The structure of the atom, chemical bond, valency and structure of molecules. The structure of atoms of chemical elements
Documentary educational films. Series "Physics".
Atom (from the Greek atomos - indivisible) - a single-nuclear, chemically indivisible particle chemical element, the carrier of the properties of matter. Substances are made up of atoms. The atom itself consists of a positively charged nucleus and a negatively charged electron cloud. In general, the atom is electrically neutral. The size of an atom is completely determined by the size of its electron cloud, since the size of the nucleus is negligible compared to the size of the electron cloud. The nucleus consists of Z positively charged protons (the proton charge corresponds to +1 in conventional units) and N neutrons that do not carry a charge (protons and neutrons are called nucleons). Thus, the charge of the nucleus is determined only by the number of protons and is equal to the serial number of the element in the periodic table. The positive charge of the nucleus is compensated by negatively charged electrons (electron charge -1 in arbitrary units), which form an electron cloud. The number of electrons is equal to the number of protons. The masses of protons and neutrons are equal (1 and 1 amu, respectively).
The mass of an atom is determined by the mass of its nucleus, since the mass of an electron is approximately 1850 times less than the mass of a proton and a neutron and is rarely taken into account in calculations. The number of neutrons can be found by the difference between the mass of an atom and the number of protons (N=A-Z). The type of atoms of any chemical element with a nucleus consisting of strictly a certain number protons (Z) and neutrons (N) is called a nuclide.
Before studying the properties of an electron and the rules for the formation of electronic levels, it is necessary to touch upon the history of the formation of ideas about the structure of an atom. We will not consider the full history of the formation of the atomic structure, but will dwell only on the most relevant and most "correct" ideas that can most clearly show how the electrons are located in the atom. The presence of atoms as the elementary constituents of matter was first suggested by the ancient Greek philosophers. After that, the history of the structure of the atom went through a difficult path and different ideas, such as the indivisibility of the atom, the Thomson model of the atom, and others. The model of the atom proposed by Ernest Rutherford in 1911 turned out to be the closest. He compared the atom to solar system, where the nucleus of an atom acted as the sun, and the electrons moved around it like planets. Placing electrons in stationary orbits was a very important step in understanding the structure of the atom. However, such planetary model structure of the atom was in conflict with classical mechanics. The fact is that when an electron moved in orbit, it had to lose potential energy and eventually "fall" onto the nucleus and the atom had to cease to exist. Such a paradox was eliminated by the introduction of postulates by Niels Bohr. According to these postulates, the electron moved in stationary orbits around the nucleus and under normal conditions did not absorb or emit energy. The postulates show that the laws of classical mechanics are not suitable for describing the atom. This model of the atom is called the Bohr-Rutherford model. continuation planetary structure atom is the quantum mechanical model of the atom, according to which we will consider the electron.
The electron is a quasi-particle showing corpuscular-wave dualism. It is both a particle (corpuscle) and a wave at the same time. The properties of a particle include the mass of an electron and its charge, and the wave properties - the ability to diffraction and interference. The relationship between the wave and corpuscular properties of an electron is reflected in the de Broglie equation.
(Lecture notes)
The structure of the atom. Introduction.
The object of study in chemistry is the chemical elements and their compounds. chemical element A group of atoms with the same positive charge is called. Atom is the smallest particle of a chemical element that retains it Chemical properties. Connecting with each other, atoms of one or different elements form more complex particles - molecules. A collection of atoms or molecules form chemicals. Each individual chemical substance is characterized by a set of individual physical properties, such as boiling and melting points, density, electrical and thermal conductivity, etc.
1. The structure of the atom and the Periodic system of elements
DI. Mendeleev.
Knowledge and understanding of the patterns of filling order Periodic system elements D.I. Mendeleev allows us to understand the following:
1. the physical essence of the existence in nature of certain elements,
2. the nature of the chemical valency of the element,
3. the ability and "ease" of an element to give or receive electrons when interacting with another element,
4. the nature of the chemical bonds that can form given element when interacting with other elements, the spatial structure of simple and complex molecules, etc., etc.
The structure of the atom.
An atom is a complex microsystem of elementary particles in motion and interacting with each other.
In the late 19th and early 20th centuries, it was found that atoms are composed of smaller particles: neutrons, protons and electrons. The last two particles are charged particles, the proton carries a positive charge, the electron is negative. Since the atoms of an element in the ground state are electrically neutral, this means that the number of protons in an atom of any element is equal to the number of electrons. The mass of atoms is determined by the sum of the masses of protons and neutrons, the number of which is equal to the difference between the mass of atoms and its serial number in the periodic system of D.I. Mendeleev.
In 1926, Schrodinger proposed to describe the motion of microparticles in the atom of an element using the wave equation he derived. When solving the Schrödinger wave equation for the hydrogen atom, three integer quantum numbers appear: n, ℓ And m ℓ , which characterize the state of an electron in three-dimensional space in the central field of the nucleus. quantum numbers n, ℓ And m ℓ take integer values. Wave function defined by three quantum numbers n, ℓ And m ℓ and obtained as a result of solving the Schrödinger equation is called an orbital. An orbital is a region of space in which an electron is most likely to be found. belonging to an atom of a chemical element. Thus, the solution of the Schrödinger equation for the hydrogen atom leads to the appearance of three quantum numbers, physical meaning which is that they characterize the three different kinds of orbitals that an atom can have. Let's take a closer look at each quantum number.
Principal quantum number n can take any positive integer values: n = 1,2,3,4,5,6,7… It characterizes the energy of the electronic level and the size of the electronic "cloud". It is characteristic that the number of the main quantum number coincides with the number of the period in which the given element is located.
Azimuthal or orbital quantum numberℓ can take integer values from ℓ = 0….up to n – 1 and determines the moment of electron motion, i.e. orbital shape. For various numerical values of ℓ, the following notation is used: ℓ = 0, 1, 2, 3, and are denoted by symbols s, p, d, f, respectively for ℓ = 0, 1, 2 and 3. In the periodic table of elements there are no elements with a spin number ℓ = 4.
Magnetic quantum numberm ℓ characterizes the spatial arrangement of electron orbitals and, consequently, the electromagnetic properties of the electron. It can take values from - ℓ to + ℓ , including zero.
The shape or, more precisely, the symmetry properties of atomic orbitals depend on quantum numbers ℓ And m ℓ . "electronic cloud", corresponding to s- orbitals has, has the shape of a ball (at the same time ℓ = 0).
Fig.1. 1s orbital
Orbitals defined by quantum numbers ℓ = 1 and m ℓ = -1, 0 and +1 are called p-orbitals. Since m ℓ has three different values, then the atom has three energetically equivalent p-orbitals (the main quantum number for them is the same and can have the value n = 2,3,4,5,6 or 7). p-Orbitals have axial symmetry and have the form of three-dimensional eights, oriented along the x, y and z axes in an external field (Fig. 1.2). Hence the origin of the symbols p x , p y and p z .
Fig.2. p x , p y and p z -orbitals
In addition, there are d- and f-atomic orbitals, for the first ℓ = 2 and m ℓ = -2, -1, 0, +1 and +2, i.e. five AO, for the second ℓ = 3 and m ℓ = -3, -2, -1, 0, +1, +2 and +3, i.e. 7 AO.
fourth quantum m s called the spin quantum number, was introduced to explain some subtle effects in the spectrum of the hydrogen atom by Goudsmit and Uhlenbeck in 1925. The spin of an electron is the angular momentum of a charged elementary particle of an electron, the orientation of which is quantized, i.e. strictly limited to certain angles. This orientation is determined by the value of the spin magnetic quantum number (s), which for an electron is ½ , therefore, for an electron, according to the quantization rules m s = ± ½. In this regard, to the set of three quantum numbers, one should add the quantum number m s . We emphasize once again that four quantum numbers determine the order in which Mendeleev's periodic table of elements is constructed and explain why there are only two elements in the first period, eight in the second and third, 18 in the fourth, and so on. However, in order to explain the structure of multielectron of atoms, the order in which electronic levels are filled as the positive charge of an atom increases, it is not enough to have an idea about the four quantum numbers that "govern" the behavior of electrons when filling electronic orbitals, but you need to know some more simple rules, namely, Pauli's principle, Gund's rule and Klechkovsky's rules.
According to the Pauli principle in the same quantum state, characterized by certain values of four quantum numbers, there cannot be more than one electron. This means that one electron can, in principle, be placed in any atomic orbital. Two electrons can be in the same atomic orbital only if they have different spin quantum numbers.
When filling three p-AOs, five d-AOs and seven f-AOs with electrons, one should be guided not only by the Pauli principle but also by the Hund rule: The filling of the orbitals of one subshell in the ground state occurs with electrons with the same spins.
When filling subshells (p, d, f) the absolute value of the sum of spins must be maximum.
Klechkovsky's rule. According to the Klechkovsky rule, when fillingd And forbital by electrons must be respectedprinciple of minimum energy. According to this principle, electrons in the ground state fill the orbits with minimum energy levels. The sublevel energy is determined by the sum of quantum numbersn + ℓ = E .
Klechkovsky's first rule: first fill those sublevels for whichn + ℓ = E minimal.
Klechkovsky's second rule: in case of equalityn + ℓ for several sublevels, the sublevel for whichn minimal .
Currently, 109 elements are known.
2. Ionization energy, electron affinity and electronegativity.
The most important characteristics of the electronic configuration of an atom are the ionization energy (EI) or ionization potential (IP) and the atom's electron affinity (SE). The ionization energy is the change in energy in the process of detachment of an electron from a free atom at 0 K: A = + + ē . The dependence of the ionization energy on the atomic number Z of the element, the size of the atomic radius has a pronounced periodic character.
Electron affinity (SE) is the change in energy that accompanies the addition of an electron to an isolated atom with the formation of a negative ion at 0 K: A + ē = A - (the atom and ion are in their ground states). In this case, the electron occupies the lowest free atomic orbital (LUAO) if the VZAO is occupied by two electrons. SE strongly depends on their orbital electronic configuration.
Changes in EI and SE correlate with changes in many properties of elements and their compounds, which is used to predict these properties from the values of EI and SE. Halogens have the highest absolute electron affinity. In each group of the periodic table of elements, the ionization potential or EI decreases with increasing element number, which is associated with an increase in atomic radius and with an increase in the number of electron layers, and which correlates well with an increase in the element's reducing power.
Table 1 of the Periodic Table of the Elements gives the values of EI and SE in eV/atom. Note that exact values SE are known only for a few atoms, their values are underlined in Table 1.
Table 1
The first ionization energy (EI), electron affinity (SE) and electronegativity χ) of atoms in the periodic table.
|
χ |
0.747 2. 1 0 0, 3 7 |
1,2 2 |
||||||||||||||||
|
χ |
0.54 1. 55 |
-0.3 1. 1 3 |
0.2 0. 91 |
1.2 5 |
-0. 1 0, 55 |
1.47 0. 59 |
3.45 0. 64 |
1 ,60 |
||||||||||
|
χ |
0. 7 4 1. 89 |
-0.3 1 . 3 1 1 . 6 0 |
0. 6 |
1.63 |
0.7 |
2.07 |
3.61 | |||||||||||
|
χ |
2.3 6 |
- 0 .6 |
1.26(α) |
-0.9 1 . 39 |
0. 18 |
1.2 |
0. 6 |
2.07 |
3.36 | |||||||||
|
χ |
2.4 8 |
-0.6 1 . 56 |
0. 2 |
2.2 | ||||||||||||||
|
χ |
2.6 7 |
2, 2 1 |
ABOUTs |
χ - Pauling electronegativity
r- atomic radius, (from "Laboratory and seminar classes in general and inorganic chemistry", N.S. Akhmetov, M.K. Azizova, L.I. Badygina)
Chemicals are the things that make up the world around us.
The properties of each chemical substance are divided into two types: these are chemical, which characterize its ability to form other substances, and physical, which are objectively observed and can be considered in isolation from chemical transformations. So, for example, the physical properties of a substance are its state of aggregation (solid, liquid or gaseous), thermal conductivity, heat capacity, solubility in various media (water, alcohol, etc.), density, color, taste, etc.
Transformations of some chemical substances into other substances are called chemical phenomena or chemical reactions. It should be noted that there are also physical phenomena, which, obviously, are accompanied by a change in some physical properties substances without being converted into other substances. Physical phenomena, for example, include the melting of ice, the freezing or evaporation of water, etc.
The fact that in the course of any process a chemical phenomenon takes place can be concluded by observing characteristics chemical reactions such as color change, precipitation, gas evolution, heat and/or light evolution.
So, for example, a conclusion about the course of chemical reactions can be made by observing:
The formation of sediment when boiling water, called scale in everyday life;
The release of heat and light during the burning of a fire;
Change the slice color fresh apple on air;
The formation of gas bubbles during the fermentation of dough, etc.
The smallest particles of matter, which in the process of chemical reactions practically do not undergo changes, but only in a new way are connected to each other, are called atoms.
The very idea of the existence of such units of matter arose in ancient greece in the minds of ancient philosophers, which actually explains the origin of the term "atom", since "atomos" literally translated from Greek means "indivisible".
However, contrary to the idea ancient Greek philosophers, atoms are not the absolute minimum of matter, i.e. themselves have a complex structure.
Each atom consists of the so-called subatomic particles - protons, neutrons and electrons, denoted respectively by the symbols p + , n o and e - . The superscript in the notation used indicates that the proton has a unit positive charge, the electron has a unit negative charge, and the neutron has no charge.
As for the qualitative structure of the atom, each atom has all the protons and neutrons concentrated in the so-called nucleus, around which the electrons form an electron shell.
The proton and neutron have practically the same masses, i.e. m p ≈ m n , and the electron mass is almost 2000 times less than the mass of each of them, i.e. m p / m e ≈ m n / m e ≈ 2000.
Since the fundamental property of an atom is its electrical neutrality, and the charge of one electron is equal to the charge of one proton, it can be concluded from this that the number of electrons in any atom is equal to the number of protons.
So, for example, the table below shows the possible composition of atoms:
The type of atoms with the same nuclear charge, i.e. from the same number protons in their nuclei are called a chemical element. Thus, from the table above, we can conclude that atom1 and atom2 belong to one chemical element, and atom3 and atom4 belong to another chemical element.
Each chemical element has its own name and individual symbol, which is read in a certain way. So, for example, the simplest chemical element, the atoms of which contain only one proton in the nucleus, has the name "hydrogen" and is denoted by the symbol "H", which is read as "ash", and the chemical element with a nuclear charge of +7 (i.e. containing 7 protons) - "nitrogen", has the symbol "N", which is read as "en".
As you can see from the table above, the atoms of one chemical element can differ in the number of neutrons in the nuclei.
Atoms belonging to the same chemical element, but having a different number of neutrons and, as a result, mass, are called isotopes.
So, for example, the chemical element hydrogen has three isotopes - 1 H, 2 H and 3 H. The indices 1, 2 and 3 above the H symbol mean the total number of neutrons and protons. Those. knowing that hydrogen is a chemical element, which is characterized by the fact that there is one proton in the nuclei of its atoms, we can conclude that there are no neutrons at all in the 1 H isotope (1-1 = 0), in the 2 H isotope - 1 neutron (2-1=1) and in the isotope 3 H - two neutrons (3-1=2). Since, as already mentioned, a neutron and a proton have the same masses, and the mass of an electron is negligible compared to them, this means that the 2 H isotope is almost twice as heavy as the 1 H isotope, and the 3 H isotope is even three times as heavy. . In connection with such a large spread in the masses of hydrogen isotopes, the 2 H and 3 H isotopes were even given separate individual names and symbols, which is not typical of any other chemical element. The 2 H isotope was named deuterium and given the symbol D, and the 3 H isotope was given the name tritium and given the symbol T.
If we take the mass of the proton and neutron as unity, and neglect the mass of the electron, in fact, the upper left index, in addition to the total number of protons and neutrons in the atom, can be considered its mass, and therefore this index is called the mass number and is denoted by the symbol A. Since the charge of the nucleus of any protons correspond to the atom, and the charge of each proton is conventionally considered equal to +1, the number of protons in the nucleus is called charge number(Z). Denoting the number of neutrons in an atom with the letter N, mathematically the relationship between mass number, charge number and number of neutrons can be expressed as:
According to modern concepts, the electron has a dual (particle-wave) nature. It has the properties of both a particle and a wave. Like a particle, an electron has a mass and a charge, but at the same time, the flow of electrons, like a wave, is characterized by the ability to diffraction.
To describe the state of an electron in an atom, representations are used quantum mechanics, according to which the electron does not have a specific trajectory of motion and can be located at any point in space, but with different probabilities.
The region of space around the nucleus where an electron is most likely to be found is called the atomic orbital.
An atomic orbital can have various form, size and orientation. An atomic orbital is also called an electron cloud.
Graphically, one atomic orbital is usually denoted as a square cell:
Quantum mechanics has an extremely complex mathematical apparatus, therefore, within the framework of a school chemistry course, only the consequences of quantum mechanical theory are considered.
According to these consequences, any atomic orbital and an electron located on it are completely characterized by 4 quantum numbers.
- The main quantum number - n - determines the total energy of an electron in a given orbital. The range of values of the main quantum number is all integers, i.e. n = 1,2,3,4, 5 etc.
- The orbital quantum number - l - characterizes the shape of the atomic orbital and can take any integer values from 0 to n-1, where n, recall, is the main quantum number.
Orbitals with l = 0 are called s-orbitals. s-orbitals are spherical and do not have a direction in space:
Orbitals with l = 1 are called p-orbitals. These orbitals have the shape of a three-dimensional figure eight, i.e. the shape obtained by rotating the figure eight around the axis of symmetry, and outwardly resemble a dumbbell:
Orbitals with l = 2 are called d-orbitals, and with l = 3 – f-orbitals. Their structure is much more complex.
3) The magnetic quantum number - m l - determines the spatial orientation of a particular atomic orbital and expresses the projection of the orbital angular momentum on the direction magnetic field. The magnetic quantum number m l corresponds to the orientation of the orbital relative to the direction of the external magnetic field strength vector and can take any integer values from –l to +l, including 0, i.e. total amount possible values equals (2l+1). So, for example, for l = 0 ml = 0 (one value), for l = 1 ml = -1, 0, +1 (three values), for l = 2 ml = -2, -1, 0, +1 , +2 (five values of the magnetic quantum number), etc.
So, for example, p-orbitals, i.e. orbitals with an orbital quantum number l = 1, having the shape of a “three-dimensional figure eight”, correspond to three values of the magnetic quantum number (-1, 0, +1), which, in turn, corresponds to three directions in space perpendicular to each other.
4) The spin quantum number (or simply spin) - m s - can be conditionally considered responsible for the direction of rotation of an electron in an atom, it can take on values. Electrons with different spins are indicated by vertical arrows pointing in different directions: ↓ and .
The set of all orbitals in an atom that have the same value of the principal quantum number is called the energy level or electron shell. Any arbitrary energy level with some number n consists of n 2 orbitals.
Many orbitals with the same values principal quantum number and orbital quantum number represents the energy sublevel.
Each energy level, which corresponds to the main quantum number n, contains n sublevels. In turn, each energy sublevel with an orbital quantum number l consists of (2l+1) orbitals. Thus, the s-sublayer consists of one s-orbital, the p-sublayer - three p-orbitals, the d-sublayer - five d-orbitals, and the f-sublayer - seven f-orbitals. Since, as already mentioned, one atomic orbital is often denoted by one square cell, the s-, p-, d- and f-sublevels can be graphically depicted as follows:
Each orbital corresponds to an individual strictly defined set of three quantum numbers n, l and m l .
The distribution of electrons in orbitals is called the electronic configuration.
The filling of atomic orbitals with electrons occurs in accordance with three conditions:
- The principle of minimum energy: Electrons fill orbitals starting from the lowest energy sublevel. The sequence of sublevels in order of increasing energy is as follows: 1s<2s<2p<3s<3p<4s≤3d<4p<5s≤4d<5p<6s…;
In order to make it easier to remember this sequence of filling electronic sublevels, the following graphic illustration is very convenient:
- Pauli principle: Each orbital can hold at most two electrons.
If there is one electron in the orbital, then it is called unpaired, and if there are two, then they are called an electron pair.
- Hund's rule: the most stable state of an atom is one in which, within one sublevel, the atom has the maximum possible number of unpaired electrons. This most stable state of the atom is called the ground state.
In fact, the above means that, for example, the placement of the 1st, 2nd, 3rd and 4th electrons on three orbitals of the p-sublevel will be carried out as follows:
The filling of atomic orbitals from hydrogen, which has a charge number of 1, to krypton (Kr) with a charge number of 36, will be carried out as follows:
A similar representation of the order in which atomic orbitals are filled is called an energy diagram. Based on the electronic diagrams of individual elements, you can write down their so-called electronic formulas (configurations). So, for example, an element with 15 protons and, as a result, 15 electrons, i.e. phosphorus (P) will have the following energy diagram:
When translated into an electronic formula, the phosphorus atom will take the form:
15 P = 1s 2 2s 2 2p 6 3s 2 3p 3
Normal-sized digits to the left of the sublevel symbol show the number of the energy level, and superscripts to the right of the sublevel symbol show the number of electrons in the corresponding sublevel.
Below are the electronic formulas of the first 36 elements of D.I. Mendeleev.
| period | Item No. | symbol | title | electronic formula |
| I | 1 | H | hydrogen | 1s 1 |
| 2 | He | helium | 1s2 | |
| II | 3 | Li | lithium | 1s2 2s1 |
| 4 | Be | beryllium | 1s2 2s2 | |
| 5 | B | boron | 1s 2 2s 2 2p 1 | |
| 6 | C | carbon | 1s 2 2s 2 2p 2 | |
| 7 | N | nitrogen | 1s 2 2s 2 2p 3 | |
| 8 | O | oxygen | 1s 2 2s 2 2p 4 | |
| 9 | F | fluorine | 1s 2 2s 2 2p 5 | |
| 10 | Ne | neon | 1s 2 2s 2 2p 6 | |
| III | 11 | Na | sodium | 1s 2 2s 2 2p 6 3s 1 |
| 12 | mg | magnesium | 1s 2 2s 2 2p 6 3s 2 | |
| 13 | Al | aluminum | 1s 2 2s 2 2p 6 3s 2 3p 1 | |
| 14 | Si | silicon | 1s 2 2s 2 2p 6 3s 2 3p 2 | |
| 15 | P | phosphorus | 1s 2 2s 2 2p 6 3s 2 3p 3 | |
| 16 | S | sulfur | 1s 2 2s 2 2p 6 3s 2 3p 4 | |
| 17 | Cl | chlorine | 1s 2 2s 2 2p 6 3s 2 3p 5 | |
| 18 | Ar | argon | 1s 2 2s 2 2p 6 3s 2 3p 6 | |
| IV | 19 | K | potassium | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 |
| 20 | Ca | calcium | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 | |
| 21 | sc | scandium | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 1 | |
| 22 | Ti | titanium | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 2 | |
| 23 | V | vanadium | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 3 | |
| 24 | Cr | chromium | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 5 s on the d sublevel | |
| 25 | Mn | manganese | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 5 | |
| 26 | Fe | iron | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 6 | |
| 27 | co | cobalt | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 7 | |
| 28 | Ni | nickel | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 8 | |
| 29 | Cu | copper | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 10 s on the d sublevel | |
| 30 | Zn | zinc | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 | |
| 31 | Ga | gallium | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 1 | |
| 32 | Ge | germanium | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 2 | |
| 33 | As | arsenic | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 3 | |
| 34 | Se | selenium | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 4 | |
| 35 | Br | bromine | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 5 | |
| 36 | kr | krypton | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 |
As already mentioned, in their ground state, electrons in atomic orbitals are arranged according to the principle of least energy. Nevertheless, in the presence of empty p-orbitals in the ground state of an atom, often, when excess energy is imparted to it, the atom can be transferred to the so-called excited state. So, for example, a boron atom in its ground state has an electronic configuration and an energy diagram of the following form:
5 B = 1s 2 2s 2 2p 1
And in the excited state (*), i.e. when imparting some energy to the boron atom, its electronic configuration and energy diagram will look like this:
5 B* = 1s 2 2s 1 2p 2
Depending on which sublevel in the atom is filled last, chemical elements are divided into s, p, d or f.
Finding s, p, d and f-elements in the table D.I. Mendeleev:
- s-elements have the last s-sublevel to be filled. These elements include elements of the main (on the left in the table cell) subgroups of groups I and II.
- For p-elements, the p-sublevel is filled. The p-elements include the last six elements of each period, except for the first and seventh, as well as elements of the main subgroups of III-VIII groups.
- d-elements are located between s- and p-elements in large periods.
- The f-elements are called lanthanides and actinides. They are placed at the bottom of the table by D.I. Mendeleev.
The lesson is devoted to the formation of ideas about the complex structure of the atom. The state of electrons in an atom is considered, the concepts of "atomic orbital and electron cloud", the forms of orbitals (s--, p-, d-orbitals) are introduced. Also considered are aspects such as the maximum number of electrons at energy levels and sublevels, the distribution of electrons over energy levels and sublevels in atoms of elements of the first four periods, valence electrons of s-, p- and d-elements. A graphical diagram of the structure of the electronic layers of atoms (electron-graphic formula) is given.
Topic: The structure of the atom. Periodic law D.I. Mendeleev
Lesson: The structure of the atom
Translated from Greek, the word " atom" means "indivisible". However, phenomena have been discovered that demonstrate the possibility of its division. These are the emission of x-rays, the emission of cathode rays, the phenomenon of the photoelectric effect, the phenomenon of radioactivity. Electrons, protons, and neutrons are the particles that make up an atom. They're called subatomic particles.
Tab. one
In addition to protons, the nucleus of most atoms contains neutrons that carry no charge. As can be seen from Table. 1, the mass of the neutron practically does not differ from the mass of the proton. Protons and neutrons make up the nucleus of an atom and are called nucleons (nucleus - nucleus). Their charges and masses in atomic mass units (a.m.u.) are shown in Table 1. When calculating the mass of an atom, the mass of an electron can be neglected.
Mass of an atom ( mass number) is equal to the sum of the masses of the protons and neutrons that make up its nucleus. The mass number is denoted by the letter BUT. From the name of this quantity, it can be seen that it is closely related to the atomic mass of the element rounded to an integer. A=Z+N
Here A- mass number of an atom (the sum of protons and neutrons), Z- nuclear charge (number of protons in the nucleus), N is the number of neutrons in the nucleus. According to the doctrine of isotopes, the concept of "chemical element" can be given the following definition:
chemical element A group of atoms with the same nuclear charge is called.
Some elements exist as multiple isotopes. "Isotopes" means "occupying the same place." Isotopes have the same number of protons, but differ in mass, i.e., the number of neutrons in the nucleus (number N). Because neutrons have little to no effect on the chemical properties of elements, all isotopes of the same element are chemically indistinguishable.
Isotopes are called varieties of atoms of the same chemical element with the same nuclear charge (that is, with the same number of protons), but with a different number of neutrons in the nucleus.
Isotopes differ from each other only in mass number. This is indicated either by a superscript in the right corner, or in a line: 12 C or C-12 . If an element contains several natural isotopes, then in the periodic table D.I. Mendeleev indicates its average atomic mass, taking into account the prevalence. For example, chlorine contains 2 natural isotopes 35 Cl and 37 Cl, the content of which is 75% and 25%, respectively. Thus, the atomic mass of chlorine will be equal to:
BUTr(Cl)=0,75 . 35+0,25 . 37=35,5
For artificially synthesized heavy atoms, one atomic mass value is given in square brackets. This is the atomic mass of the most stable isotope of that element.
Basic models of the structure of the atom
Historically, the Thomson model of the atom was the first in 1897.
Rice. 1. Model of the structure of the atom by J. Thomson
The English physicist J. J. Thomson suggested that atoms consist of a positively charged sphere in which electrons are interspersed (Fig. 1). This model is figuratively called "plum pudding", a bun with raisins (where "raisins" are electrons), or "watermelon" with "seeds" - electrons. However, this model was abandoned, since experimental data were obtained that contradicted it.
Rice. 2. Model of the structure of the atom by E. Rutherford
In 1910, the English physicist Ernst Rutherford, with his students Geiger and Marsden, conducted an experiment that gave amazing results that were inexplicable from the point of view of the Thomson model. Ernst Rutherford proved by experience that in the center of the atom there is a positively charged nucleus (Fig. 2), around which, like planets around the Sun, electrons revolve. The atom as a whole is electrically neutral, and the electrons are held in the atom due to the forces of electrostatic attraction (Coulomb forces). This model had many contradictions and, most importantly, did not explain why electrons do not fall on the nucleus, as well as the possibility of absorption and emission of energy by it.
The Danish physicist N. Bohr in 1913, taking Rutherford's model of the atom as a basis, proposed a model of the atom in which electron-particles revolve around the atomic nucleus in much the same way as the planets revolve around the Sun.
Rice. 3. Planetary model of N. Bohr
Bohr suggested that electrons in an atom can only exist stably in orbits at strictly defined distances from the nucleus. These orbits he called stationary. An electron cannot exist outside stationary orbits. Why this is so, Bohr could not explain at the time. But he showed that such a model (Fig. 3) makes it possible to explain many experimental facts.
Currently used to describe the structure of the atom quantum mechanics. This is a science, the main aspect of which is that the electron has the properties of a particle and a wave at the same time, i.e., wave-particle duality. According to quantum mechanics, the region of space in which the probability of finding an electron is greatest is calledorbital. The farther the electron is from the nucleus, the lower its interaction energy with the nucleus. Electrons with close energies form energy level. Number of energy levels equals period number, in which this element is located in the table D.I. Mendeleev. There are various shapes of atomic orbitals. (Fig. 4). The d-orbital and f-orbital have a more complex shape.
Rice. 4. Shapes of atomic orbitals
There are exactly as many electrons in the electron shell of any atom as there are protons in its nucleus, so the atom as a whole is electrically neutral. Electrons in an atom are arranged so that their energy is minimal. The farther the electron is from the nucleus, the more orbitals and the more complex they are in shape. Each level and sublevel can only hold a certain number of electrons. The sublevels, in turn, consist of orbitals.
At the first energy level, closest to the nucleus, there can be one spherical orbital ( 1 s). At the second energy level - a spherical orbital, large in size and three p-orbitals: 2 s2 ppp. On the third level: 3 s3 ppp3 dddd.
In addition to movement around the nucleus, electrons also have movement, which can be represented as their movement around their own axis. This rotation is called spin ( in lane from English. "spindle"). Only two electrons with opposite (antiparallel) spins can be in one orbital.
Maximum number of electrons per energy level is determined by the formula N=2 n 2.
Where n is the main quantum number (energy level number). See table. 2
Tab. 2
Depending on which orbital the last electron is in, they distinguish s-, p-, d-elements. Elements of the main subgroups belong to s-, p-elements. In the side subgroups are d-elements
Graphic diagram of the structure of the electronic layers of atoms (electronic graphic formula).
To describe the arrangement of electrons in atomic orbitals, the electronic configuration is used. To write it in a line, orbitals are written in the legend ( s--, p-, d-,f-orbitals), and in front of them are numbers indicating the number of the energy level. The larger the number, the further the electron is from the nucleus. In upper case, above the designation of the orbital, the number of electrons in this orbital is written (Fig. 5).
Rice. five
Graphically, the distribution of electrons in atomic orbitals can be represented as cells. Each cell corresponds to one orbital. There will be three such cells for the p-orbital, five for the d-orbital, and seven for the f-orbital. One cell can contain 1 or 2 electrons. According to Gund's rule, electrons are distributed in orbitals of the same energy (for example, in three p-orbitals), first one at a time, and only when there is already one electron in each such orbital, the filling of these orbitals with second electrons begins. Such electrons are called paired. This is explained by the fact that in neighboring cells, electrons repel each other less, as similarly charged particles.
See fig. 6 for atom 7 N.
Rice. 6
The electronic configuration of the scandium atom
21 sc: 1 s 2 2 s 2 2 p 6 3 s 2 3 p 6 4 s 2 3 d 1
Electrons in the outer energy level are called valence electrons. 21 sc refers to d-elements.
Summing up the lesson
At the lesson, the structure of the atom, the state of electrons in the atom were considered, the concept of "atomic orbital and electron cloud" was introduced. Students learned what the shape of orbitals is ( s-, p-, d-orbitals), what is the maximum number of electrons at energy levels and sublevels, the distribution of electrons over energy levels, what is s-, p- And d-elements. A graphical diagram of the structure of the electronic layers of atoms (electron-graphic formula) is given.
Bibliography
1. Rudzitis G.E. Chemistry. Fundamentals of General Chemistry. Grade 11: textbook for educational institutions: basic level / G.E. Rudzitis, F.G. Feldman. - 14th ed. - M.: Education, 2012.
2. Popel P.P. Chemistry: 8th grade: a textbook for general educational institutions / P.P. Popel, L.S. Krivlya. - K .: Information Center "Academy", 2008. - 240 p.: ill.
3. A.V. Manuilov, V.I. Rodionov. Fundamentals of chemistry. Internet tutorial.
Homework
1. No. 5-7 (p. 22) Rudzitis G.E. Chemistry. Fundamentals of General Chemistry. Grade 11: textbook for educational institutions: basic level / G.E. Rudzitis, F.G. Feldman. - 14th ed. - M.: Education, 2012.
2. Write electronic formulas for the following elements: 6 C, 12 Mg, 16 S, 21 Sc.
3. Elements have the following electronic formulas: a) 1s 2 2s 2 2p 4 .b) 1s 2 2s 2 2p 6 3s 2 3p 1. c) 1s 2 2s 2 2p 6 3s 2 3p 6 3d 6 4s 2 . What are these elements?
The composition of the atom.
An atom is made up of atomic nucleus And electron shell.
The nucleus of an atom is made up of protons ( p+) and neutrons ( n 0). Most hydrogen atoms have a single proton nucleus.
Number of protons N(p+) is equal to the nuclear charge ( Z) and the ordinal number of the element in the natural series of elements (and in the periodic system of elements).
N(p +) = Z
The sum of the number of neutrons N(n 0), denoted simply by the letter N, and the number of protons Z called mass number and is marked with the letter BUT.
A = Z + N
The electron shell of an atom consists of electrons moving around the nucleus ( e -).
Number of electrons N(e-) in the electron shell of a neutral atom is equal to the number of protons Z at its core.
The mass of a proton is approximately equal to the mass of a neutron and 1840 times the mass of an electron, so the mass of an atom is practically equal to the mass of the nucleus.
The shape of an atom is spherical. The radius of the nucleus is about 100,000 times smaller than the radius of the atom.
Chemical element- type of atoms (set of atoms) with the same nuclear charge (with the same number of protons in the nucleus).
Isotope- a set of atoms of one element with the same number of neutrons in the nucleus (or a type of atoms with the same number of protons and the same number of neutrons in the nucleus).
Different isotopes differ from each other in the number of neutrons in the nuclei of their atoms.
Designation of a single atom or isotope: (E - element symbol), for example: .
The structure of the electron shell of the atom
atomic orbital is the state of an electron in an atom. Orbital symbol - . Each orbital corresponds to an electron cloud.
The orbitals of real atoms in the ground (unexcited) state are of four types: s, p, d And f.
electronic cloud- the part of space in which an electron can be found with a probability of 90 (or more) percent.
Note: sometimes the concepts of "atomic orbital" and "electron cloud" are not distinguished, calling both of them "atomic orbital".
The electron shell of an atom is layered. Electronic layer formed by electron clouds of the same size. Orbitals of one layer form electronic ("energy") level, their energies are the same for the hydrogen atom, but different for other atoms.
Orbitals of the same level are grouped into electronic (energy) sublevels:
s- sublevel (consists of one s-orbitals), symbol - .
p sublevel (consists of three p
d sublevel (consists of five d-orbitals), symbol - .
f sublevel (consists of seven f-orbitals), symbol - .
The energies of the orbitals of the same sublevel are the same.
When designating sublevels, the number of the layer (electronic level) is added to the sublevel symbol, for example: 2 s, 3p, 5d means s- sublevel of the second level, p- sublevel of the third level, d- sublevel of the fifth level.
The total number of sublevels in one level is equal to the level number n. The total number of orbitals in one level is n 2. Accordingly, the total number of clouds in one layer is also n 2 .
Designations: - free orbital (without electrons), - orbital with an unpaired electron, - orbital with an electron pair (with two electrons).
The order in which electrons fill the orbitals of an atom is determined by three laws of nature (formulations are given in a simplified way):
1. The principle of least energy - electrons fill the orbitals in order of increasing energy of the orbitals.
2. Pauli's principle - there cannot be more than two electrons in one orbital.
3. Hund's rule - within the sublevel, electrons first fill free orbitals (one at a time), and only after that they form electron pairs.
The total number of electrons in the electronic level (or in the electronic layer) is 2 n 2 .
The distribution of sublevels by energy is expressed next (in order of increasing energy):
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p ...
Visually, this sequence is expressed by the energy diagram:
The distribution of electrons of an atom by levels, sublevels and orbitals (the electronic configuration of an atom) can be depicted as an electronic formula, an energy diagram, or, more simply, as a diagram of electronic layers ("electronic diagram").
Examples of the electronic structure of atoms:
Valence electrons- electrons of an atom that can take part in the formation of chemical bonds. For any atom, these are all the outer electrons plus those pre-outer electrons whose energy is greater than that of the outer ones. For example: Ca atom has 4 outer electrons s 2, they are also valence; the Fe atom has external electrons - 4 s 2 but he has 3 d 6, hence the iron atom has 8 valence electrons. The valence electronic formula of the calcium atom is 4 s 2, and iron atoms - 4 s 2 3d 6 .
Periodic system of chemical elements of D. I. Mendeleev
(natural system of chemical elements)
Periodic law of chemical elements(modern formulation): the properties of chemical elements, as well as simple and complex substances formed by them, are in a periodic dependence on the value of the charge from atomic nuclei.
Periodic system- graphical expression of the periodic law.
Natural range of chemical elements- a number of chemical elements, arranged according to the increase in the number of protons in the nuclei of their atoms, or, what is the same, according to the increase in the charges of the nuclei of these atoms. The serial number of an element in this series is equal to the number of protons in the nucleus of any atom of this element.
The table of chemical elements is constructed by "cutting" the natural series of chemical elements into periods(horizontal rows of the table) and groupings (vertical columns of the table) of elements with a similar electronic structure of atoms.
Depending on how elements are combined into groups, a table can be long period(elements with the same number and type of valence electrons are collected in groups) and short-term(elements with the same number of valence electrons are collected in groups).
The groups of the short period table are divided into subgroups ( main And side effects), coinciding with the groups of the long-period table.
All atoms of elements of the same period have the same number of electron layers, equal to the number of the period.
The number of elements in the periods: 2, 8, 8, 18, 18, 32, 32. Most of the elements of the eighth period were obtained artificially, the last elements of this period have not yet been synthesized. All periods except the first start with an alkali metal forming element (Li, Na, K, etc.) and end with a noble gas forming element (He, Ne, Ar, Kr, etc.).
In the short period table - eight groups, each of which is divided into two subgroups (main and secondary), in the long period table - sixteen groups, which are numbered in Roman numerals with the letters A or B, for example: IA, IIIB, VIA, VIIB. Group IA of the long period table corresponds to the main subgroup of the first group of the short period table; group VIIB - secondary subgroup of the seventh group: the rest - similarly.
The characteristics of chemical elements naturally change in groups and periods.
In periods (with increasing serial number)
- the nuclear charge increases
- the number of outer electrons increases,
- the radius of the atoms decreases,
- the bond strength of electrons with the nucleus increases (ionization energy),
- electronegativity increases.
- the oxidizing properties of simple substances are enhanced ("non-metallicity"),
- the reducing properties of simple substances ("metallicity") weaken,
- weakens the basic character of hydroxides and the corresponding oxides,
- the acidic character of hydroxides and corresponding oxides increases.
In groups (with increasing serial number)
- the nuclear charge increases
- the radius of atoms increases (only in A-groups),
- the strength of the bond between electrons and the nucleus decreases (ionization energy; only in A-groups),
- electronegativity decreases (only in A-groups),
- weaken the oxidizing properties of simple substances ("non-metallicity"; only in A-groups),
- the reducing properties of simple substances are enhanced ("metallicity"; only in A-groups),
- the basic character of hydroxides and the corresponding oxides increases (only in A-groups),
- the acidic nature of hydroxides and the corresponding oxides weakens (only in A-groups),
- the stability of hydrogen compounds decreases (their reducing activity increases; only in A-groups).
Tasks and tests on the topic "Topic 9. "The structure of the atom. Periodic law and periodic system of chemical elements of D. I. Mendeleev (PSCE)"."
- Periodic Law - Periodic law and structure of atoms Grade 8–9
You should know: the laws of filling orbitals with electrons (principle of least energy, Pauli's principle, Hund's rule), the structure of the periodic system of elements.You should be able to: determine the composition of an atom by the position of an element in the periodic system, and, conversely, find an element in the periodic system, knowing its composition; depict the structure diagram, the electronic configuration of an atom, ion, and, conversely, determine the position of a chemical element in the PSCE from the diagram and electronic configuration; characterize the element and the substances it forms according to its position in the PSCE; determine changes in the radius of atoms, the properties of chemical elements and the substances they form within one period and one main subgroup of the periodic system.
Example 1 Determine the number of orbitals in the third electronic level. What are these orbitals?
To determine the number of orbitals, we use the formula N orbitals = n 2 , where n- level number. N orbitals = 3 2 = 9. One 3 s-, three 3 p- and five 3 d-orbitals.Example 2 Determine the atom of which element has the electronic formula 1 s 2 2s 2 2p 6 3s 2 3p 1 .
In order to determine which element it is, you need to find out its serial number, which is equal to the total number of electrons in the atom. In this case: 2 + 2 + 6 + 2 + 1 = 13. This is aluminum.After making sure that everything you need is learned, proceed to the tasks. We wish you success.
Recommended literature:- O. S. Gabrielyan and others. Chemistry, 11th grade. M., Bustard, 2002;
- G. E. Rudzitis, F. G. Feldman. Chemistry 11 cells. M., Education, 2001.
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