Mendeleev called horizontal rows of elements, within which the properties of elements change sequentially periods(start with an alkali metal (Li, Na, K, Rb, Cs, Fr) and end with a noble gas (He, Ne, Ar, Kr, Xe, Rn)).
Exceptions: the first period, which begins with hydrogen, and the seventh period, which is incomplete.
The periods are divided into small And big. Small periods consist of one horizontal row. The first, second and third periods are small, they contain 2 elements (1st period) or 8 elements (2nd, 3rd periods). Large periods consist of two horizontal rows. The fourth, fifth and sixth periods are large, containing 18 elements (4th, 5th periods) or 32 elements (6th, 7th periods). Top rows long periods are called even, the bottom rows are odd.
In the sixth period, the lanthanides and in the seventh period, the actinides are located at the bottom of the periodic table.
In each period, from left to right, the metallic properties of the elements weaken, and the non-metallic properties increase.
In even rows of large periods there are only metals.
As a result, the table has 7 periods, 10 rows and 8 vertical columns, called groups – is a collection of elements that have the same highest valency in oxides and in other compounds. This valency is equal to the group number.
Exceptions:
In group VIII, only Ru and Os have the highest valency VIII.
Groups are vertical sequences of elements, they are numbered with Roman numerals from I to VIII and Russian letters A and B. Each group consists of two subgroups: main and secondary. Main subgroup– A, contains elements of small and large periods. Side subgroup - B, contains elements only of large periods. They include elements of periods starting from the fourth.
In the main subgroups, from top to bottom, metallic properties are strengthened, and non-metallic properties are weakened. All elements of secondary subgroups are metals.
Quantum numbers
The principal quantum number n determines full energy electron. Each number corresponds to an energy level. n=1,2,3,4…or K,L,M,N…
The orbital quantum number l determines the sublevels at the energy level. Quantum number l determines the shape of orbitals (n-1) 0,1,2…
The magnetic quantum number ml determines the number of orbitals at the sublevel. …-2,-1,0,+1,+2… Total number orbitals at the sublevel is 2l+1
The spin quantum number ms refers to two different orientations +1/2 -1/2 in each orbital there can only be two electrons with opposite spins.
Rule for filling energy levels and sublevels of elements of the periodic table
Klechkovsky's first rule: as the charge of the atomic nucleus increases, the filling of energy levels occurs from orbitals with a smaller value of the sum of the principal and orbital * quantum numbers (n+l) to orbitals with a larger value of this sum. Therefore, the 4s sublevel (n+l=4) should be filled earlier than the 3d (n+l=5).
The second Klechkovsky rule, according to which, for the same values of the sum (n+l), the orbitals are filled in order of increasing principal quantum number n. The 3d sublevel is filled in ten elements from Sc to Zn. These are atoms of d-elements. Then the formation of the 4p sublevel begins. The order of filling sublevels in accordance with Klechkovsky's rules can be written as a sequence: 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d 7p.
Features of the electronic structure of atoms of elements of the periodic table
Features of the electronic structure of atoms of elements in the main and secondary subgroups, families of lanthanides and actinides
Shielding and penetration effects
Due to shielding, the attraction of valence electrons to the nucleus is weakened. At the same time, the opposite role is played by the penetrating ability of valence electrons to the nucleus, which enhances the interaction with the nucleus. Overall result the attraction of valence electrons to the nucleus depends on the relative contribution to their interaction of the screening influence of electrons in the inner layers and the penetrating ability of valence electrons to the nucleus.
Periodic nature of the properties of elements associated with the structures of their electronic shells
Changes in the acid-base properties of oxides and hydroxides in periods and groups
The acidic properties of elemental oxides increase in periods from left to right and in groups from bottom to top.!
Oxidation states of elements
Oxidation state (oxidation number, formal charge) - an auxiliary conventional value for recording the processes of oxidation, reduction and redox reactions, a numerical value electric charge, assigned to an atom in a molecule under the assumption that the electron pairs performing the bond are completely biased towards more electronegative atoms.
Ideas about the degree of oxidation form the basis for the classification and nomenclature of inorganic compounds.
The oxidation number corresponds to the charge of an ion or the formal charge of an atom in a molecule or chemical formal unit, for example:
The oxidation number is indicated above the element symbol. Unlike indicating the charge of an atom, when indicating the oxidation state, the sign is given first, and then the numerical value, and not vice versa.
1) s-block in the periodic table of elements - an electron shell that includes the first two layers of s-electrons. This block includes alkali metals, alkaline earth metals, hydrogen and helium. These elements differ in that in the atomic state the high-energy electron is located in the s orbital. With the exception of hydrogen and helium, these electrons are very easily transferred and formed into positive ions when chemical reaction. The configuration of helium is chemically very stable, which is why helium does not have stable isotopes; sometimes, due to this property, it is combined with inert gases. The remaining elements that have this block, without exception, are strong reducing agents and therefore are not found in free form in nature. The element in metallic form can only be obtained by electrolysis of a salt dissolved in water. Davy Humphrey, in 1807 and 1808, became the first to detach acid salts from s-block metals, with the exception of lithium, beryllium, rubidium and cesium. Beryllium was first separated from salts independently by two scientists: F. Wooler and A. A. Bazi in 1828, while lithium was separated only in 1854 by R. Bunsen, who, after studying rubidium, separated it 9 years later. Cesium was not isolated in its pure form until 1881, after Carl Setterberg electrolyzed cesium cyanide. The hardness of elements having an s-block in compact form (under normal conditions) can vary from very low (all alkali metals - they can be cut with a knife) to quite high (beryllium). With the exception of beryllium and magnesium, the metals are very reactive and can be used in alloys with lead in small quantities (<2 %). Бериллий и магний, ввиду их высокой стоимости, могут быть ценными компонентами для деталей, где требуется твёрдость и лёгкость. Эти металлы являются чрезвычайно важными, поскольку позволяют сэкономить средства при добыче титана, циркония, тория и тантала из их минеральных форм; могут находить своё применение как восстановители в органической химии.
Danger and storage
All elements with an s-shell are hazardous substances. They are fire hazardous and require special fire extinguishing, with the exception of beryllium and magnesium. Must be stored in an inert atmosphere of argon or hydrocarbons. React violently with water, the reaction product is hydrogen, for example:
Excluding magnesium, which reacts slowly, and beryllium, which reacts only when its oxide film is removed with mercury. Lithium has similar properties to magnesium, since it is located, relative to the periodic table, next to magnesium.
The P block in the periodic table of elements is the electron shell of atoms whose highest energy valence electrons occupy the p orbital.
The p-block includes the last six groups, excluding helium (which is in the s-block). This block contains all nonmetals (excluding hydrogen and helium) and semimetals, as well as some metals.
The P-block contains elements that have various properties, both physical and mechanical. P-block non-metals are, as a rule, highly reactive substances with strong electronegativity, p-metals are moderately active metals, and their activity increases towards the bottom of the table of chemical elements
Properties of d- and f-elements. Give examples.
The D block in the periodic table of elements is the electron shell of atoms whose highest energy valence electrons occupy the d orbital.
This block is part of the periodic table; it includes elements from groups 3 to 12. The elements of this block fill the d-shell with d-electrons, which for the elements begins with s2d1 (third group) and ends with s2d10 (twelfth group). However, there are some irregularities in this sequence, for example, in chromium s1d5 (but not s2d4) the entire eleventh group has the configuration s1d10 (but not s2d9). The eleventh group has filled s and d electrons.
D-block elements are also known as transition metals or transition elements. However, the exact boundaries separating transition metals from other groups of chemical elements have not yet been drawn. Although some authors believe that the elements included in the d-block are transition elements in which the d-electrons are partially filled or in neutral atoms or ions where the oxidation state is zero. IUPAC currently accepts such studies as reliable, and reports that this applies only to 3-12 groups of chemical elements. Group 12 metals do not have clearly defined chemical and physical properties, this is explained by the incomplete filling of the d subshell, so they can also be considered post-transition metals. The historical use of the term "transition elements" and d-block was also revised.
In the s-block and p-block of the periodic table, similar properties are, as a rule, not observed across periods: the most important properties are enhanced vertically in the lower elements of these groups. It is noteworthy that the differences between the elements included in the d-block horizontally, through periods, become more pronounced.
Lutetium and lawrencium are in the d-block and are not considered transition metals, but the lanthanides and actinides, remarkably, are considered so by IUPAC. Although the twelfth group of chemical elements is located in the d-block, it is believed that the elements included in it are post-transition elements
The p-elements of the periodic table include elements with a valence p-sublevel. These elements are located in III, IV, V, VI, VII, VIII groups, main subgroups. During the period, the orbital radii of atoms decrease with increasing atomic number, but generally increase. In subgroups of elements, as the element number increases, the sizes of atoms generally increase and decrease. p-elements of group III Group III p-elements include gallium Ga, indium In and thallium Tl. By the nature of these elements, boron is a typical non-metal, the rest are metals. Within the subgroup there is a sharp transition from non-metals to metals. The properties and behavior of boron are similar, which is the result of the diagonal affinity of elements in the periodic table, according to which a shift in a period to the right causes an increase in non-metallic character, and down the group - a metallic character, therefore elements with similar properties are located diagonally next to each other, for example Li and Mg, Ber and Al, B and Si.
The electronic structure of the valence sublevels of atoms of group III p-elements in the ground state has the form ns 2 np 1 . In compounds, boron and trivalent, gallium and indium, in addition, can form compounds with +1, and for thallium the latter is quite characteristic.
p-elements of group VIII Group VIII p-elements include helium He, neon Ne, argon Ar, krypton Kr, xenon Xe and radon Rh, which form the main subgroup. The atoms of these elements have complete outer electronic layers, so the electronic configuration of the valence sublevels of their atoms in the ground state is 1s 2 (He) and ns 2 np 6 (other elements). Due to the very high stability of electronic configurations, they are generally characterized by high ionization energies and chemical inertness, which is why they are called noble (inert) gases. In a free state, they exist in the form of atoms (monatomic molecules). The atoms of helium (1s 2), neon (2s 2 2p 6) and argon (3s 2 3p 6) have a particularly stable electronic structure, so valence-type compounds are unknown for them.
Krypton (4s 2 4p 6), xenon (5s 2 5p 6) and radon (6s 2 6p 6) differ from the previous noble gases in their larger atomic sizes and, accordingly, lower ionization energies. They are capable of forming compounds that often have low stability.
Elements in Mendeleev's periodic table are divided into s-, p-, d-elements. This division is carried out on the basis of how many levels the electron shell of an element’s atom has and at what level the filling of the shell with electrons ends.
TO s-elements include elements IA-groups – alkali metals. Electronic formula of the valence shell of alkali metal atoms ns1. The stable oxidation state is +1. Elements IA-groups have similar properties due to the similar structure of the electron shell. As the radius in the Li-Fr group increases, the bond between the valence electron and the nucleus weakens and the ionization energy decreases. Atoms of alkaline elements easily give up their valence electron, which characterizes them as strong reducing agents.
The reducing properties increase with increasing serial number.
TO p-elements include 30 elements IIIA-VIIIA-groups periodic system; p-elements are located in the second and third minor periods, as well as in the fourth to sixth major periods. Elements IIIA-groups have one electron in the p orbital. IN IVA-VIIIA-groups the filling of the p-sublevel with up to 6 electrons is observed. General electronic formula of p-elements ns2np6. In periods with increasing nuclear charge, the atomic radii and ionic radii of p-elements decrease, ionization energy and electron affinity increase, electronegativity increases, the oxidative activity of compounds and the non-metallic properties of elements increase. In groups, the radii of atoms increase. From 2p elements to 6p elements, the ionization energy decreases. The metallic properties of the p-element in the group increase with increasing atomic number.
TO d-elements There are 32 elements of the periodic table IV–VII major periods. IN IIIB-group atoms have the first electron in the d-orbital, in subsequent B-groups the d-sublevel is filled with up to 10 electrons. General formula for the outer electron shell (n-1)dansb, where a=1?10, b=1?2. With an increase in the ordinal number, the properties of d-elements change slightly. The d-elements slowly increase in atomic radius, and they also have a variable valency associated with the incompleteness of the outer d-electron sublevel. In lower oxidation states, d-elements exhibit metallic properties; with an increase in the atomic number in groups B, they decrease. In solutions, d-elements with the highest oxidation state exhibit acidic and oxidizing properties, and vice versa at lower oxidation states. Elements with intermediate oxidation states exhibit amphoteric properties.
8. Covalent bond. Valence bond method
A chemical bond carried out by common electron pairs arising in the shells of bonded atoms having antiparallel spins is called atomic or covalent bond. The covalent bond is two-electron and two-center (holds the nuclei). It is formed by atoms of one type - covalent non-polar– a new electron pair, arising from two unpaired electrons, becomes common to two chlorine atoms; and atoms of different types, similar in chemical character - covalent polar. Elements with greater electronegativity (Cl) will withdraw shared electrons from elements with less electronegativity (H). Atoms with unpaired electrons having parallel spins repel each other - no chemical bond occurs. The method of forming a covalent bond is called exchange mechanism.
Properties of covalent bonds. Link length – internuclear distance. The shorter this distance, the stronger the chemical bond. Communication energy – the amount of energy required to break a bond. The bond multiplicity is directly proportional to the bond energy and inversely proportional to the bond length. Communication direction – a specific arrangement of electron clouds in a molecule. Saturability– the ability of an atom to form a certain number of covalent bonds. A chemical bond formed by overlapping electron clouds along an axis connecting the centers of atoms is called ?-connection. A bond formed by overlapping electron clouds perpendicular to the axis connecting the centers of atoms is called ?-connection. The spatial orientation of a covalent bond is characterized by the angles between the bonds. These angles are called bond angles. Hybridization – the process of restructuring electron clouds of unequal shape and energy, leading to the formation of hybrid clouds identical in the same parameters. Valence– number of chemical bonds (covalent ), through which an atom is connected to others. Electrons involved in the formation of chemical bonds are called valence. The number of bonds between atoms is equal to the number of its unpaired electrons participating in the formation of common electron pairs, therefore valence does not take into account polarity and has no sign. In compounds in which there is no covalent bond, there is oxidation state – the conventional charge of an atom, based on the assumption that it consists of positively or negatively charged ions. The concept of oxidation state applies to most inorganic compounds.