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.
2. The structure of nuclei and electron shells of atoms
2.7. Distribution of electrons in an atom
The state of the electrons in an atom is indicated using a specific notation. For example, for a helium atom we have:
The distribution of electrons in an atom is indicated by:
a) electronic circuits, in which only the number of electrons in each layer is noted. For example: Mg 2e , 8e , 2e ; Cl2e, 8e, 7e.
Often used graphics electronic circuits, for example, for the chlorine atom:
b) electronic configurations; in this case, the number of the layer (level), the nature of the sublevels, and the number of electrons on them are shown. For example:
Mg 1s 2 2s 2 2p 6 3s 2 ;
in) electronic graphic schemes, on which orbitals are depicted, for example, in the form of a cell, and electrons are depicted by arrows (Fig. 2.6).
Rice. 2.6. Electronic graphic scheme for the magnesium atom
In addition to the full formulas of electronic configurations, abbreviated ones are widely used. In this case, the noble gas portion of the electron configuration is indicated by the noble gas symbol in square brackets. For example: 12 Mg3s 2 , 19 K4s 1 .
There are certain principles and rules for filling energy levels and sublevels with electrons:
1. The principle of minimum total energy of an atom, according to which the population of AO with electrons occurs in such a way that the total energy of the atom is minimal. The following sequence of AO filling was experimentally established:
1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p ... .
2. One AO can contain no more than two electrons, and in this case their spins must be antiparallel.
3. Within a given energy sublevel, electrons fill the AO gradually, first one at a time (first all vacant ones, and only then two), and the orientation of all unpaired electrons should be the same, i.e. such
but not like that
In almost any atom, only s - and p -AO are external (Fig. 2.7), therefore The outer electron layer of any atom cannot have more than eight electrons.. An outer electron layer containing eight electrons (two in the case of helium) is called complete.
Rice. 2.7. Electronic graphic schemes for atoms K (a) and S (b)
Electronic configurations of atoms of elements of the 4th period of the periodic systemThe energy values of different energy sublevels for different atoms are not constant, but depend on the charge of the nucleus Z of an element atom: for atoms of elements with Z = 1–20 Е 3 d > E 4 s and Е 3 d > E 4 p ; for atoms of elements with Z ≥ 21 vice versa: E 3 d< E 4 s и Е 3 d < E 4 p (рис. 2.8). Кроме того, чем больше Z , тем меньше различаются подуровни по энергии, а кривые, выражающие зависимость энергии подуровней от Z , пересекаются.
Rice. 2.8. Diagram of energy sublevels of atoms of elements with Z = 1–20 (a), Z ≥ 21 (b)
The electronic configurations of atoms (ground state) K and Ca are as follows (see Fig. 2.8):
19 K: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 ,
20 Ca: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 .
Starting from scandium (Z = 21), the 3d sublevel is filled, and 4s electrons remain in the outer layer. The general electronic formula of atoms of elements from Sc to Zn is 3d 1−10 4s 1−2. For example:
21 Sc: 3d 1 4s 2 ,
25 Mn: 3d 5 4s 2 ,
28 Ni: 3d 8 4s 2 .
30 Zn: 3d 10 4s 2 .
For chromium and copper, a slip (failure) of the 4s electron to the 3d sublevel is observed: Cr - 3d 5 4s 1, Cu - 3d 10 4s 1. Such a jump from the ns - to the (n - 1)d -sublevel is also observed in atoms of other elements (Mo, Ag, Au, Pt) and is explained by the closeness of the energies of the ns - and (n - 1)d -sublevels, as well as the stability of half and completely filled d-sublevels.
Further in the 4th period after 10 d-elements follow from Ga ( 3d 10 4s 2 4p 1) to Kr ( 3d 10 4s 2 4p 6) p-elements.
The formation of cations of d-elements is associated with the loss of first external ns-, then (n - 1)d-electrons, for example:
Ti: 3d 2 4s 2 → − 2 e − Ti 2+ : 3d 2 → − 1 e − Ti 3+ : 3d 1
Mn: 3d 5 4s 2 → − 2 e − Mn 2+ : 3d 5 → − 2 e − Mn 4+ : 3d 3
Note that in the formulas of electronic configurations, it is customary to first write down all the electrons with a lower value of n, and then proceed to indicate the electrons with a higher value of the principal quantum number. Therefore, the order of filling and the order of recording energy sublevels for 3d elements do not match. For example, in the electronic formula of the scandium atom, the 3d orbital is indicated before the 4s orbital, although the 4s orbital is filled earlier.
A natural question arises: why is the 4s sublevel filled earlier in the atoms of 3d elements, although its energy is greater than the energy of the 3d sublevel? Why, for example, does the Sc atom not have the electronic configuration 1s 2 2s 2 2p 6 3s 2 3p 6 3d 3 in the ground state?
This is because the ratio of the energies of various electronic states of an atom does not always coincide with the ratio of the energies of individual energy sublevels. The energy of the 4s sublevel for 3d elements is greater than the energy of the 3d sublevel, but the energy of the state
3d 1 4s 2 is less than the energy of the state 3d 3 .
This is explained by the fact that the interelectronic repulsion, and, accordingly, the energy of the entire state for the configuration ... 3d 3 (with three electrons on the same energy sublevel) is greater than for the configuration ... 3d 1 4s 2 (with three electrons, at different energy levels).
It is written in the form of so-called electronic formulas. In electronic formulas, the letters s, p, d, f denote the energy sublevels of electrons; the numbers in front of the letters indicate the energy level in which the given electron is located, and the index at the top right is the number of electrons in this sublevel. To compose the electronic formula of an atom of any element, it is enough to know the number of this element in the periodic system and fulfill the basic provisions that govern the distribution of electrons in an atom.
The structure of the electron shell of an atom can also be depicted in the form of an arrangement of electrons in energy cells.
For iron atoms, such a scheme has the following form:
This diagram clearly shows the implementation of Hund's rule. On the 3d sublevel maximum amount, cells (four) are filled with unpaired electrons. The image of the structure of the electron shell in the atom in the form of electronic formulas and in the form of diagrams does not clearly reflect the wave properties of the electron.
The wording of the periodic law as amended YES. Mendeleev : the properties of simple bodies, as well as the forms and properties of the compounds of elements, are in a periodic dependence on the magnitude of the atomic weights of the elements.
Modern formulation of the Periodic Law: the properties of the elements, as well as the forms and properties of their compounds, are in a periodic dependence on the magnitude of the charge of the nucleus of their atoms.
Thus, the positive charge of the nucleus (rather than atomic mass) turned out to be a more accurate argument on which the properties of elements and their compounds depend.
Valence- is the number of chemical bonds that one atom is bonded to another.
The valence possibilities of an atom are determined by the number of unpaired electrons and the presence of free atomic orbitals at the outer level. The structure of the outer energy levels of atoms of chemical elements determines mainly the properties of their atoms. Therefore, these levels are called valence levels. The electrons of these levels, and sometimes of the pre-external levels, can take part in the formation of chemical bonds. Such electrons are also called valence electrons.
Stoichiometric valency chemical element - is the number of equivalents that a given atom can attach to itself, or is the number of equivalents in the atom.
Equivalents are determined by the number of attached or substituted hydrogen atoms, therefore, the stoichiometric valence is equal to the number of hydrogen atoms with which this atom interacts. But not all elements interact freely, but almost everything interacts with oxygen, so the stoichiometric valency can be defined as twice the number of attached oxygen atoms.
For example, the stoichiometric valency of sulfur in hydrogen sulfide H 2 S is 2, in oxide SO 2 - 4, in oxide SO 3 -6.
When determining the stoichiometric valence of an element according to the formula of a binary compound, one should be guided by the rule: the total valency of all atoms of one element must be equal to the total valence of all atoms of another element.
Oxidation state also characterizes the composition of the substance and is equal to the stoichiometric valence with a plus sign (for a metal or a more electropositive element in a molecule) or minus.
1. In simple substances, the oxidation state of elements is zero.
2. The oxidation state of fluorine in all compounds is -1. The remaining halogens (chlorine, bromine, iodine) with metals, hydrogen and other more electropositive elements also have an oxidation state of -1, but in compounds with more electronegative elements they have positive oxidation states.
3. Oxygen in compounds has an oxidation state of -2; the exceptions are hydrogen peroxide H 2 O 2 and its derivatives (Na 2 O 2, BaO 2, etc., in which oxygen has an oxidation state of -1, as well as oxygen fluoride OF 2, in which the oxidation state of oxygen is +2.
4. Alkaline elements (Li, Na, K, etc.) and elements of the main subgroup of the second group of the Periodic system (Be, Mg, Ca, etc.) always have an oxidation state equal to the group number, that is, +1 and +2, respectively .
5. All elements of the third group, except for thallium, have a constant oxidation state equal to the group number, i.e. +3.
6. The highest oxidation state of an element is equal to the group number of the Periodic system, and the lowest is the difference: group number - 8. For example, highest degree nitrogen oxidation (it is located in the fifth group) is +5 (in nitric acid and its salts), and the lowest is -3 (in ammonia and ammonium salts).
7. The oxidation states of the elements in the compound compensate each other so that their sum for all atoms in a molecule or a neutral formula unit is zero, and for an ion - its charge.
These rules can be used to determine the unknown oxidation state of an element in a compound, if the oxidation states of the others are known, and to formulate multi-element compounds.
Degree of oxidation (oxidation number,) — auxiliary conditional value for recording the processes of oxidation, reduction and redox reactions.
concept oxidation state often used in inorganic chemistry instead of the concept valence. The oxidation state of an atom is equal to the numerical value of the electric charge attributed to the atom, assuming that the electron pairs that carry out the bond are completely biased towards more electronegative atoms (that is, based on the assumption that the compound consists only of ions).
The oxidation state corresponds to the number of electrons that must be added to a positive ion to reduce it to a neutral atom, or taken from a negative ion to oxidize it to a neutral atom:
Al 3+ + 3e − → Al
S 2− → S + 2e − (S 2− − 2e − → S)
The properties of the elements, depending on the structure of the electron shell of the atom, change according to the periods and groups of the periodic system. Since in a number of analogous elements the electronic structures are only similar, but not identical, when moving from one element in a group to another, not a simple repetition of properties is observed for them, but their more or less clearly expressed regular change.
The chemical nature of an element is determined by the ability of its atom to lose or gain electrons. This ability is quantified by the values of ionization energies and electron affinity.
Ionization energy (Ei) is the minimum amount of energy required for the detachment and complete removal of an electron from an atom in the gas phase at T = 0
K without transferring kinetic energy to the released electron with the transformation of the atom into a positively charged ion: E + Ei = E + + e-. The ionization energy is a positive value and has the lowest values for alkali metal atoms and the highest for noble (inert) gas atoms.
Electron affinity (Ee) is the energy released or absorbed when an electron is attached to an atom in the gas phase at T = 0
K with the transformation of the atom into a negatively charged ion without transferring kinetic energy to the particle:
E + e- = E- + Ee.
Halogens, especially fluorine, have the maximum electron affinity (Ee = -328 kJ/mol).
The values of Ei and Ee are expressed in kilojoules per mol (kJ/mol) or in electron volts per atom (eV).
The ability of a bound atom to displace the electrons of chemical bonds towards itself, increasing the electron density around itself is called electronegativity.
This concept was introduced into science by L. Pauling. Electronegativitydenoted by the symbol ÷ and characterizes the tendency of a given atom to attach electrons when it forms a chemical bond.
According to R. Maliken, the electronegativity of an atom is estimated by half the sum of the ionization energies and the electron affinity of free atoms h = (Ee + Ei)/2
In periods, there is a general tendency for an increase in the ionization energy and electronegativity with an increase in the charge of the atomic nucleus; in groups, these values decrease with an increase in the ordinal number of the element.
It should be emphasized that an element cannot be assigned a constant value of electronegativity, since it depends on many factors, in particular, on the valence state of the element, the type of compound in which it is included, the number and type of neighboring atoms.
Atomic and ionic radii. The dimensions of atoms and ions are determined by the dimensions of the electron shell. According to quantum mechanical concepts, the electron shell does not have strictly defined boundaries. Therefore, for the radius of a free atom or ion, we can take theoretically calculated distance from the core to the position of the main maximum density of the outer electron clouds. This distance is called the orbital radius. In practice, the values of the radii of atoms and ions in compounds, calculated from experimental data, are usually used. In this case, covalent and metallic radii of atoms are distinguished.
The dependence of atomic and ionic radii on the charge of the nucleus of an atom of an element and is periodic. In periods, as the atomic number increases, the radii tend to decrease. The greatest decrease is typical for elements of small periods, since the outer electronic level is filled in them. In large periods in the families of d- and f-elements, this change is less sharp, since the filling of electrons in them occurs in the preexternal layer. In subgroups, the radii of atoms and ions of the same type generally increase.
The periodic system of elements is a clear example of the manifestation of various kinds of periodicity in the properties of elements, which is observed horizontally (in a period from left to right), vertically (in a group, for example, from top to bottom), diagonally, i.e. some property of the atom increases or decreases, but the periodicity is preserved.
In the period from left to right (→), the oxidizing and non-metallic properties of the elements increase, while the reducing and metallic properties decrease. So, of all the elements of period 3, sodium will be the most active metal and the strongest reducing agent, and chlorine will be the strongest oxidizing agent.
chemical bond- this is the interconnection of atoms in a molecule, or crystal lattice, as a result of the action of electric forces of attraction between atoms.
This is the interaction of all electrons and all nuclei, leading to the formation of a stable, polyatomic system (radical, molecular ion, molecule, crystal).
Chemical bonding is carried out by valence electrons. According to modern concepts, the chemical bond has an electronic nature, but it is carried out in different ways. Therefore, there are three main types of chemical bonds: covalent, ionic, metallic. Between molecules arises hydrogen bond, and happen van der Waals interactions.
The main characteristics of a chemical bond are:
- bond length - is the internuclear distance between chemically bonded atoms.
It depends on the nature of the interacting atoms and on the multiplicity of the bond. With an increase in the multiplicity, the bond length decreases, and, consequently, its strength increases;
- bond multiplicity - is determined by the number of electron pairs linking two atoms. As the multiplicity increases, the binding energy increases;
- connection angle- the angle between imaginary straight lines passing through the nuclei of two chemically interconnected neighboring atoms;
Binding energy E CB - this is the energy that is released during the formation of this bond and is spent on breaking it, kJ / mol.
covalent bond - A chemical bond formed by sharing a pair of electrons with two atoms.
The explanation of the chemical bond by the appearance of common electron pairs between atoms formed the basis of the spin theory of valence, the tool of which is valence bond method (MVS) , discovered by Lewis in 1916. For the quantum mechanical description of the chemical bond and the structure of molecules, another method is used - molecular orbital method (MMO) .
Valence bond method
The basic principles of the formation of a chemical bond according to MVS:
1. A chemical bond is formed due to valence (unpaired) electrons.
2. Electrons with antiparallel spins belonging to two different atoms become common.
3. A chemical bond is formed only if, when two or more atoms approach each other, the total energy of the system decreases.
4. The main forces acting in the molecule are of electrical, Coulomb origin.
5. The stronger the connection, the more the interacting electron clouds overlap.
There are two mechanisms for the formation of a covalent bond:
exchange mechanism. The bond is formed by sharing the valence electrons of two neutral atoms. Each atom gives one unpaired electron to a common electron pair:
Rice. 7. Exchange mechanism for the formation of a covalent bond: a- non-polar; b- polar
Donor-acceptor mechanism. One atom (donor) provides an electron pair, and another atom (acceptor) provides an empty orbital for this pair.
connections, educated according to the donor-acceptor mechanism, belong to complex compounds
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Rice. 8. Donor-acceptor mechanism of covalent bond formation
A covalent bond has certain characteristics.
Saturability - the property of atoms to form a strictly defined number of covalent bonds. Due to the saturation of the bonds, the molecules have a certain composition.
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Orientation - t . e. the connection is formed in the direction of maximum overlap of electron clouds . With respect to the line connecting the centers of atoms forming a bond, there are: σ and π (Fig. 9): σ-bond - formed by overlapping AO along the line connecting the centers of interacting atoms; A π-bond is a bond that occurs in the direction of an axis perpendicular to the straight line connecting the nuclei of an atom. The orientation of the bond determines the spatial structure of the molecules, i.e., their geometric shape. hybridization - it is a change in the shape of some orbitals in the formation of a covalent bond in order to achieve a more efficient overlap of orbitals. The chemical bond formed with the participation of electrons of hybrid orbitals is stronger than the bond with the participation of electrons of non-hybrid s- and p-orbitals, since there is more overlap. There are the following types of hybridization (Fig. 10, Table 31): |
sp hybridization - one s-orbital and one p-orbital turn into two identical "hybrid" orbitals, the angle between the axes of which is 180°. Molecules in which sp hybridization occurs have a linear geometry (BeCl 2).sp 2 hybridization- one s-orbital and two p-orbitals turn into three identical "hybrid" orbitals, the angle between the axes of which is 120°. Molecules in which sp 2 hybridization is carried out have a flat geometry (BF 3 , AlCl 3).
sp 3-hybridization- one s-orbital and three p-orbitals turn into four identical "hybrid" orbitals, the angle between the axes of which is 109 ° 28 ". Molecules in which sp 3 hybridization occurs have a tetrahedral geometry (CH 4 , NH3).
Rice. 10. Types of hybridizations of valence orbitals: a - sp-hybridization of valence orbitals; b - sp2- hybridization of valence orbitals; in - sp 3 - hybridization of valence orbitals
The structure of the electron shells of atoms play an important role in chemistry, determine the chemical properties of substances. The most important characteristic of the motion of an electron in a certain orbit is the energy of its connection with the nucleus. Electrons in an atom differ in a certain energy, and, as experiments show, some are attracted to the nucleus more strongly, others weaker. This is explained by the remoteness of electrons from the nucleus. The closer the electrons to the nucleus, the greater their bond with the nucleus, but the less energy. As the distance from the nucleus of the atom, the force of attraction of the electron to the nucleus decreases, and the energy supply increases. This is how electron layers are formed in the electron shell of an atom. Electrons with similar energy values form a single electron layer, or energy level. The energy of electrons in an atom and the energy level is determined by the main quantum number n and takes integer values 1, 2, 3, 4, 5, 6 and 7. The larger the value of n, the greater the energy of the electron in the atom. The maximum number of electrons that can be in a particular energy level is determined by the formula:
Where N is the maximum number of electrons per level;
n is the number of the energy level.
It has been established that no more than two electrons are located on the first shell, no more than eight on the second, no more than 18 on the third, and no more than 32 on the fourth. We will not consider the filling of more distant shells. It is known that the external energy level can contain no more than eight electrons, it is called complete. Electronic layers that do not contain the maximum number of electrons are called incomplete.
The number of electrons in the outer energy level of the electron shell of an atom is equal to the group number for the chemical elements of the main subgroups.
As previously mentioned, the electron does not move in an orbit, but in an orbit and has no trajectory.
The space around the nucleus where a given electron is most likely to be is called that electron's orbital, or electron cloud.
Orbitals, or sublevels, as they are also called, can have different shapes, and their number corresponds to the level number, but does not exceed four. The first energy level has one sublevel (s), the second has two (s,p), the third has three (s,p,d), and so on. Electrons of different sublevels of the same level have different shapes of the electron cloud: spherical (s), dumbbell-shaped (p), and more complex configurations (d) and (f). Scientists agreed to call the spherical atomic orbital s-orbital. It is the most stable and is located quite close to the core.
The greater the energy of an electron in an atom, the faster it rotates, the more the region of its stay is extended, and, finally, it turns into a dumbbell-shaped p-orbital:
An electron cloud of this form can occupy three positions in an atom along the coordinate axes of space x, y and z. This is easily explained: after all, all electrons are negatively charged, so the electron clouds repel each other and tend to be located as far as possible from each other.
So, p There can be three orbitals. Their energy, of course, is the same, but their location in space is different.
Draw a diagram of the sequential filling of energy levels with electrons
Now we can draw up a diagram of the structure of the electron shells of atoms:
1. Determine the total number of electrons on the shell by the element's serial number.
2. Determine the number of energy levels in the electron shell. Their number is equal to the number of the period in the table of D. I. Mendeleev, in which the element is located.
3. Determine the number of electrons at each energy level.
4. Using Arabic numerals to designate the level and designating the orbitals with the letters s and p, and the number of electrons in a given orbital Arabic numeral at the top right above the letter, we depict the structure of atoms with more complete electronic formulas. Scientists agreed to designate each atomic orbital as a quantum cell - a square on the energy diagram:
On the s A sublevel can contain one atomic orbital
and on p-there may already be three sublevels -
(according to the three coordinate axes):
Orbitals d- and f- sublevels in an atom can already be five and seven, respectively:
The nucleus of a hydrogen atom has a charge of +1, so only one electron moves around its nucleus at a single energy level. Let's write down the electronic configuration of the hydrogen atom
To establish a connection between the structure of the atom of a chemical element and its properties, we will consider a few more chemical elements.
The next element after hydrogen is helium. The nucleus of a helium atom has a charge of +2, so a helium atom contains two electrons in the first energy level:
Since the first energy level can contain no more than two electrons, it is considered complete.
Element number 3 - lithium. The lithium nucleus has a charge of +3, therefore, there are three electrons in the lithium atom. Two of them are at the first energy level, and the third electron begins to fill the second energy level. First, the s-orbital of the first level is filled, then the s-orbital of the second level. The electron in the second level is weaker bound to the nucleus than the other two.
For the carbon atom, we can already assume three possible schemes filling of electron shells in accordance with electron-graphic formulas:
An analysis of the atomic spectrum shows that the latter scheme is correct. Using this rule, it is not difficult to draw up a diagram of the electronic structure for the nitrogen atom:
This scheme corresponds to the formula 1s22s22p3. Then the pairwise placement of electrons into 2p orbitals begins. Electronic formulas of the remaining atoms of the second period:
The filling of the second energy level of the neon atom ends, and the construction of the second period of the system of elements is completed.
Find the chemical sign of lithium in the periodic system, from lithium to neon Ne, the charge of atomic nuclei naturally increases. The second layer is gradually filled with electrons. With an increase in the number of electrons in the second layer, the metallic properties of the elements gradually weaken and are replaced by non-metallic ones.
The third period, like the second, begins with two elements (Na, Mg), in which the electrons are located on the s-sublevel of the outer electron layer. This is followed by six elements (from Al to Ar), in which the p-sublevel of the outer electron layer is formed. The structure of the outer electronic layer of the corresponding elements of the second and third periods is similar. In other words, with an increase in the charge of the nucleus, the electronic structure of the outer layers of atoms is periodically repeated. If the elements have the same external energy levels, then the properties of these elements are similar. For example, argon and neon contain eight electrons at the outer level, and therefore they are inert, that is, they almost do not enter into chemical reactions. In the free form, argon and neon are gases that have monatomic molecules.
The atoms of lithium, sodium and potassium contain one electron at the outer level and have similar properties, so they are placed in the same group of the periodic system.
III. Conclusions.
1. The properties of chemical elements, arranged in ascending order of the nuclear charge, are periodically repeated, since the structure of the external energy levels of the elements' atoms is periodically repeated.
2. A smooth change in the properties of chemical elements within one period can be explained by a gradual increase in the number of electrons at the external energy level.
3. The reason for the similarity of the properties of chemical elements belonging to the same family lies in the same structure of the external energy levels of their atoms.
The electronic structure of an atom can be shown by an electronic formula and an electronic graphic diagram. In electronic formulas, the energy levels and sublevels are sequentially written in the order of their filling and the total number of electrons in the sublevel. In this case, the state of an individual electron, in particular its magnetic and spin quantum numbers, is not reflected in the electronic formula. In electronic graphic schemes, each electron is “visible” completely, i.e. it can be characterized by all four quantum numbers. Electronic graphic diagrams are usually given for external electrons.
Example 1 Write the electronic formula of fluorine, express the state of external electrons with an electronic graphic diagram. How many unpaired electrons are in an atom of this element?
Solution. The atomic number of fluorine is nine, so there are nine electrons in its atom. In accordance with the principle of least energy, using Fig. 7 and taking into account the consequences of the Pauli principle, we write down the electronic formula of fluorine: 1s 2 2s 2 2p 5 . For external electrons (the second energy level), we draw up an electronic graphic diagram (Fig. 8), from which it follows that there is one unpaired electron in the fluorine atom.
Rice. 8. Electron-graphic scheme of valence electrons of a fluorine atom
Example 2 Make electron-graphic diagrams of possible states of the nitrogen atom. Which of them reflect the normal state, and which - excited?
Solution. The electronic formula of nitrogen is 1s 2 s 2 2p 3 , the formula of external electrons is 2s 2 2p 3 . Sublevel 2p is incomplete, because the number of electrons on it is less than six. Possible options the distributions of three electrons on the 2p sublevel are shown in Figs. 9.
Rice. 9. Electron-graphic diagrams of possible states of the 2p sublevel in the nitrogen atom.
The maximum (in absolute value) value of the spin (3 / 2) corresponds to states 1 and 2, therefore, they are ground, and the rest are excited.
Example 3 Determine the quantum numbers that determine the state of the last electron in a vanadium atom?
Solution. The atomic number of vanadium is Z = 23, therefore, the full electronic formula of the element is: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 3. The electronic graphic scheme of external electrons (4s 2 3d 3) is as follows (Fig. 10):
Rice. 10. Electron-graphic scheme of valence electrons of the vanadium atom
Principal quantum number of the last electron n = 3 (third energy level), orbital l= 2 (sublevel d). The magnetic quantum number for each of the three d-electrons is different: for the first it is -2, for the second -1, for the third - 0. The spin quantum number for all three electrons is the same: m s \u003d + 1 / 2. Thus, the state of the last electron in the vanadium atom is characterized by quantum numbers: n = 3; l= 2; m = 0; m s = + 1 / 2 .
7. Paired and unpaired electrons
Electrons that fill orbitals in pairs are called paired, and single electrons are called unpaired. Unpaired electrons provide the chemical bond of an atom with other atoms. The presence of unpaired electrons is established experimentally by studying the magnetic properties. Substances with unpaired electrons paramagnetic(they are drawn into a magnetic field due to the interaction of electron spins, like elementary magnets, with an external magnetic field). Substances that have only paired electrons diamagnetic(external magnetic field does not act on them). Unpaired electrons are located only on the outer energy level of an atom and their number can be determined from its electronic graphic scheme.
Example 4 Determine the number of unpaired electrons in a sulfur atom.
Solution. The atomic number of sulfur is Z = 16, therefore, the full electronic formula of the element is: 1s 2 2s 2 2p 6 3s 2 3p 4. The electronic graphic scheme of external electrons is as follows (Fig. 11).
Rice. 11. Electron-graphic scheme of valence electrons of a sulfur atom
It follows from the electron-graphic scheme that there are two unpaired electrons in the sulfur atom.
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