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. At the 3d sublevel, the maximum number of cells (four) is 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 chemical elements and determines basically the properties of their atoms. Therefore, these levels are called valence. 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 valence chemical element - is the number of equivalents that a given atom can attach to itself, or is the number of equivalents in an 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 valency 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 valency 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 is 8. For example, the highest oxidation state of nitrogen (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 the series of analogous elements electronic structures are only similar, but not identical, then in the transition from one element in the group to another, for them there is not a simple repetition of properties, 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 enters, 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, while chlorine is 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
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Anyone who went to school remembers that one of the required subjects to study was chemistry. She could like it, or she could not like it - it does not matter. And it is likely that much knowledge in this discipline has already been forgotten and is not applied in life. However, everyone probably remembers the table of chemical elements of D. I. Mendeleev. For many, it has remained a multi-colored table, where certain letters are inscribed in each square, denoting the names of chemical elements. But here we will not talk about chemistry as such, and describe hundreds of chemical reactions and processes, but we will talk about how the periodic table appeared in general - this story will be of interest to any person, and indeed to all those who want interesting and useful information .
A little background
Back in 1668, the outstanding Irish chemist, physicist and theologian Robert Boyle published a book in which many myths about alchemy were debunked, and in which he talked about the need to search for indecomposable chemical elements. The scientist also gave a list of them, consisting of only 15 elements, but allowed the idea that there may be more elements. This became the starting point not only in the search for new elements, but also in their systematization.
A hundred years later, the French chemist Antoine Lavoisier compiled a new list, which already included 35 elements. 23 of them were later found to be indecomposable. But the search for new elements continued by scientists around the world. And the main role in this process was played by the famous Russian chemist Dmitry Ivanovich Mendeleev - he was the first to put forward the hypothesis that there could be a relationship between the atomic mass of elements and their location in the system.
Thanks to painstaking work and comparison of chemical elements, Mendeleev was able to discover a relationship between elements in which they can be one, and their properties are not something taken for granted, but are a periodically repeating phenomenon. As a result, in February 1869, Mendeleev formulated the first periodic law, and already in March, his report “The relationship of properties with the atomic weight of elements” was submitted to the Russian Chemical Society by the historian of chemistry N. A. Menshutkin. Then in the same year, Mendeleev's publication was published in the journal Zeitschrift fur Chemie in Germany, and in 1871 a new extensive publication of the scientist dedicated to his discovery was published by another German journal Annalen der Chemie.
Creating a Periodic Table
By 1869, the main idea had already been formed by Mendeleev, and in a fairly short time, but he could not formalize it into any sort of ordered system that clearly displays what was what, for a long time he could not. In one of the conversations with his colleague A. A. Inostrantsev, he even said that everything had already worked out in his head, but he could not bring everything to the table. After that, according to Mendeleev's biographers, he began painstaking work on his table, which lasted three days without a break for sleep. All sorts of ways to organize the elements in a table were sorted out, and the work was complicated by the fact that at that time science did not yet know about all the chemical elements. But, despite this, the table was still created, and the elements were systematized.
Legend of Mendeleev's dream
Many have heard the story that D. I. Mendeleev dreamed of his table. This version was actively distributed by the aforementioned colleague of Mendeleev, A. A. Inostrantsev, as a funny story with which he entertained his students. He said that Dmitry Ivanovich went to bed and in a dream he clearly saw his table, in which all the chemical elements were arranged in the right order. After that, the students even joked that 40° vodka was discovered in the same way. But there were still real prerequisites for the sleep story: as already mentioned, Mendeleev worked on the table without sleep and rest, and Inostrantsev once found him tired and exhausted. In the afternoon, Mendeleev decided to take a break, and some time later, he woke up abruptly, immediately took a piece of paper and depicted a ready-made table on it. But the scientist himself refuted this whole story with a dream, saying: “I’ve been thinking about it for maybe twenty years, and you think: I was sitting and suddenly ... it’s ready.” So the legend of the dream may be very attractive, but the creation of the table was only possible through hard work.
Further work
In the period from 1869 to 1871, Mendeleev developed the ideas of periodicity, to which the scientific community was inclined. And one of the important stages of this process was the understanding that any element in the system should be located based on the totality of its properties in comparison with the properties of other elements. Based on this, and also based on the results of research in the change of glass-forming oxides, the chemist managed to amend the values of the atomic masses of some elements, among which were uranium, indium, beryllium and others.
Of course, Mendeleev wanted to fill the empty cells that remained in the table as soon as possible, and in 1870 he predicted that chemical elements unknown to science would soon be discovered, the atomic masses and properties of which he was able to calculate. The first of these were gallium (discovered in 1875), scandium (discovered in 1879) and germanium (discovered in 1885). Then the forecasts continued to be realized, and eight more new elements were discovered, including: polonium (1898), rhenium (1925), technetium (1937), francium (1939) and astatine (1942-1943). By the way, in 1900, D. I. Mendeleev and the Scottish chemist William Ramsay came to the conclusion that the elements of the zero group should also be included in the table - until 1962 they were called inert, and after - noble gases.
Organization of the periodic system
The chemical elements in the table of D. I. Mendeleev are arranged in rows, in accordance with the increase in their mass, and the length of the rows is chosen so that the elements in them have similar properties. For example, noble gases such as radon, xenon, krypton, argon, neon, and helium do not easily react with other elements, and also have low chemical activity, which is why they are located in the far right column. And the elements of the left column (potassium, sodium, lithium, etc.) react perfectly with other elements, and the reactions themselves are explosive. To put it simply, within each column, the elements have similar properties, varying from one column to the next. All elements up to No. 92 are found in nature, and with No. 93 begin artificial elements that can only be created in the laboratory.
In its original version, the periodic system was understood only as a reflection of the order existing in nature, and there were no explanations why everything should be that way. And only when quantum mechanics appeared, the true meaning of the order of elements in the table became clear.
Creative Process Lessons
Speaking about what lessons of the creative process can be drawn from the entire history of the creation of the periodic table of D. I. Mendeleev, one can cite as an example the ideas of the English researcher in the field of creative thinking Graham Wallace and the French scientist Henri Poincaré. Let's take them briefly.
According to Poincaré (1908) and Graham Wallace (1926), there are four main stages in creative thinking:
- Training- the stage of formulating the main task and the first attempts to solve it;
- Incubation- the stage during which there is a temporary distraction from the process, but work on finding a solution to the problem is carried out at a subconscious level;
- insight- the stage at which the intuitive solution is found. Moreover, this solution can be found in a situation that is absolutely not relevant to the task;
- Examination- the stage of testing and implementation of the solution, at which the verification of this solution and its possible further development takes place.
As we can see, in the process of creating his table, Mendeleev intuitively followed these four stages. How effective this is can be judged by the results, i.e. because the table was created. And given that its creation was a huge step forward not only for chemical science, but for the whole of humanity, the above four stages can be applied both to the implementation of small projects and to the implementation of global plans. The main thing to remember is that not a single discovery, not a single solution to a problem can be found on its own, no matter how much we want to see them in a dream and no matter how much we sleep. In order to succeed, whether it is the creation of a table of chemical elements or the development of a new marketing plan, you need to have certain knowledge and skills, as well as skillfully use your potential and work hard.
We wish you success in your endeavors and successful implementation of your plans!
- First, we write down the sign of the chem. element, where below to the left of the sign we indicate its serial number.
- Further, by the number of the period (from which the element) we determine the number of energy levels and draw next to the sign of the chemical element such a number of arcs.
- Then, according to the group number, the number of electrons in the outer level is written under the arc.
- At the 1st level, the maximum possible is 2, at the second it is already 8, at the third - as many as 18. We begin to put numbers under the corresponding arcs.
- The number of electrons at the penultimate level must be calculated as follows: the number of already affixed electrons is subtracted from the serial number of the element.
- It remains to turn our circuit into an electronic formula:
- We write the chemical element and its serial number. The number shows the number of electrons in the atom.
- We make a formula. To do this, you need to find out the number of energy levels, the basis for determining the number of the period of the element is taken.
- We break the levels into sub-levels.
Below you can see an example of how to correctly compose electronic formulas of chemical elements.
- first we fill in the s-sublevel, and then the p-, d-b f-sublevels;
- according to the Klechkovsky rule, electrons fill orbitals in order of increasing energy of these orbitals;
- according to Hund's rule, electrons within one sublevel occupy free orbitals one at a time, and then form pairs;
- According to the Pauli principle, there are no more than 2 electrons in one orbital.
The electronic formula of a chemical element shows how many electron layers and how many electrons are contained in an atom and how they are distributed over the layers.
To compile the electronic formula of a chemical element, you need to look at the periodic table and use the information obtained for this element. The serial number of the element in the periodic table corresponds to the number of electrons in the atom. The number of electron layers corresponds to the period number, the number of electrons in the last electron layer corresponds to the group number.
It must be remembered that the first layer has a maximum of 2 1s2 electrons, the second - a maximum of 8 (two s and six p: 2s2 2p6), the third - a maximum of 18 (two s, six p, and ten d: 3s2 3p6 3d10).
For example, the electronic formula of carbon: C 1s2 2s2 2p2 (serial number 6, period number 2, group number 4).
Electronic formula of sodium: Na 1s2 2s2 2p6 3s1 (serial number 11, period number 3, group number 1).
To check the correctness of writing an electronic formula, you can look at the site www.alhimikov.net.
Drawing up an electronic formula of chemical elements at first glance may seem like a rather complicated task, but everything will become clear if you adhere to the following scheme:
- write the orbitals first
- we insert numbers in front of the orbitals that indicate the number of the energy level. Don't forget the formula for determining maximum number electrons at the energy level: N=2n2
And how to find out the number of energy levels? Just look at the periodic table: this number is equal to the number of the period in which this element is located.
- above the orbital icon we write a number that indicates the number of electrons that are in this orbital.
For example, the electronic formula for scandium would look like this.
The task of compiling the electronic formula of a chemical element is not the easiest.
So, the algorithm for compiling electronic formulas of elements is as follows:
Here are the electronic formulas of some chemical elements:
You need to compose the electronic formulas of chemical elements in this way: you need to look at the number of the element in the periodic table, thus finding out how many electrons it has. Then you need to find out the number of levels, which is equal to the period. Then the sublevels are written and filled in:
First of all, you need to determine the number of atoms according to the periodic table.
To compile an electronic formula, you will need the periodic system of Mendeleev. Find your chemical element there and look at the period - it will is equal to the number energy levels. The group number will correspond numerically to the number of electrons in the last level. The element number will be quantitatively equal to the number of its electrons. You also clearly need to know that there are a maximum of 2 electrons on the first level, 8 on the second, and 18 on the third.
These are the highlights. In addition, on the Internet (including our website) you can find information with a ready-made electronic formula for each element, so you can check yourself.
Compiling electronic formulas of chemical elements is a very complex process, you can’t do without special tables, and you need to use a whole bunch of formulas. To summarize, you need to go through these steps:
It is necessary to draw up an orbital diagram in which there will be a concept of the difference between electrons from each other. Orbitals and electrons are highlighted in the diagram.
Electrons are filled in levels, from bottom to top and have several sublevels.
So first we find out the total number of electrons of a given atom.
We fill in the formula according to a certain scheme and write it down - this will be the electronic formula.
For example, for Nitrogen, this formula looks like this, first we deal with electrons:
And write down the formula:
To understand the principle of compiling the electronic formula of a chemical element, first you need to determine the total number of electrons in the atom by the number in the periodic table. After that, you need to determine the number of energy levels, taking as a basis the number of the period in which the element is located.
After that, the levels are broken down into sublevels, which are filled with electrons, based on the Principle of Least Energy.
You can check the correctness of your reasoning by looking, for example, here.
By compiling the electronic formula of a chemical element, you can find out how many electrons and electron layers are in a particular atom, as well as the order in which they are distributed among the layers.
To begin with, we determine the serial number of the element according to the periodic table, it corresponds to the number of electrons. The number of electron layers indicates the period number, and the number of electrons in the last layer of the atom corresponds to the group number.
6.6. Features of the electronic structure of atoms of chromium, copper and some other elements
If you carefully looked at Appendix 4, you probably noticed that for atoms of some elements, the sequence of filling orbitals with electrons is violated. Sometimes these violations are called "exceptions", but this is not so - there are no exceptions to the laws of Nature!
The first element with such a violation is chromium. Let us consider in more detail its electronic structure (Fig. 6.16 a). The chromium atom has 4 s-sublevel is not two, as one would expect, but only one electron. But for 3 d-sublevel five electrons, but this sublevel is filled after 4 s-sublevel (see Fig. 6.4). To understand why this happens, let's look at what electron clouds are 3 d sublevel of this atom.
Each of the five 3 d-clouds in this case is formed by one electron. As you already know from § 4 of this chapter, the common electron cloud of these five electrons is spherical, or, as they say, spherically symmetrical. By the nature of the electron density distribution in different directions, it is similar to 1 s-EO. The energy of the sublevel whose electrons form such a cloud turns out to be lower than in the case of a less symmetrical cloud. In this case, the energy of orbitals 3 d-sublevel is equal to energy 4 s-orbitals. When the symmetry is broken, for example, when the sixth electron appears, the energy of the orbitals is 3 d-sublevel again becomes more than energy 4 s-orbitals. Therefore, the manganese atom again has a second electron for 4 s-AO.
Spherical symmetry has a common cloud of any sublevel filled with electrons both half and completely. The decrease in energy in these cases is of a general nature and does not depend on whether any sublevel is half or completely filled with electrons. And if so, then we must look for the next violation in the atom, in the electron shell of which the ninth “comes” last d-electron. Indeed, the copper atom has 3 d-sublevel 10 electrons, and 4 s- there is only one sublevel (Fig. 6.16 b).
The decrease in the energy of the orbitals of a fully or half-filled sublevel is the cause of a number of important chemical phenomena, some of which you will become familiar with.
6.7. Outer and valence electrons, orbitals and sublevels
In chemistry, the properties of isolated atoms, as a rule, are not studied, since almost all atoms, being part of various substances, form chemical bonds. Chemical bonds are formed during the interaction of the electron shells of atoms. For all atoms (except hydrogen), not all electrons take part in the formation of chemical bonds: for boron, three out of five electrons, for carbon, four out of six, and, for example, for barium, two out of fifty-six. These "active" electrons are called valence electrons.
Sometimes valence electrons are confused with external electrons, but they are not the same thing.
The electron clouds of outer electrons have the maximum radius (and the maximum value of the principal quantum number).
It is the outer electrons that take part in the formation of bonds in the first place, if only because when the atoms approach each other, the electron clouds formed by these electrons come into contact first of all. But along with them, part of the electrons can also take part in the formation of a bond. pre-external(penultimate) layer, but only if they have an energy not much different from the energy of the outer electrons. Both those and other electrons of the atom are valence. (In lanthanides and actinides, even some "pre-external" electrons are valence)
The energy of valence electrons is much greater than the energy of other electrons of the atom, and valence electrons differ much less in energy from each other.
Outer electrons are always valence only if the atom can form chemical bonds at all. So, both electrons of the helium atom are external, but they cannot be called valence, since the helium atom does not form any chemical bonds at all.
Valence electrons occupy valence orbitals, which in turn form valence sublevels.
As an example, consider an iron atom whose electronic configuration is shown in Fig. 6.17. Of the electrons of the iron atom, the maximum principal quantum number ( n= 4) have only two 4 s-electron. Therefore, they are the outer electrons of this atom. The outer orbitals of the iron atom are all orbitals with n= 4, and the outer sublevels are all the sublevels formed by these orbitals, that is, 4 s-,
4p-, 4d- and 4 f-EPU.
Outer electrons are always valence, therefore, 4 s-electrons of an iron atom are valence electrons. And if so, then 3 d-electrons with a slightly higher energy will also be valence. At the outer level of the iron atom, in addition to the filled 4 s-AO there are still free 4 p-, 4d- and 4 f-AO. All of them are external, but only 4 are valence R-AO, since the energy of the remaining orbitals is much higher, and the appearance of electrons in these orbitals is not beneficial for the iron atom.
So, the iron atom
external electronic level - the fourth,
outer sublevels - 4 s-, 4p-, 4d- and 4 f-EPU,
outer orbitals - 4 s-, 4p-, 4d- and 4 f-AO,
outer electrons - two 4 s-electron (4 s 2),
the outer electron layer is the fourth,
external electron cloud - 4 s-EO
valence sublevels - 4 s-, 4p-, and 3 d-EPU,
valence orbitals - 4 s-, 4p-, and 3 d-AO,
valence electrons - two 4 s-electron (4 s 2) and six 3 d-electrons (3 d 6).
Valence sublevels can be partially or completely filled with electrons, or they can remain free at all. With an increase in the charge of the nucleus, the energy values of all sublevels decrease, but due to the interaction of electrons with each other, the energy of different sublevels decreases with different "speed". The energy of fully filled d- and f-sublevels decreases so much that they cease to be valence.
As an example, consider the atoms of titanium and arsenic (Fig. 6.18).
In the case of titanium atom 3 d-EPU is only partially filled with electrons, and its energy is greater than the energy of 4 s-EPU, and 3 d-electrons are valence. At the arsenic atom 3 d-EPU is completely filled with electrons, and its energy is much less than energy 4 s-EPU, and therefore 3 d-electrons are not valence.
In these examples, we analyzed valence electronic configuration titanium and arsenic atoms.
The valence electronic configuration of an atom is depicted as valence electronic formula, or in the form energy diagram of valence sublevels.
VALENCE ELECTRONS, EXTERNAL ELECTRONS, VALENCE EPU, VALENCE AO, VALENCE ELECTRON CONFIGURATION OF THE ATOM, VALENCE ELECTRON FORMULA, VALENCE SUBLEVEL DIAGRAM.
1. On the energy diagrams you have compiled and in the full electronic formulas of the atoms Na, Mg, Al, Si, P, S, Cl, Ar, indicate the external and valence electrons. Write the valence electronic formulas of these atoms. On the energy diagrams, highlight the parts corresponding to the energy diagrams of the valence sublevels.
2. What is common between the electronic configurations of atoms a) Li and Na, B and Al, O and S, Ne and Ar; b) Zn and Mg, Sc and Al, Cr and S, Ti and Si; c) H and He, Li and O, K and Kr, Sc and Ga. What are their differences
3. How many valence sublevels are in the electron shell of an atom of each of the elements: a) hydrogen, helium and lithium, b) nitrogen, sodium and sulfur, c) potassium, cobalt and germanium
4. How many valence orbitals are completely filled at the atom of a) boron, b) fluorine, c) sodium?
5. How many orbitals with an unpaired electron does an atom have a) boron, b) fluorine, c) iron
6. How many free outer orbitals does a manganese atom have? How many free valences?
7. For the next lesson, prepare a strip of paper 20 mm wide, divide it into cells (20 × 20 mm), and apply a natural series of elements to this strip (from hydrogen to meitnerium).
8. In each cell, place the symbol of the element, its serial number and the valence electronic formula, as shown in fig. 6.19 (use appendix 4).
6.8. Systematization of atoms according to the structure of their electron shells
The systematization of chemical elements is based on the natural series of elements
and principle of similarity of electron shells their atoms.
You are already familiar with the natural range of chemical elements. Now let's get acquainted with the principle of similarity of electron shells.
Considering the valence electronic formulas of atoms in the NRE, it is easy to find that for some atoms they differ only in the values of the main quantum number. For example, 1 s 1 for hydrogen, 2 s 1 for lithium, 3 s 1 for sodium, etc. Or 2 s 2 2p 5 for fluorine, 3 s 2 3p 5 for chlorine, 4 s 2 4p 5 for bromine, etc. This means that outer areas clouds of valence electrons of such atoms are very similar in shape and differ only in size (and, of course, electron density). And if so, then the electron clouds of such atoms and their corresponding valence configurations can be called similar. For atoms of different elements with similar electronic configurations, we can write common valence electronic formulas: ns 1 in the first case and ns 2 np 5 in the second. Moving along the natural series of elements, one can find other groups of atoms with similar valence configurations.
Thus, in the natural series of elements, atoms with similar valence electronic configurations regularly occur.
This is the principle of similarity of electron shells.
Let us try to reveal the form of this regularity. To do this, we will use the natural series of elements you made.
NRE begins with hydrogen, whose valence electronic formula is 1 s one . In search of similar valence configurations, we cut the natural series of elements in front of elements with a common valence electronic formula ns 1 (that is, before lithium, before sodium, etc.). We have received so-called "periods" of elements. Let's add the resulting "periods" so that they become table rows (see Figure 6.20). As a result, only the atoms of the first two columns of the table will have such electronic configurations.
Let's try to achieve similarity of valence electronic configurations in other columns of the table. To do this, we cut out elements with numbers 58 - 71 and 90 -103 from the 6th and 7th periods (they have 4 f- and 5 f-sublevels) and place them under the table. The symbols of the remaining elements will be shifted horizontally as shown in the figure. After that, the atoms of the elements in the same column of the table will have similar valence configurations, which can be expressed in general valence electronic formulas: ns 1 , ns 2 , ns 2 (n–1)d 1 , ns 2 (n–1)d 2 and so on until ns 2 np 6. All deviations from the general valence formulas are explained by the same reasons as in the case of chromium and copper (see paragraph 6.6).
As you can see, using the NRE and applying the principle of similarity of electron shells, we managed to systematize the chemical elements. Such a system of chemical elements is called natural, as it is based solely on the laws of Nature. The table we received (Fig. 6.21) is one of the ways to graphically depict a natural system of elements and is called long period table of chemical elements.
PRINCIPLE OF SIMILARITY OF ELECTRONIC SHELLS, NATURAL SYSTEM OF CHEMICAL ELEMENTS ("PERIODIC" SYSTEM), TABLE OF CHEMICAL ELEMENTS.
6.9. Long period table of chemical elements
Let's get acquainted in more detail with the structure of the long-period table of chemical elements.
The rows of this table, as you already know, are called "periods" of the elements. Periods are numbered with Arabic numerals from 1 to 7. There are only two elements in the first period. The second and third periods, containing eight elements each, are called short periods. The fourth and fifth periods, containing 18 elements each, are called long periods. The sixth and seventh periods, containing 32 elements each, are called extra long periods.
The columns of this table are called groups elements. Group numbers are indicated by Roman numerals with Latin letters A or B.
The elements of some groups have their own common (group) names: elements of the IA group (Li, Na, K, Rb, Cs, Fr) - alkaline elements(or alkali metal elements); group IIA elements (Ca, Sr, Ba and Ra) - alkaline earth elements(or alkaline earth metal elements)(the name "alkali metals" and alkaline earth metals" refer to simple substances formed by the corresponding elements and should not be used as names of groups of elements); elements of group VIA (O, S, Se, Te, Po) - chalcogens, elements of group VIIA (F, Cl, Br, I, At) – halogens, elements of group VIIIA (He, Ne, Ar, Kr, Xe, Rn) – noble gas elements.(The traditional name "noble gases" also applies to simple substances)
Taken out usually in lower part table elements with serial numbers 58 - 71 (Ce - Lu) are called lanthanides("following lanthanum"), and elements with serial numbers 90 - 103 (Th - Lr) - actinides("following actinium"). There is a variant of the long-period table, in which the lanthanides and actinides are not cut out of the NRE, but remain in their places in extra-long periods. This table is sometimes called extra long period.
The long period table is divided into four block(or sections).
s-block includes elements of IA and IIA groups with common valence electronic formulas ns 1 and ns 2
(s-elements).
p-block includes elements from group IIIA to VIIIA with common valence electronic formulas from ns 2 np 1 to ns 2 np 6 (p-elements).
d-block includes elements from IIIB to IIB group with common valence electronic formulas from ns 2 (n–1)d 1 to ns 2 (n–1)d 10 (d-elements).
f-block includes lanthanides and actinides ( f-elements).
Elements s- and p-blocks form A-groups, and elements d-block - B-group of a system of chemical elements. All f-elements are formally included in group IIIB.
The elements of the first period - hydrogen and helium - are s-elements and can be placed in IA and IIA groups. But helium is more often placed in group VIIIA as the element with which the period ends, which is fully consistent with its properties (helium, like all other simple substances formed by elements of this group, is a noble gas). Hydrogen is often placed in group VIIA, since its properties are much closer to halogens than to alkaline elements.
Each of the periods of the system begins with an element that has a valence configuration of atoms ns 1 , since it is from these atoms that the formation of the next electron layer begins, and ends with an element with the valence configuration of atoms ns 2 np 6 (except for the first period). This makes it easy to identify groups of sublevels in the energy diagram that are filled with electrons at the atoms of each of the periods (Fig. 6.22). Do this work with all the sublevels shown in the copy you made of Figure 6.4. The sublevels highlighted in Figure 6.22 (except for fully filled d- and f-sublevels) are valence for atoms of all elements of a given period.
Appearance in periods s-, p-, d- or f-elements are fully consistent with the sequence of filling s-, p-, d- or f- sublevels of electrons. This feature of the system of elements allows, knowing the period and group, which includes a given element, to immediately write down its valence electronic formula.
LONG-PERIOD TABLE OF CHEMICAL ELEMENTS, BLOCKS, PERIODS, GROUPS, ALKALINE ELEMENTS, ALKALINE EARTH ELEMENTS, CHALCOGENES, HALOGENS, NOBLE GAS ELEMENTS, LANTHANOIDES, ACTINIDES.
Write down the general valence electronic formulas of the atoms of the elements a) IVA and IVB groups, b) IIIA and VIIB groups?
2. What is common between the electronic configurations of atoms of elements A and B groups? How do they differ?
3. How many groups of elements are included in a) s-block, b) R-block, c) d-block?
4. Continue Figure 30 in the direction of increasing the energy of the sublevels and select the groups of sublevels that are filled with electrons in the 4th, 5th and 6th periods.
5. List the valence sublevels of atoms a) calcium, b) phosphorus, c) titanium, d) chlorine, e) sodium. 6. Formulate how s-, p- and d-elements differ from each other.
7. Explain why an atom belongs to any element is determined by the number of protons in the nucleus, and not by the mass of this atom.
8. For atoms of lithium, aluminum, strontium, selenium, iron and lead, make valence, complete and abbreviated electronic formulas and draw energy diagrams of valence sublevels. 9. The atoms of which elements correspond to the following valence electronic formulas: 3 s 1 , 4s 1 3d 1 , 2s 2 2 p 6 , 5s 2 5p 2 , 5s 2 4d 2 ?
6.10. Types of electronic formulas of the atom. The algorithm for their compilation
For different purposes, we need to know either the full or valence configuration of an atom. Each of these electronic configurations can be represented both by a formula and by an energy diagram. I.e, complete electronic configuration of an atom expressed the full electronic formula of the atom, or full energy diagram of an atom. In its turn, valence electron configuration of an atom expressed valence(or, as it is often called, " short ") the electronic formula of the atom, or diagram of valence sublevels of an atom(Fig. 6.23).
Previously, we made electronic formulas of atoms using the ordinal numbers of the elements. At the same time, we determined the sequence of filling sublevels with electrons according to the energy diagram: 1 s, 2s,
2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p,
6s, 4f, 5d, 6p, 7s etc. And only by writing down the full electronic formula, we could also write down the valence formula.
It is more convenient to write the valence electronic formula of the atom, which is most often used, based on the position of the element in the system of chemical elements, according to the period-group coordinates.
Let's consider in detail how this is done for elements s-, p- and d-blocks.
For elements s-block valence electronic formula of an atom consists of three characters. In general, it can be written like this:
In the first place (in the place of a large cell) is the period number (equal to the main quantum number of these s-electrons), and on the third (in the superscript) - the number of the group (equal to the number of valence electrons). Taking as an example a magnesium atom (3rd period, group IIA), we get:
For elements p-block valence electronic formula of an atom consists of six symbols:
Here, in place of large cells, the period number is also put (equal to the main quantum number of these s- and p-electrons), and the group number (equal to the number of valence electrons) turns out to be equal to the sum of the superscripts. For the oxygen atom (2nd period, VIA group) we get:
2s 2 2p 4 .
Valence electronic formula of most elements d block can be written like this:
As in previous cases, here instead of the first cell, the period number is put (equal to the main quantum number of these s-electrons). The number in the second cell turns out to be one less, since the main quantum number of these d-electrons. The group number here is also equal to the sum of the indices. An example is the valence electronic formula of titanium (4th period, IVB group): 4 s 2 3d 2 .
The group number is equal to the sum of the indices and for the elements of the VIB group, but they, as you remember, on the valence s-sublevel has only one electron, and the general valence electronic formula ns 1 (n–1)d 5 . Therefore, the valence electronic formula, for example, of molybdenum (5th period) is 5 s 1 4d 5 .
It is also easy to compose the valence electronic formula of any element of the IB group, for example, gold (6th period)>–>6 s 1 5d 10 , but in this case you need to remember that d- the electrons of the atoms of the elements of this group still remain valence, and some of them can participate in the formation of chemical bonds.
The general valence electronic formula of atoms of group IIB elements is - ns 2 (n – 1)d ten . Therefore, the valence electronic formula, for example, of a zinc atom is 4 s 2 3d 10 .
The valence electronic formulas of the elements of the first triad (Fe, Co and Ni) also obey the general rules. Iron, an element of group VIIIB, has a valence electronic formula of 4 s 2 3d 6. The cobalt atom has one d-electron more (4 s 2 3d 7), while the nickel atom has two (4 s 2 3d 8).
Using only these rules for writing valence electronic formulas, it is impossible to compose the electronic formulas of atoms of some d-elements (Nb, Ru, Rh, Pd, Ir, Pt), since in them, due to the tendency to highly symmetric electron shells, the filling of valence sublevels with electrons has some additional features.
Knowing the valence electronic formula, one can also write down the complete electronic formula of the atom (see below).
Often, instead of cumbersome full electronic formulas, they write down abbreviated electronic formulas atoms. To compile them in the electronic formula, all the electrons of the atom except the valence ones are selected, their symbols are placed in square brackets and the part of the electronic formula corresponding to the electronic formula of the atom of the last element of the previous period (the element that forms the noble gas) is replaced by the symbol of this atom.
Examples of electronic formulas of different types are shown in Table 14.
Table 14 Examples of electronic formulas of atoms
Electronic formulas |
|||
abbreviated |
Valence |
||
1s 2 2s 2 2p 3 |
2s 2 2p 3 |
2s 2 2p 3 |
|
1s 2 2s 2 2p 6 3s 2 3p 5 |
3s 2 3p 5 |
3s 2 3p 5 |
|
1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 5 |
4s 2 3d 5 |
4s 2 3d 5 |
|
1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 3 |
4s 2 4p 3 |
4s 2 4p 3 |
|
1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 |
4s 2 4p 6 |
4s 2 4p 6 |
|
Algorithm for compiling electronic formulas of atoms (on the example of an iodine atom)
№ |
Operation |
Result |
|
Determine the coordinates of the atom in the table of elements. |
Period 5, group VIIA |
||
Write the valence electronic formula. |
5s 2 5p 5 |
||
Add the symbols of the inner electrons in the order in which they fill the sublevels. |
1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 5 |
||
Taking into account the decrease in the energy of completely filled d- and f- sublevels, write down the full electronic formula. |
|
||
Label the valence electrons. |
1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 4d 10 5s 2 5p 5 |
||
Select the electronic configuration of the preceding noble gas atom. |
|||
Write down the abbreviated electronic formula, combining in square brackets all non-valent electrons. |
5s 2 5p 5 |
Notes
1. For elements of the 2nd and 3rd periods, the third operation (without the fourth) immediately leads to a complete electronic formula.
2. (n – 1)d 10 - Electrons remain valence at the atoms of the elements of the IB group.
COMPLETE ELECTRONIC FORMULA, VALENCE ELECTRONIC FORMULA, abbreviated ELECTRONIC FORMULA, ALGORITHM FOR COMPOSING ELECTRONIC FORMULA OF ATOMS.
1. Compose the valence electronic formula of the atom of the element a) the second period of the third A group, b) the third period of the second A group, c) the fourth period of the fourth A group.
2. Make abbreviated electronic formulas of magnesium, phosphorus, potassium, iron, bromine and argon atoms.
6.11. Short Period Table of Chemical Elements
Over the more than 100 years that have passed since the discovery of the natural system of elements, several hundred of the most diverse tables have been proposed that graphically reflect this system. Of these, in addition to the long-period table, the so-called short-period table of elements of D. I. Mendeleev is most widely used. A short-period table is obtained from a long-period one, if the 4th, 5th, 6th and 7th periods are cut before the elements of the IB group, moved apart and the resulting rows are added in the same way as we added the periods before. The result is shown in figure 6.24.
The lanthanides and actinides are also placed under the main table here.
AT groups this table contains elements whose atoms have the same number of valence electrons no matter what orbitals these electrons are in. So, the elements chlorine (a typical element that forms a non-metal; 3 s 2 3p 5) and manganese (metal-forming element; 4 s 2 3d 5), not possessing the similarity of electron shells, fall here into the same seventh group. The need to distinguish between such elements makes it necessary to single out in groups subgroups: main- analogues of A-groups of the long-period table and side effects are analogues of B-groups. In Figure 34, the symbols of the elements of the main subgroups are shifted to the left, and the symbols of the elements of the secondary subgroups are shifted to the right.
True, such an arrangement of elements in the table also has its advantages, because it is the number of valence electrons that primarily determines the valence capabilities of an atom.
The long-period table reflects the laws of the electronic structure of atoms, the similarity and patterns of changes in the properties of simple substances and compounds by groups of elements, the regular change in a number of physical quantities characterizing atoms, simple substances and compounds throughout the system of elements, and much more. The short period table is less convenient in this respect.
SHORT-PERIOD TABLE, MAIN SUB-GROUPS, SECONDARY SUB-GROUPS.
1. Convert the long-period table you built from the natural series of elements into a short-period table. Carry out the reverse transformation.
2. Is it possible to make a general valence electronic formula of atoms of elements of one group of a short period table? Why?
6.12. Atom sizes. Orbital radii
.The atom has no clear boundaries. What is considered the size of an isolated atom? The nucleus of an atom is surrounded by an electron shell, and the shell consists of electron clouds. The size of the EO is characterized by a radius r oo. All clouds in the outer layer have approximately the same radius. Therefore, the size of an atom can be characterized by this radius. It is called orbital radius of an atom(r 0).
The values of the orbital radii of atoms are given in Appendix 5.
The radius of the EO depends on the charge of the nucleus and on which orbital the electron that forms this cloud is located. Consequently, the orbital radius of an atom also depends on these same characteristics.
Consider the electron shells of hydrogen and helium atoms. Both in the hydrogen atom and in the helium atom, electrons are located on 1 s-AO, and their clouds would have the same size if the charges of the nuclei of these atoms were the same. But the charge of the nucleus of a helium atom is twice that of the charge of the nucleus of a hydrogen atom. According to Coulomb's law, the force of attraction acting on each of the electrons of a helium atom is twice the force of attraction of an electron to the nucleus of a hydrogen atom. Therefore, the radius of a helium atom must be much smaller than the radius of a hydrogen atom. And there is: r 0 (He) / r 0 (H) \u003d 0.291 E / 0.529 E 0.55.
The lithium atom has an outer electron at 2 s-AO, that is, forms a cloud of the second layer. Naturally, its radius should be larger. Really: r 0 (Li) = 1.586 E.
The atoms of the remaining elements of the second period have external electrons (and 2 s, and 2 p) are placed in the same second electron layer, and the charge of the nucleus of these atoms increases with increasing serial number. Electrons are more strongly attracted to the nucleus, and, naturally, the radii of atoms decrease. We could repeat these arguments for the atoms of the elements of other periods, but with one clarification: the orbital radius monotonically decreases only when each of the sublevels is filled.
But if we ignore the particulars, then the general nature of the change in the size of atoms in a system of elements is as follows: with an increase in the serial number in a period, the orbital radii of atoms decrease, and in a group they increase. The largest atom is a cesium atom, and the smallest is a helium atom, but of the atoms of the elements that form chemical compounds (helium and neon do not form them), the smallest is a fluorine atom.
Most of the atoms of the elements, standing in the natural series after the lanthanides, have orbital radii somewhat smaller than one would expect, based on general laws. This is due to the fact that 14 lanthanides are located between lanthanum and hafnium in the system of elements, and, consequently, the nuclear charge of the hafnium atom is 14 e more than lanthanum. Therefore, the outer electrons of these atoms are attracted to the nucleus more strongly than they would be attracted in the absence of lanthanides (this effect is often called "lanthanide contraction").
Please note that when passing from atoms of elements of group VIIIA to atoms of elements of group IA, the orbital radius increases abruptly. Consequently, our choice of the first elements of each period (see § 7) turned out to be correct.
ORBITAL RADIUS OF THE ATOM, ITS CHANGE IN THE SYSTEM OF ELEMENTS.
1. According to the data given in Appendix 5, plot on graph paper the dependence of the orbital radius of the atom on the element's serial number for elements with Z from 1 to 40. The length of the horizontal axis is 200 mm, the length of the vertical axis is 100 mm.
2. How can you characterize the appearance of the resulting broken line?
6.13. Ionization energy of an atom
If you give an electron in an atom additional energy (you will learn how to do this from a physics course), then the electron can go to another AO, that is, the atom will end up in excited state. This state is unstable, and the electron will almost immediately return to its original state, and excess energy will be released. But if the energy imparted to the electron is large enough, the electron can completely break away from the atom, while the atom ionized, that is, it turns into a positively charged ion ( cation). The energy needed to do this is called ionization energy of an atom(E and).
It is quite difficult to tear off an electron from a single atom and measure the energy required for this, therefore, it is practically determined and used molar ionization energy(E and m).
Molar ionization energy shows what is the smallest energy required to detach 1 mole of electrons from 1 mole of atoms (one electron from each atom). This value is usually measured in kilojoules per mole. The values of the molar ionization energy of the first electron for most elements are given in Appendix 6.
How does the ionization energy of an atom depend on the position of the element in the system of elements, that is, how does it change in the group and period?
In physical terms, the ionization energy is equal to the work that must be spent to overcome the force of attraction of an electron to an atom when moving an electron from an atom to an infinite distance from it.
where q is the charge of an electron, Q is the charge of the cation remaining after the removal of an electron, and r o is the orbital radius of the atom.
And q, and Q are constant values, and it can be concluded that, the work of detaching an electron BUT, and with it the ionization energy E and, are inversely proportional to the orbital radius of the atom.
After analyzing the values of the orbital radii of atoms of various elements and the corresponding values of the ionization energy given in Appendices 5 and 6, you can see that the dependence between these values is close to proportional, but somewhat different from it. The reason that our conclusion does not agree well with the experimental data is that we used a very rough model that does not take into account many significant factors. But even this rough model allowed us to draw the correct conclusion that with an increase in the orbital radius, the ionization energy of an atom decreases and, conversely, with a decrease in the radius, it increases.
Since the orbital radius of atoms decreases in a period with an increase in the serial number, the ionization energy increases. In a group, as the atomic number increases, the orbital radius of the atoms, as a rule, increases, and the ionization energy decreases. The highest molar ionization energy is in the smallest atoms, helium atoms (2372 kJ/mol), and of the atoms capable of forming chemical bonds, in fluorine atoms (1681 kJ/mol). The smallest is for the largest atoms, cesium atoms (376 kJ/mol). In a system of elements, the direction of increasing ionization energy can be schematically shown as follows: 
In chemistry, it is important that the ionization energy characterizes the propensity of an atom to donate "its" electrons: the greater the ionization energy, the less inclined the atom is to donate electrons, and vice versa.
Excited state, ionization, cation, ionization energy, molar ionization energy, change in ionization energy in a system of elements.
1. Using the data given in Appendix 6, determine how much energy you need to spend to tear off one electron from all sodium atoms with a total mass of 1 g.
2. Using the data given in Appendix 6, determine how many times more energy needs to be spent to detach one electron from all sodium atoms with a mass of 3 g than from all potassium atoms of the same mass. Why does this ratio differ from the ratio of the molar ionization energies of the same atoms?
3. According to the data given in Appendix 6, plot the dependence of the molar ionization energy on the serial number for elements with Z from 1 to 40. The dimensions of the graph are the same as in the task for the previous paragraph. See if this graph matches the choice of "periods" of the system of elements.
6.14. Electron affinity energy
.The second most important energy characteristic of an atom is electron affinity energy(E with).
In practice, as in the case of ionization energy, the corresponding molar quantity is usually used - molar electron affinity energy().
The molar electron affinity energy shows what is the energy released when one mole of electrons is added to one mole of neutral atoms (one electron to each atom). Like the molar ionization energy, this quantity is also measured in kilojoules per mole.
At first glance, it may seem that energy should not be released in this case, because an atom is a neutral particle, and there are no electrostatic forces of attraction between a neutral atom and a negatively charged electron. On the contrary, approaching the atom, the electron, it would seem, should be repelled by the same negatively charged electrons that form the electron shell. Actually this is not true. Remember if you have ever dealt with atomic chlorine. Of course not. After all, it exists only at very high temperatures. Even more stable molecular chlorine is practically not found in nature - if necessary, it has to be obtained using chemical reactions. And you have to deal with sodium chloride (common salt) all the time. After all, table salt is consumed by a person with food every day. And it is quite common in nature. But after all, table salt contains chloride ions, that is, chlorine atoms that have attached one "extra" electron each. One of the reasons for this prevalence of chloride ions is that chlorine atoms have a tendency to attach electrons, that is, when chloride ions are formed from chlorine atoms and electrons, energy is released.
One of the reasons for the release of energy is already known to you - it is associated with an increase in the symmetry of the electron shell of the chlorine atom during the transition to a singly charged anion. At the same time, as you remember, energy 3 p- sublevel decreases. There are other more complex reasons.
Due to the fact that several factors influence the value of the electron affinity energy, the nature of the change in this value in a system of elements is much more complex than the nature of the change in the ionization energy. You can verify this by analyzing the table given in Appendix 7. But since the value of this quantity is determined, first of all, by the same electrostatic interaction as the values of the ionization energy, then its change in the system of elements (at least in A- groups) in general terms is similar to a change in the ionization energy, that is, the energy of electron affinity in a group decreases, and in a period it increases. It is maximum at the atoms of fluorine (328 kJ/mol) and chlorine (349 kJ/mol). The nature of the change in the electron affinity energy in the system of elements resembles the nature of the change in the ionization energy, that is, the direction of the increase in the electron affinity energy can be schematically shown as follows:
2. On the same scale along the horizontal axis as in the previous tasks, plot the dependence of the molar energy of electron affinity on the serial number for atoms of elements with Z from 1 to 40 using app 7.
3. What is the physical meaning of negative electron affinity energies?
4. Why, of all the atoms of the elements of the 2nd period, only beryllium, nitrogen and neon have negative values of the molar energy of electron affinity?
6.15. The tendency of atoms to donate and gain electrons
You already know that the propensity of an atom to donate its own and accept foreign electrons depends on its energy characteristics (ionization energy and electron affinity energy). What atoms are more inclined to donate their electrons, and which ones are more inclined to accept strangers?
To answer this question, let us summarize in Table 15 everything that we know about the change in these inclinations in the system of elements.
Table 15
How to use the periodic table? For an uninitiated person, reading the periodic table is the same as looking at the ancient runes of elves for a dwarf. And the periodic table, by the way, if used correctly, can tell a lot about the world. In addition to serving you in the exam, it is also simply indispensable for solving a huge number of chemical and physical problems. But how to read it? Fortunately, today everyone can learn this art. In this article we will tell you how to understand the periodic table.
The periodic system of chemical elements (Mendeleev's table) is a classification of chemical elements that establishes the dependence of various properties of elements on the charge of the atomic nucleus.
History of the creation of the Table
Dmitri Ivanovich Mendeleev was not a simple chemist, if someone thinks so. He was a chemist, physicist, geologist, metrologist, ecologist, economist, oilman, aeronaut, instrument maker and teacher. During his life, the scientist managed to conduct a lot of fundamental research in various fields of knowledge. For example, it is widely believed that it was Mendeleev who calculated the ideal strength of vodka - 40 degrees. We do not know how Mendeleev treated vodka, but it is known for sure that his dissertation on the topic “Discourse on the combination of alcohol with water” had nothing to do with vodka and considered alcohol concentrations from 70 degrees. With all the merits of the scientist, the discovery of the periodic law of chemical elements - one of the fundamental laws of nature, brought him the widest fame.
There is a legend according to which the scientist dreamed of the periodic system, after which he only had to finalize the idea that had appeared. But, if everything were so simple .. This version of the creation of the periodic table, apparently, is nothing more than a legend. When asked how the table was opened, Dmitry Ivanovich himself answered: “ I’ve been thinking about it for maybe twenty years, and you think: I sat and suddenly ... it’s ready. ”
In the middle of the nineteenth century, attempts to streamline the known chemical elements (63 elements were known) were simultaneously undertaken by several scientists. For example, in 1862 Alexandre Emile Chancourtois placed the elements along a helix and noted the cyclical repetition chemical properties. Chemist and musician John Alexander Newlands proposed his version of the periodic table in 1866. An interesting fact is that in the arrangement of the elements the scientist tried to discover some mystical musical harmony. Among other attempts was the attempt of Mendeleev, which was crowned with success.
In 1869, the first scheme of the table was published, and the day of March 1, 1869 is considered the day of the discovery of the periodic law. The essence of Mendeleev's discovery was that the properties of elements with increasing atomic mass do not change monotonously, but periodically. The first version of the table contained only 63 elements, but Mendeleev made a number of very non-standard decisions. So, he guessed to leave a place in the table for yet undiscovered elements, and also changed the atomic masses of some elements. The fundamental correctness of the law derived by Mendeleev was confirmed very soon after the discovery of gallium, scandium and germanium, the existence of which was predicted by scientists.
Modern view of the periodic table
Below is the table itself.
Today, instead of atomic weight (atomic mass), the concept of atomic number (the number of protons in the nucleus) is used to order elements. The table contains 120 elements, which are arranged from left to right in ascending order of atomic number (number of protons)
The columns of the table are so-called groups, and the rows are periods. There are 18 groups and 8 periods in the table.
- The metallic properties of elements decrease when moving along the period from left to right, and increase in the opposite direction.
- The dimensions of atoms decrease as they move from left to right along the periods.
- When moving from top to bottom in the group, the reducing metallic properties increase.
- Oxidizing and non-metallic properties increase along the period from left to right. I.
What do we learn about the element from the table? For example, let's take the third element in the table - lithium, and consider it in detail.
First of all, we see the symbol of the element itself and its name under it. In the upper left corner is the atomic number of the element, in the order in which the element is located in the table. The atomic number, as already mentioned, is equal to the number of protons in the nucleus. The number of positive protons is usually equal to the number of negative electrons in an atom (with the exception of isotopes).
The atomic mass is indicated under the atomic number (in this version of the table). If we round the atomic mass to the nearest integer, we get the so-called mass number. The difference between the mass number and the atomic number gives the number of neutrons in the nucleus. So, the number of neutrons in a helium nucleus is two, and in lithium - four.
So our course "Mendeleev's Table for Dummies" has ended. In conclusion, we invite you to watch a thematic video, and we hope that the question of how to use the periodic table of Mendeleev has become clearer to you. We remind you that learning a new subject is always more effective not alone, but with the help of an experienced mentor. That is why, you should never forget about those who will gladly share their knowledge and experience with you.