclei of nitrogen (atomic number 14), silicon (atomic number 28) and barium (atomic number 137). |
| Solution | The number of protons in the nucleus of an atom of a chemical element is determined by its serial number in the Periodic Table, and the number of neutrons is the difference between the mass number (M) and the nuclear charge (Z). Nitrogen: n(N)=M-Z=14-7=7. Silicon: n(Si) \u003d M -Z \u003d 28-14 \u003d 14. Barium: n (Ba) \u003d M -Z \u003d 137-56 \u003d 81. |
| Answer | The number of protons in the nitrogen nucleus is 7, neutrons - 7; in the nucleus of a flint atom there are 14 protons, 14 neutrons; in the nucleus of a barium atom, there are 56 protons and 81 neutrons. |
EXAMPLE 2
| Exercise | Arrange the energy sublevels in the sequence of their filling with electrons: a) 3p, 3d, 4s, 4p; b) 4d , 5s, 5p, 6s; c) 4f , 5s , 6p; 4d , 6s; d) 5d, 6s, 6p, 7s, 4f . |
|
| Solution | The energy sublevels are filled with electrons in accordance with the Klechkovsky rules. A prerequisite is the minimum value of the sum of the principal and orbital quantum numbers. The s-sublevel is characterized by the number 0, p - 1, d - 2 and f-3. The second condition is that the sublevel with the lowest value of the main quantum number is filled first. | |
| Answer | a) Orbitals 3p, 3d, 4s, 4p will correspond to the numbers 4, 5, 4 and 5. Therefore, the filling with electrons will occur in the following sequence: 3p, 4s, 3d, 4p. b) Orbitals 4d , 5s, 5p, 6s will correspond to the numbers 7, 5, 6 and 6. Therefore, filling with electrons will occur in the following sequence: 5s, 5p, 6s, 4d. c) Orbitals 4f , 5s , 6p; 4d , 6s will correspond to the numbers 7, 5, 76 and 6. Therefore, filling with electrons will occur in the following sequence: 5s, 4d , 6s, 4f, 6p. d) Orbitals 5d, 6s, 6p, 7s, 4f will correspond to the numbers 7, 6, 7, 7 and 7. Therefore, the filling with electrons will occur in the following sequence: 6s, 4f, 5d, 6p, 7s. |
Electrons
The concept of an atom originated in the ancient world to denote the particles of matter. In Greek, atom means "indivisible".
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 a unit (-1). Thomson even managed to determine the speed of the electron (the speed of an electron in orbit is inversely proportional to the orbit number n. The radii of the orbits grow in proportion to the square of the orbit number. In the first orbit of the hydrogen atom (n=1; Z=1), the speed is ≈ 2.2 106 m / c, that is, about a hundred times less than the speed of light c=3 108 m/s.) and the mass of an electron (it is almost 2000 times less than the mass of a hydrogen atom).
The state of electrons in an atom
The state of an electron in an atom is a set of information about the energy of a particular electron and the space in which it is located. An electron in an atom does not have a trajectory of motion, i.e., one can only speak of the probability of 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 points. Overlaying countless such photographs would result in a picture of an electron cloud with the highest density where there will be most of these points.
The space around the atomic nucleus, in which the electron is most likely to be found, is called the orbital. It contains approximately 90% e-cloud, and this means that about 90% of the time the electron is in this part of space. Distinguished by shape 4 currently known types of orbitals, which are denoted by 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 electron 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 main 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 next levels will be characterized by a large amount of energy. Consequently, the electrons of the outer level are the least strongly bound to the nucleus of the atom.
The largest number of electrons in the energy level is determined by the formula:
N = 2n2,
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.
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 sublevel; the second - two; third - three; fourth - four sublevels. Sublevels, in turn, are formed by orbitals. Each valuen corresponds to the number of orbitals equal to n.
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.
Protons and neutrons
An atom of any chemical element is comparable to a tiny solar system. Therefore, such a model of the atom, proposed by E. Rutherford, is called planetary.
The atomic nucleus, in which the entire mass of the atom is concentrated, consists of particles of two types - protons and neutrons.
Protons have a charge equal to the charge of electrons, but opposite in sign (+1), and a mass equal to the mass of a hydrogen atom (it is accepted in chemistry as a unit). Neutrons carry no charge, they are neutral and have a mass equal to that of a proton.
Protons and neutrons are collectively called nucleons (from the Latin nucleus - nucleus). The sum of the number of protons and neutrons in an atom is called the mass number. For example, the mass number of an aluminum atom:
13 + 14 = 27
number of protons 13, number of neutrons 14, mass number 27
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 represent e - .
Because the atom electrically neutral, it is also obvious that the number of protons and electrons in an atom is the same. It is equal to the serial number of the chemical element assigned to it in the Periodic system. The mass of an atom is made up 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, you can find the number of neutrons (N) using the formula:
N=A-Z
For example, the number of neutrons in an iron atom is:
56 — 26 = 30
isotopes
Varieties of atoms of the same element that have the same nuclear charge but different mass numbers are called isotopes. Chemical elements found in nature are a mixture of isotopes. So, carbon has three isotopes with a mass of 12, 13, 14; oxygen - three isotopes with a mass of 16, 17, 18, etc. The relative atomic mass of a chemical element usually given in the Periodic System is the average value of the atomic masses of a natural mixture of isotopes of a given element, taking into account their relative content in nature. The chemical properties of the isotopes of most chemical elements are exactly the same. However, hydrogen isotopes differ greatly in properties due to the dramatic fold increase in their relative atomic mass; they have even been given individual names and chemical symbols.
Elements of the first period
Scheme of the electronic structure of the hydrogen atom:
Schemes of the electronic structure of atoms show the distribution of electrons over electronic layers (energy levels).
The graphical electronic formula of the hydrogen atom (shows 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 orbits.
In a helium atom, the first electron layer is completed - it has 2 electrons. Hydrogen and helium are s-elements; for these atoms, the s-orbital is filled with electrons.
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 completed - it has 8 electrons.
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.
A 3s electron orbital is completed at the magnesium atom. Na and Mg are s-elements.
For aluminum and subsequent elements, the 3p sublevel is filled with electrons.
The elements of the third period have unfilled 3d orbitals.
All elements from Al to Ar are p-elements. s- and p-elements form the main subgroups in the Periodic system.
Elements of the fourth - seventh periods
A fourth electron layer appears at the potassium and calcium atoms, the 4s sublevel is filled, since it has less energy than the 3d sublevel.
K, Ca - s-elements included in the main subgroups. For atoms from Sc to Zn, the 3d sublevel is filled with electrons. These are 3d elements. They are included in the secondary subgroups, they have a pre-external electron layer filled, they are referred to as transition elements.
Pay attention to the structure of the electron shells of chromium and copper atoms. In them, a “failure” of one electron from the 4s- to the 3d-sublevel occurs, which is explained by the greater energy stability of the resulting electronic configurations 3d 5 and 3d 10:
In the zinc atom, the third electron layer is completed - all the 3s, 3p and 3d sublevels are filled in it, in total there are 18 electrons on them. In the elements following zinc, the fourth electron layer continues to be filled, the 4p sublevel.
Elements from Ga to Kr are p-elements.
The outer layer (fourth) of the krypton atom is complete and has 8 electrons. But there can only be 32 electrons in the fourth electron layer; the 4d- and 4f-sublevels of the krypton atom still remain unfilled. The elements of the fifth period are filling the sub-levels in the following order: 5s - 4d - 5p. And there are also exceptions related to " failure» electrons, y 41 Nb, 42 Mo, 44 Ru, 45 Rh, 46 Pd, 47 Ag.
In the sixth and seventh periods, f-elements appear, i.e., elements in which the 4f- and 5f-sublevels of the third outer electronic layer are filled, respectively.
4f elements are called lanthanides.
5f elements are 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 x - 5d element; 58 Ce - 71 Lu - 4f elements; 72 Hf - 80 Hg - 5d elements; 81 T1 - 86 Rn - 6d elements. But even here 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 are divided into four electronic 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 III-VIII groups.
- 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 intercalary decades of large periods located between s- and p-elements. They are also called transition elements.
- f-elements. The f-sublevel of the third outside level of the atom is filled with electrons; these include the lanthanides and antinoids.
The Swiss physicist W. Pauli in 1925 established that in an atom in one orbital there can be no more than two electrons having opposite (antiparallel) spins (translated from English - “spindle”), i.e. having such properties that can be conditionally imagined as the rotation of an electron around its imaginary axis: clockwise or counterclockwise.
This principle is called Pauli principle. If there is one electron in the orbital, then it is called unpaired, if there are two, then these are paired electrons, that is, electrons with opposite spins. The figure shows a diagram of the division of energy levels into sublevels and the order in which they are filled.
Very often, the structure of the 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 and F. Hund's rule, according to which electrons occupy free cells, first one at a time and at the same time have the same spin value, and only then pair, but the spins, according to the Pauli principle, will already be oppositely directed.
Hund's rule and Pauli's principle
Hund's rule- the rule of quantum chemistry, which determines the order of filling the orbitals of a certain sublayer and is formulated as follows: the total value of the spin quantum number of electrons of this sublayer should be maximum. Formulated by Friedrich Hund in 1925.
This means that in each of the orbitals of the sublayer, one electron is first filled, and only after the unfilled orbitals are exhausted, a second electron is added to this orbital. In this case, there are two electrons with half-integer spins of the opposite sign in one orbital, which pair (form a two-electron cloud) and, as a result, the total spin of the orbital becomes equal to zero.
Other wording: Below in energy lies the atomic term for which two conditions are satisfied.
- Multiplicity is maximum
- When the multiplicities coincide, the total orbital momentum L is maximum.
Let's analyze this rule using the example of filling the orbitals of the p-sublevel p- elements of the second period (that is, from boron to neon (in the diagram below, horizontal lines indicate orbitals, vertical arrows indicate electrons, and the direction of the arrow indicates the orientation of the spin).
Klechkovsky's rule
Klechkovsky's rule - as the total number of electrons in atoms increases (with an increase in the charges of their nuclei, or the ordinal numbers of chemical elements), atomic orbitals are populated in such a way that the appearance of electrons in higher-energy orbitals depends only on the principal quantum number n and does not depend on all other quantum numbers. numbers, including those from l. Physically, this means that in a hydrogen-like atom (in the absence of interelectron repulsion) the orbital energy of an electron is determined only by the spatial remoteness of the electron charge density from the nucleus and does not depend on the features of its motion in the field of the nucleus.
Klechkovsky's empirical rule and the sequence of sequences of a somewhat contradictory real energy sequence of atomic orbitals arising from it only in two cases of the same type: for atoms Cr, Cu, Nb, Mo, Ru, Rh, Pd, Ag, Pt, Au, there is a “failure” of an electron with s - sublevel of the outer layer to the d-sublevel of the previous layer, which leads to an energetically more stable state of the atom, namely: after filling the orbital 6 with two electrons s
