ENERGY LEVELS
| Parameter name | Meaning |
| Article subject: | ENERGY LEVELS |
| Rubric (thematic category) | Education |
STRUCTURE OF THE ATOM
1. Development of the theory of the structure of the atom. With
2. The nucleus and electron shell of the atom. With
3. The structure of the nucleus of an atom. With
4. Nuclides, isotopes, mass number. With
5. Energy levels.
6. Quantum-mechanical explanation of the structure.
6.1. Orbital model of the atom.
6.2. Rules for filling orbitals.
6.3. Orbitals with s-electrons (atomic s-orbitals).
6.4. Orbitals with p-electrons (atomic p-orbitals).
6.5. Orbitals with d-f electrons
7. Energy sublevels of a multielectron atom. quantum numbers.
ENERGY LEVELS
Structure electron shell atom is determined by the different energy reserves of individual electrons in the atom. In accordance with the Bohr model of the atom, electrons can occupy positions in the atom that correspond to precisely defined (quantized) energy states. These states are called energy levels.
The number of electrons that can be on a separate energy level is determined by the formula 2n 2, where n is the number of the level, which is denoted by Arabic numerals 1 - 7. The maximum filling of the first four energy levels in. in accordance with the formula 2n 2 is: for the first level - 2 electrons, for the second - 8, for the third -18 and for the fourth level - 32 electrons. The maximum filling of higher energy levels in atoms of known elements with electrons has not been achieved.
Rice. 1 shows the filling of the energy levels of the first twenty elements with electrons (from hydrogen H to calcium Ca, black circles). By filling in the energy levels in the indicated order, the simplest models of the atoms of the elements are obtained, while observing the order of filling (from bottom to top and from left to right in the figure) in such a way that the last electron points to the symbol of the corresponding element At the third energy level M(maximum capacity is 18 e -) for elements Na - Ar contains only 8 electrons, then the fourth energy level begins to build up N- two electrons appear on it for the elements K and Ca. The next 10 electrons again occupy the level M(elements Sc – Zn (not shown), and then the filling of the N level with six more electrons continues (elements Ca-Kr, white circles).
| Rice. one | Rice. 2 |
If the atom is in the ground state, then its electrons occupy levels with a minimum energy, i.e., each subsequent electron occupies the energetically most favorable position, such as in Fig. 1. With an external impact on an atom associated with the transfer of energy to it, for example, by heating, electrons are transferred to higher energy levels (Fig. 2). This state of the atom is called excited. The place vacated at the lower energy level is filled (as an advantageous position) by an electron from a higher energy level. During the transition, the electron gives off a certain amount of energy, ĸᴏᴛᴏᴩᴏᴇ corresponds to the energy difference between the levels. As a result of electronic transitions, characteristic radiation arises. From the spectral lines of the absorbed (emitted) light, one can make a quantitative conclusion about the energy levels of the atom.
In accordance with the Bohr quantum model of the atom, an electron having a certain energy state moves in a circular orbit in the atom. Electrons with the same energy reserve are located at equal distances from the nucleus, each energy level corresponds to its own set of electrons, called the electron layer by Bohr. Τᴀᴋᴎᴍ ᴏϬᴩᴀᴈᴏᴍ, according to Bohr, the electrons of one layer move along a spherical surface, the electrons of the next layer along another spherical surface. all spheres are inscribed one into another with the center corresponding to the atomic nucleus.
ENERGY LEVELS - concept and types. Classification and features of the category "ENERGY LEVELS" 2017, 2018.
E.N.FRENKEL
Chemistry tutorial
A guide for those who do not know, but want to learn and understand chemistry
Part I. Elements general chemistry
(first level of difficulty)
Continuation. See the beginning in No. 13, 18, 23/2007
Chapter 3. Elementary information about the structure of the atom.
Periodic law of D.I. Mendeleev
Remember what an atom is, what an atom consists of, whether an atom changes in chemical reactions.
An atom is an electrically neutral particle consisting of a positively charged nucleus and negatively charged electrons.
The number of electrons during chemical processes can change, but nuclear charge always stays the same. Knowing the distribution of electrons in an atom (the structure of an atom), one can predict many properties of a given atom, as well as the properties of simple and complex substances, of which it is included.
The structure of the atom, i.e. the composition of the nucleus and the distribution of electrons around the nucleus, it is easy to determine by the position of the element in periodic system.
In the periodic system of D.I. Mendeleev, chemical elements are arranged in a certain sequence. This sequence is closely related to the structure of the atoms of these elements. Each chemical element in the system is assigned serial number, in addition, for it you can specify the period number, group number, subgroup type.
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Knowing the exact "address" of a chemical element - a group, subgroup and period number, one can unambiguously determine the structure of its atom.
Period is a horizontal row of chemical elements. There are seven periods in the modern periodic system. The first three periods small, because they contain 2 or 8 elements:
1st period - H, He - 2 elements;
2nd period - Li ... Ne - 8 elements;
3rd period - Na ... Ar - 8 elements.
Other periods - large. Each of them contains 2-3 rows of elements:
4th period (2 rows) - K ... Kr - 18 elements;
6th period (3 rows) - Cs ... Rn - 32 elements. This period includes a number of lanthanides.
Group is a vertical row of chemical elements. There are eight groups in total. Each group consists of two subgroups: main subgroup and secondary subgroup. For example:
The main subgroup is formed by chemical elements of small periods (for example, N, P) and large periods (for example, As, Sb, Bi).
A side subgroup is formed by chemical elements of only large periods (for example, V, Nb,
Ta).
Visually, these subgroups are easy to distinguish. The main subgroup is “high”, it starts from the 1st or 2nd period. The secondary subgroup is “low”, starting from the 4th period.
So, each chemical element of the periodic system has its own address: period, group, subgroup, ordinal number.
For example, vanadium V is a chemical element of the 4th period, group V, secondary subgroup, serial number 23.
Task 3.1. Specify the period, group and subgroup for chemical elements with serial numbers 8, 26, 31, 35, 54.
Task 3.2. Specify the serial number and name of the chemical element, if it is known that it is located:
a) in the 4th period, group VI, secondary subgroup;
b) in the 5th period, group IV, main subgroup.
How can information about the position of an element in the periodic system be related to the structure of its atom?
An atom is made up of a nucleus (positively charged) and electrons (negatively charged). In general, the atom is electrically neutral.
Positive charge of the nucleus of an atom equal to the atomic number of the chemical element.
The nucleus of an atom is a complex particle. Almost all the mass of an atom is concentrated in the nucleus. Since a chemical element is a collection of atoms with the same nuclear charge, the following coordinates are indicated near the symbol of the element:
Based on these data, the composition of the nucleus can be determined. The nucleus is made up of protons and neutrons.
Proton p has a mass of 1 (1.0073 amu) and a charge of +1. Neutron n it has no charge (neutral), and its mass is approximately equal to the mass of a proton (1.0087 amu).
The nuclear charge is determined by the protons. And the number of protons is(by size) charge of the nucleus of an atom, i.e. serial number.
Number of neutrons N determined by the difference between the quantities: "mass of the nucleus" BUT and "serial number" Z. So, for an aluminum atom:
N = BUT – Z = 27 –13 = 14n,
Task 3.3. Determine composition nuclei of atoms if the chemical element is in:
a) 3rd period, group VII, main subgroup;
b) 4th period, group IV, secondary subgroup;
c) 5th period, group I, main subgroup.
Attention! When determining the mass number of the nucleus of an atom, it is necessary to round off the atomic mass indicated in the periodic system. This is done because the masses of the proton and neutron are practically integer, and the mass of electrons can be neglected.
Let us determine which of the nuclei below belong to the same chemical element:
A (20 R + 20n),
B (19 R + 20n),
IN 20 R + 19n).
Atoms of the same chemical element have nuclei A and B, since they contain the same number of protons, i.e., the charges of these nuclei are the same. Studies show that the mass of an atom does not significantly affect its Chemical properties.
Isotopes are called atoms of the same chemical element (the same number of protons), differing in mass ( different number neutrons).
Isotopes and their chemical compounds differ from each other in physical properties, but the chemical properties of isotopes of one chemical element are the same. Thus, isotopes of carbon-14 (14 C) have the same chemical properties as carbon-12 (12 C), which enter the tissues of any living organism. The difference is manifested only in radioactivity (isotope 14 C). Therefore, isotopes are used for the diagnosis and treatment of various diseases, for scientific research.
Let us return to the description of the structure of the atom. As you know, the nucleus of an atom does not change in chemical processes. What is changing? The variable turns out to be total number electrons in an atom and the distribution of electrons. General number of electrons in a neutral atom it is easy to determine - it is equal to the serial number, i.e. charge of the nucleus of an atom:
Electrons have a negative charge of -1, and their mass is negligible: 1/1840 of the mass of a proton.
Negatively charged electrons repel each other and are at different distances from the nucleus. Wherein electrons that have an approximately equal amount of energy are at approximately equal distance away from the nucleus and form an energy level.
The number of energy levels in an atom is equal to the number of the period in which the chemical element is located. Energy levels are conventionally designated as follows (for example, for Al):
Task 3.4. Determine the number of energy levels in the atoms of oxygen, magnesium, calcium, lead.
Each energy level can contain a limited number of electrons:
On the first - no more than two electrons;
On the second - no more than eight electrons;
On the third - no more than eighteen electrons.
These numbers show that, for example, the second energy level can have 2, 5, or 7 electrons, but not 9 or 12 electrons.
It is important to know that regardless of the energy level number on external level(last) cannot be more than eight electrons. The outer eight-electron energy level is the most stable and is called complete. Such energy levels are found in the most inactive elements - the noble gases.
How to determine the number of electrons in the outer level of the remaining atoms? There is a simple rule for this: number of outer electrons equals:
For elements of the main subgroups - the number of the group;
For elements of secondary subgroups, it cannot be more than two.
For example (Fig. 5):
Task 3.5. Specify the number of external electrons for chemical elements with serial numbers 15, 25, 30, 53.
Task 3.6. Find chemical elements in the periodic table, in the atoms of which there is a completed external level.
It is very important to correctly determine the number of external electrons, because It is with them that the most important properties of the atom are associated. Yes, in chemical reactions atoms tend to acquire a stable, complete external level (8 e). Therefore, atoms, on the outer level of which there are few electrons, prefer to give them away.
Chemical elements whose atoms can only donate electrons are called metals. Obviously, there should be few electrons at the outer level of the metal atom: 1, 2, 3.
If there are many electrons on the external energy level of an atom, then such atoms tend to accept electrons before the completion of the external energy level, i.e. up to eight electrons. Such elements are called non-metals.
Question. Do the chemical elements of the secondary subgroups belong to metals or non-metals? Why?
Answer. Metals and non-metals of the main subgroups in the periodic table are separated by a line that can be drawn from boron to astatine. Above this line (and on the line) are non-metals, below - metals. All elements of secondary subgroups are below this line.
Task 3.7. Determine whether metals or non-metals include: phosphorus, vanadium, cobalt, selenium, bismuth. Use the position of the element in the periodic table of chemical elements and the number of electrons in the outer level.
In order to compose the distribution of electrons over the remaining levels and sublevels, the following algorithm should be used.
1. Determine the total number of electrons in the atom (by serial number).
2. Determine the number of energy levels (by period number).
3. Determine the number of external electrons (according to the type of subgroup and group number).
4. Indicate the number of electrons at all levels except the penultimate one.
For example, according to points 1–4 for the manganese atom, it is determined:
Total 25 e; distributed (2 + 8 + 2) = 12 e; so, on the third level is: 25 - 12 = 13 e.
The distribution of electrons in the manganese atom was obtained:
Task 3.8. Work out the algorithm by drawing up atomic structure diagrams for elements No. 16, 26, 33, 37. Indicate whether they are metals or non-metals. Explain the answer.
When compiling the above diagrams of the structure of the atom, we did not take into account that the electrons in the atom occupy not only levels, but also certain sublevels each level. Types of sublevels are indicated by Latin letters: s, p, d.
The number of possible sublevels is equal to the level number. The first level consists of one
s-sublevel. The second level consists of two sublevels - s and R. The third level - from three sublevels - s, p and d.
Each sublevel can contain a strictly limited number of electrons:
at the s-sublevel - no more than 2e;
at the p-sublevel - no more than 6e;
at the d-sublevel - no more than 10e.
Sublevels of one level are filled in a strictly defined order: spd.
Thus, R- sublevel can't start to fill if not full s-sublevel of a given energy level, etc. Based on this rule, it is easy to compose the electronic configuration of the manganese atom:
Generally electronic configuration of an atom manganese is written like this:
25 Mn 1 s 2 2s 2 2p 6 3s 2 3p 6 3d 5 4s 2 .
Task 3.9. Make electronic configurations of atoms for chemical elements No. 16, 26, 33, 37.
Why is it necessary to make electronic configurations of atoms? To determine the properties of these chemical elements. It should be remembered that only valence electrons.
Valence electrons are in the outer energy level and incomplete
d-sublevel of the pre-outer level.
Let's determine the number of valence electrons for manganese:
or abbreviated: Mn ... 3 d 5 4s 2 .
What can be determined by the formula for the electronic configuration of an atom?
1. What element is it - metal or non-metal?
Manganese is a metal, because the outer (fourth) level contains two electrons.
2. What process is typical for metal?
Manganese atoms always donate electrons in reactions.
3. What electrons and how many will give a manganese atom?
In reactions, the manganese atom gives up two outer electrons (they are farthest from the nucleus and are weaker attracted by it), as well as five pre-outer d-electrons. The total number of valence electrons is seven (2 + 5). In this case, eight electrons will remain at the third level of the atom, i.e. complete outer level is formed.
All these reasoning and conclusions can be reflected using the scheme (Fig. 6):
The resulting conditional charges of an atom are called oxidation states.
Considering the structure of the atom, in a similar way it can be shown that the typical oxidation states for oxygen are -2, and for hydrogen +1.
Question. With which of the chemical elements can manganese form compounds, if we take into account the degrees of its oxidation obtained above?
Answer: Only with oxygen, tk. its atom has the opposite charge in its oxidation state. The formulas of the corresponding manganese oxides (here the oxidation states correspond to the valences of these chemical elements):
The structure of the manganese atom suggests that manganese cannot have a higher degree of oxidation, because in this case, one would have to touch upon the stable, now completed, pre-outer level. Therefore, the +7 oxidation state is the highest, and the corresponding Mn 2 O 7 oxide is the highest manganese oxide.
To consolidate all these concepts, consider the structure of the tellurium atom and some of its properties:
As a non-metal, the Te atom can accept 2 electrons before the completion of the outer level and donate "extra" 6 electrons:
Task 3.10. Draw the electronic configurations of Na, Rb, Cl, I, Si, Sn atoms. Determine the properties of these chemical elements, the formulas of their simplest compounds (with oxygen and hydrogen).
Practical Conclusions
1. Only valence electrons participate in chemical reactions, which can only be in the last two levels.
2. Metal atoms can only donate valence electrons (all or a few), taking positive oxidation states.
3. Non-metal atoms can accept electrons (missing - up to eight), while acquiring negative oxidation states, and donate valence electrons (all or a few), while they acquire positive oxidation states.
Let us now compare the properties of the chemical elements of one subgroup, for example, sodium and rubidium:
Na...3 s 1 and Rb...5 s 1 .
What is common in the structure of the atoms of these elements? At the outer level of each atom, one electron is active metals. metal activity associated with the ability to donate electrons: the easier an atom gives off electrons, the more pronounced its metallic properties.
What holds electrons in an atom? attraction to the nucleus. The closer the electrons are to the nucleus, the stronger they are attracted by the nucleus of the atom, the more difficult it is to “tear them off”.
Based on this, we will answer the question: which element - Na or Rb - gives away an external electron more easily? Which element is the more active metal? Obviously, rubidium, because its valence electrons are farther away from the nucleus (and are less strongly held by the nucleus).
Conclusion. In the main subgroups, from top to bottom, the metallic properties are enhanced, because the radius of the atom increases, and valence electrons are weaker attracted to the nucleus.
Let's compare the properties of chemical elements of group VIIa: Cl …3 s 2 3p 5 and I...5 s 2 5p 5 .
Both chemical elements are non-metals, because. one electron is missing before the completion of the outer level. These atoms will actively attract the missing electron. Moreover, the stronger the missing electron attracts the nonmetal atom, the stronger its manifestations are. non-metallic properties(ability to accept electrons).
What causes the attraction of an electron? Due to the positive charge of the nucleus of the atom. In addition, the closer the electron to the nucleus, the stronger their mutual attraction, the more active the non-metal.
Question. Which element has more pronounced non-metallic properties: chlorine or iodine?
Answer: Obviously, chlorine, because. its valence electrons are closer to the nucleus.
Conclusion. The activity of non-metals in subgroups decreases from top to bottom, because the radius of the atom increases and it is more and more difficult for the nucleus to attract the missing electrons.
Let us compare the properties of silicon and tin: Si …3 s 2 3p 2 and Sn…5 s 2 5p 2 .
Both atoms have four electrons at the outer level. Nevertheless, these elements in the periodic table are on opposite sides of the line connecting boron and astatine. Therefore, for silicon, the symbol of which is above the B–At line, nonmetallic properties are more pronounced. On the contrary, tin, whose symbol is below the B–At line, has stronger metallic properties. This is due to the fact that in the tin atom, four valence electrons are removed from the nucleus. Therefore, the attachment of the missing four electrons is difficult. At the same time, the return of electrons from the fifth energy level occurs quite easily. For silicon, both processes are possible, with the first (acceptance of electrons) predominating.
Conclusions on chapter 3. The fewer external electrons in an atom and the farther they are from the nucleus, the stronger the metallic properties are manifested.
The more external electrons in an atom and the closer they are to the nucleus, the more non-metallic properties are manifested.
Based on the conclusions formulated in this chapter, for any chemical element of the periodic system, you can make a "characteristic".
Property Description Algorithm
chemical element by its position
in the periodic system
1. Draw up a diagram of the structure of the atom, i.e. determine the composition of the nucleus and the distribution of electrons by energy levels and sublevels:
Determine the total number of protons, electrons and neutrons in an atom (by serial number and relative atomic mass);
Determine the number of energy levels (by period number);
Determine the number of external electrons (by type of subgroup and group number);
Indicate the number of electrons at all energy levels except the penultimate one;
2. Determine the number of valence electrons.
3. Determine which properties - metal or non-metal - are more pronounced for a given chemical element.
4. Determine the number of given (received) electrons.
5. Determine the highest and lowest oxidation states of a chemical element.
6. Compose for these oxidation states chemical formulas the simplest compounds with oxygen and hydrogen.
7. Determine the nature of the oxide and write an equation for its reaction with water.
8. For the substances specified in paragraph 6, draw up equations characteristic reactions(see chapter 2).
Task 3.11. According to the above scheme, make descriptions of the atoms of sulfur, selenium, calcium and strontium and the properties of these chemical elements. What kind general properties exhibit their oxides and hydroxides?
If you have completed exercises 3.10 and 3.11, then it is easy to see that not only the atoms of the elements of one subgroup, but also their compounds have common properties and a similar composition.
Periodic law of D.I. Mendeleev:the properties of chemical elements, as well as the properties of simple and complex substances formed by them, are in a periodic dependence on the charge of the nuclei of their atoms.
The physical meaning of the periodic law: the properties of chemical elements are periodically repeated because the configurations of valence electrons (the distribution of electrons of the outer and penultimate levels) are periodically repeated.
So, the chemical elements of the same subgroup have the same distribution of valence electrons and, therefore, similar properties.
For example, the chemical elements of the fifth group have five valence electrons. At the same time, in the atoms of chemical elements of the main subgroups- all valence electrons are in the outer level: ... ns 2 np 3 , where n– period number.
At atoms elements of secondary subgroups only 1 or 2 electrons are in the outer level, the rest are in d- sublevel of the pre-external level: ... ( n – 1)d 3 ns 2 , where n– period number.
Task 3.12. Make short electronic formulas for atoms of chemical elements No. 35 and 42, and then make up the distribution of electrons in these atoms according to the algorithm. Make sure your prediction comes true.
Exercises for chapter 3
1. Formulate the definitions of the concepts "period", "group", "subgroup". What do the chemical elements that make up: a) period; b) a group; c) subgroup?
2. What are isotopes? What properties - physical or chemical - do isotopes have in common? Why?
3. Formulate the periodic law of DIMendeleev. Explain it physical meaning and illustrate with examples.
4. What are the metallic properties of chemical elements? How do they change in a group and in a period? Why?
5. What are the non-metallic properties of chemical elements? How do they change in a group and in a period? Why?
6. Make brief electronic formulas of chemical elements No. 43, 51, 38. Confirm your assumptions by describing the structure of the atoms of these elements according to the above algorithm. Specify the properties of these elements.
7. By short electronic formulas
a) ...4 s 2 4p 1 ;
b) …4 d 1 5s 2 ;
in 3 d 5 4s 1
determine the position of the corresponding chemical elements in the periodic system of D.I. Mendeleev. Name these chemical elements. Confirm your assumptions with a description of the structure of atoms of these chemical elements according to the algorithm. Specify the properties of these chemical elements.
To be continued
Rice. 7. Image shapes and orientations
s-,p-,d-, orbitals using boundary surfaces.
Quantum numberm l called magnetic . It determines the spatial arrangement of the atomic orbital and takes integer values from - l to + l through zero, that is 2 l+ 1 values (Table 27).
Orbitals of the same sublevel ( l= const) have the same energy. Such a state is called degenerate in energy. So p-orbital - three times, d- five times, and f are seven times degenerate. Boundary surfaces s-,p-,d-, orbitals are shown in fig. 7.
s -orbitals spherically symmetrical for any n and differ from each other only by the size of the sphere. Their maximally symmetrical shape is due to the fact that at l= 0 and μ l = 0.
Table 27
Number of orbitals on energy sublevels
|
Orbital quantum number |
Magnetic quantum number |
Number of orbitals with a given value l |
|
m l | ||
|
–2, –1, 0, +1, +2 | ||
|
–3, –2, –1, 0, +1, +2, +3 |
p -orbitals exist at n≥ 2 and l= 1, so there are three possible orientations in space: m l= -1, 0, +1. All p-orbitals have a nodal plane dividing the orbital into two regions, so the boundary surfaces are dumbbell-shaped, oriented in space at an angle of 90° relative to each other. The axes of symmetry for them are the coordinate axes, which are denoted p x , p y , p z .
d -orbitals determined by the quantum number l = 2 (n≥ 3), at which m l= –2, –1, 0, +1, +2, that is, they are characterized by five variants of orientation in space. d-orbitals oriented with blades along the coordinate axes are denoted d z² and d x ²– y², and oriented by the blades along the bisectors of the coordinate angles - d xy , d yz , d xz .
Seven f -orbitals corresponding l = 3 (n≥ 4) are shown as boundary surfaces.
quantum numbers n, l and m do not fully characterize the state of an electron in an atom. It has been experimentally established that the electron has one more property - spin. Simplistically, spin can be represented as the rotation of an electron around its own axis. Spin quantum number m s has only two meanings m s= ±1/2, which are two projections of the angular momentum of the electron on the selected axis. electrons with different m s indicated by arrows pointing up and down.
The sequence of filling atomic orbitals
The population of atomic orbitals (AO) with electrons is carried out according to the principle of least energy, the Paulia principle, the Hund rule, and for many-electron atoms, the Klechkovsky rule.
The principle of least energy requires that electrons populate the AO in order of increasing electron energy in these orbitals. This reflects the general rule - the maximum stability of the system corresponds to the minimum of its energy.
Principle pauli (1925) forbids electrons with the same set of quantum numbers to be in a multi-electron atom. This means that any two electrons in an atom (or molecule, or ion) must differ from each other by the value of at least one quantum number, that is, there can be no more than two electrons with different spins (paired electrons) in one orbital. Each sublevel contains 2 l+ 1 orbitals containing no more than 2(2 l+ 1) electrons. It follows from this that the capacitance s-orbitals - 2, p-orbitals - 6, d-orbitals - 10 and f-orbitals - 14 electrons. If the number of electrons for a given l sum from 0 to n– 1, then we get the formula Bora–Bury, which determines the total number of electrons in a level with a given n:
This formula does not take into account the interelectronic interaction and ceases to be valid when n ≥ 3.
Orbitals with the same energy (degenerate) are filled according to rule Gunda : the electron configuration with the maximum spin has the lowest energy. This means that if there are three electrons in the p-orbital, then they are arranged as follows: , and the total spin S=3/2, not like this: , S=1/2.
Klechkovsky's rule (principle of least energy). In multielectron atoms, as in the hydrogen atom, the state of the electron is determined by the values of the same four quantum numbers, but in this case the electron is not only in the field of the nucleus, but also in the field of other electrons. Therefore, the energy in many-electron atoms is determined not only by the main, but also by the orbital quantum number, or rather, their sum: the energy of atomic orbitals increases as the sum increasesn + l; with the same amount, the level with the smaller one is filled firstnand bigl. The energy of atomic orbitals increases according to the series:
|
1s<2s<2p<3s<3p<4s≈3d<4p<5s≈4d<5p<6s≈4f≈5d<6p<7s≈5f≈6d<7p. |
So, four quantum numbers describe the state of an electron in an atom and characterize the energy of the electron, its spin, the shape of the electron cloud and its orientation in space. When an atom passes from one state to another, the electron cloud is restructured, that is, the values of quantum numbers change, which is accompanied by absorption or emission of energy quanta by the atom.
Malyugin 14. External and internal energy levels. Completion of the energy level.
Let us briefly recall what we already know about the structure of the electron shell of atoms:
ü the number of energy levels of the atom = the number of the period in which the element is located;
ü the maximum capacity of each energy level is calculated by the formula 2n2
ü the outer energy shell cannot contain more than 2 electrons for elements of the 1st period, more than 8 electrons for elements of other periods
Once again, let us return to the analysis of the scheme for filling energy levels in elements of small periods:
Table 1. Filling of energy levels
for elements of small periods
Period number | Number of energy levels = period number | Element symbol, its ordinal number | Total electrons | Distribution of electrons by energy levels | Group number |
|
|
H +1 )1 | +1 H, 1e- | |||||
|
He + 2 ) 2 | +2 No, 2nd | |||||
|
Li + 3 ) 2 ) 1 | + 3 Li, 2e-, 1e- | |||||
|
Be +4 ) 2 )2 | + 4 Be, 2e-,2 e- | |||||
|
B +5 ) 2 )3 | +5 B, 2e-, 3e- | |||||
|
C +6 ) 2 )4 | +6 C, 2e-, 4e- | |||||
|
N + 7 ) 2 ) 5 | + 7 N, 2e-,5 e- | |||||
|
O + 8 ) 2 ) 6 | + 8 O, 2e-,6 e- | |||||
|
F + 9 ) 2 ) 7 | + 9 F, 2e-,7 e- | |||||
|
Ne + 10 ) 2 ) 8 | + 10 Ne, 2e-,8 e- | |||||
|
Na + 11 ) 2 ) 8 )1 | +1 1 Na, 2e-, 8e-, 1e- | |||||
|
mg + 12 ) 2 ) 8 )2 | +1 2 mg, 2e-, 8e-, 2 e- | |||||
|
Al + 13 ) 2 ) 8 )3 | +1 3 Al, 2e-, 8e-, 3 e- | |||||
|
Si + 14 ) 2 ) 8 )4 | +1 4 Si, 2e-, 8e-, 4 e- | |||||
|
P + 15 ) 2 ) 8 )5 | +1 5 P, 2e-, 8e-, 5 e- | |||||
|
S + 16 ) 2 ) 8 )6 | +1 5 P, 2e-, 8e-, 6 e- | |||||
|
Cl + 17 ) 2 ) 8 )7 | +1 7 Cl, 2e-, 8e-, 7 e- | |||||
18 Ar |
Ar+ 18 ) 2 ) 8 )8 | +1 8 Ar, 2e-, 8e-, 8 e- |
Analyze table 1. Compare the number of electrons in the last energy level and the number of the group in which the chemical element is located.
Have you noticed that the number of electrons in the outer energy level of atoms is the same as the group number, in which the element is located (the exception is helium)?
!!! This rule is true only for elements major subgroups.
Each period of the system ends with an inert element(helium He, neon Ne, argon Ar). The external energy level of these elements contains the maximum possible number of electrons: helium -2, the remaining elements - 8. These are elements of group VIII of the main subgroup. The energy level similar to the structure of the energy level of an inert gas is called completed. This is a kind of strength limit of the energy level for each element of the Periodic system. Molecules of simple substances - inert gases, consist of one atom and are distinguished by chemical inertness, i.e., they practically do not enter into chemical reactions.
For the remaining elements of the PSCE, the energy level differs from the energy level of the inert element, such levels are called unfinished. The atoms of these elements strive to complete the outer energy level by donating or accepting electrons.
Questions for self-control
1. What energy level is called external?
2. What energy level is called internal?
3. What energy level is called complete?
4. Elements of which group and subgroup have a completed energy level?
5. What is the number of electrons in the outer energy level of the elements of the main subgroups?
6. How are the elements of one main subgroup similar in the structure of the electronic level
7. How many electrons at the outer level contain the elements of a) group IIA;
b) IVA group; c) Group VII A
View answer
1. Last
2. Any but the last
3. The one that contains the maximum number of electrons. As well as the outer level, if it contains 8 electrons for period I - 2 electrons.
4. Elements of group VIIIA (inert elements)
5. The number of the group in which the element is located
6. All elements of the main subgroups on the external energy level contain as many electrons as the group number
7. a) the elements of group IIA have 2 electrons in the outer level; b) group IVA elements have 4 electrons; c) elements of group VII A have 7 electrons.
Tasks for independent solution
1. Determine the element according to the following criteria: a) it has 2 electronic levels, on the outer - 3 electrons; b) has 3 electronic levels, on the outer - 5 electrons. Write down the distribution of electrons over the energy levels of these atoms.
2. What two atoms have the same number of filled energy levels?
View answer:
1. a) Let's establish the "coordinates" of the chemical element: 2 electronic levels - II period; 3 electrons at the outer level - III A group. This is a 5B bur. Scheme of distribution of electrons by energy levels: 2e-, 3e-
b) III period, VA group, element phosphorus 15Р. Scheme of distribution of electrons by energy levels: 2e-, 8e-, 5e-
2. d) sodium and chlorine.
Explanation: a) sodium: +11 )2)8 )1 (filled 2) ←→ hydrogen: +1)1
b) helium: +2 )2 (filled 1) ←→ hydrogen: hydrogen: +1)1
c) helium: +2 )2 (filled 1) ←→ neon: +10 )2)8 (filled 2)
*G) sodium: +11 )2)8 )1 (filled 2) ←→ chlorine: +17 )2)8 )7 (filled 2)
4. Ten. Number of electrons = serial number
5 c) arsenic and phosphorus. Atoms located in the same subgroup have the same number of electrons.
Explanations:
a) sodium and magnesium (in different groups); b) calcium and zinc (in the same group, but different subgroups); * c) arsenic and phosphorus (in one, main, subgroup) d) oxygen and fluorine (in different groups).
7. d) the number of electrons in the outer level
8. b) the number of energy levels
9. a) lithium (located in group IA of period II)
10. c) silicon (IVA group, III period)
11. b) boron (2 levels - IIperiod, 3 electrons in the outer level - IIIAGroup)
2. The structure of nuclei and electron shells of atoms
2.6. Energy levels and sublevels
The most important characteristic of the state of an electron in an atom is the energy of the electron, which, according to the laws of quantum mechanics, does not change continuously, but abruptly, i.e. can only take on well-defined values. Thus, we can speak about the presence of a set of energy levels in the atom.
Energy level- set of AO with close energy values.
Energy levels are numbered with principal quantum number n, which can only take positive integer values (n = 1, 2, 3, ...). The larger the value of n, the higher the energy of the electron and the given energy level. Each atom contains an infinite number of energy levels, some of which are populated by electrons in the ground state of the atom, and some are not (these energy levels are populated in the excited state of the atom).
Electronic layer- a set of electrons that are at a given energy level.
In other words, an electron layer is an energy level containing electrons.
The set of electron layers forms the electron shell of an atom.
Within the same electron layer, electrons can differ somewhat in energy, and therefore they say that energy levels are split into energy sublevels(sublayers). The number of sublevels into which a given energy level is split is equal to the number of the main quantum number of the energy level:
N (subur) \u003d n (level) . (2.4)
Sublevels are depicted using numbers and letters: the number corresponds to the number of the energy level (electronic layer), the letter corresponds to the nature of the AO that forms the sublevels (s -, p -, d -, f -), for example: 2p - sublevel (2p - AO, 2p -electron).
Thus, the first energy level (Fig. 2.5) consists of one sublevel (1s), the second - of two (2s and 2p), the third - of three (3s, 3p and 3d), the fourth of four (4s, 4p, 4d and 4f ), etc. Each sublevel contains a certain number of AO:
N (AO) = n 2 . (2.5)
Rice. 2.5. Scheme of energy levels and sublevels for the first three electron layers
1. s-type AOs are present at all energy levels, p-type appear starting from the second energy level, d-type - from the third, f-type - from the fourth, etc.
2. At a given energy level, there can be one s -, three p -, five d -, seven f -orbitals.
3. The larger the main quantum number, the larger the size of the AO.
Since there cannot be more than two electrons on one AO, the total (maximum) number of electrons at a given energy level is 2 times greater than the number of AOs and is equal to:
N (e) = 2n 2 . (2.6)
Thus, at a given energy level, there can be a maximum of 2 s-type electrons, 6 p-type electrons and 10 d-type electrons. In total, at the first energy level, the maximum number of electrons is 2, at the second - 8 (2 s-type and 6 p-type), at the third - 18 (2 s-type, 6 p-type and 10 d-type). These findings are conveniently summarized in Table 1. 2.2.
Table 2.2
The relationship between the principal quantum number, the number e
