Periodicity in changes in the properties of chemical elements based on the electronic structure of their atoms

Therefore, the methodological technique for compiling electronic formulas of elements based on the periodic system is that we sequentially consider the electronic shell of each element along the way to a given one, identifying by its “coordinates” where its next electron went in the shell.

The first two elements of the first period, hydrogen H and helium He, belong to the s-family. Two of their electrons enter the s-sublevel of the first level. We write down: The first period ends here, the first energy level also. The next two elements of the second period in order - lithium Li and beryllium Be are in the main subgroups of groups I and II. These are also s-elements. Their next electrons will be located on the s sublevel of the 2nd level. We write down 6 elements of the 2nd period follow in a row: boron B, carbon C, nitrogen N, oxygen O, fluorine F and neon Ne. According to the location of these elements in the main subgroups of the III - Vl groups, their next electrons, among the six, will be located on the p-sublevel of the 2nd level. We write down: The inert element neon ends the second period, the second energy level is also completed. This is followed by two elements of the third period of the main subgroups of groups I and II: sodium Na and magnesium Mg. These are s-elements and their next electrons are located on the s-sublevel of the 3rd level. Then there are six elements of the 3rd period: aluminum Al, silicon Si, phosphorus P, sulfur S, chlorine C1, argon Ar. According to the location of these elements in the main subgroups of groups III - UI, their next electrons, among the six, will be located on the p-sublevel of the 3rd level - The inert element argon has completed the 3rd period, but the 3rd energy level has not yet been completed, as long as there are no electrons on its third possible d-sublevel.

This is followed by 2 elements of the 4th period of the main subgroups of groups I and II: potassium K and calcium Ca. These are s-elements again. Their next electrons will be at the s-sublevel, but already at the 4th level. It is energetically more favorable for these next electrons to start filling the 4th level, which is more distant from the nucleus, than to fill the 3d sublevel. We write down: The following ten elements of the 4th period from No. 21 scandium Sc to No. 30 zinc Zn are in secondary subgroups III - V - VI - VII - VIII - I - II groups. Since they are all d-elements, their next electrons are located on the d-sublevel before the outer level, i.e., the third from the nucleus. We write down:

The following six elements of the 4th period: gallium Ga, germanium Ge, arsenic As, selenium Se, bromine Br, krypton Kr - are in the main subgroups of groups III - VIIJ. Their next 6 electrons are located on the p-sublevel of the outer, i.e., 4th level: 3b elements were considered; the fourth period is completed by the inert element krypton; The 3rd energy level is also completed. However, at level 4, only two sublevels are completely filled: s and p (out of 4 possible).

This is followed by 2 elements of the 5th period of the main subgroups of groups I and II: No. 37 rubidium Rb and No. 38 strontium Sr. These are elements of the s-family, and their next electrons are located on the s-sublevel of the 5th level: The last 2 elements - No. 39 yttrium YU No. 40 zirconium Zr - are already in secondary subgroups, i.e. they belong to the d-family. Their next two electrons will go to the d-sublevel, before the outer one, i.e. 4th level Summing up all the records sequentially, we compose the electronic formula for the zirconium atom No. 40. The derived electronic formula for the zirconium atom can be slightly modified by arranging the sublevels in the order of numbering their levels:

The derived formula can, of course, be simplified into the distribution of electrons only among energy levels: Zr – 2|8| 18 |8 + 2| 2 (the arrow indicates the entry point of the next electron; valence electrons are underlined). The physical meaning of the category of subgroups lies not only in the difference in the place where the next electron enters the shell of the atom, but also in the levels at which the valence electrons are located. From a comparison of simplified electronic formulas, for example, chlorine (3rd period, main subgroup of group VII), zirconium (5th period, secondary subgroup of group IV) and uranium (7th period, lanthanide-actinide subgroup)

№17, С1-2|8|7

No. 40, Zr - 2|8|18|8+ 2| 2

No. 92, U - 2|8|18 | 32 |18 + 3|8 + 1|2

It can be seen that for elements of any main subgroup, only electrons of the outer level (s and p) can be valence. For elements of side subgroups, valence electrons can be the electrons of the outer and partially pre-outer level (s and d). In lanthanides and especially actinides, valence electrons can be located at three levels: external, pre-external and pre-external. Typically, the total number of valence electrons is equal to the group number.

The main subgroup of group V of the periodic table includes nitrogen, phosphorus, arsenic, antimony and bismuth.

These elements, having five electrons in the outer layer of the atom, are generally characterized as nonmetals. However, their ability to add electrons is much less pronounced than that of the corresponding elements of groups VI and VII. Due to the presence of five outer electrons, the highest positive oxidation of the elements of this subgroup is -5, and the negative - 3. Due to the relatively lower electronegativity, the bond of the elements in question with hydrogen is less polar than the bond with hydrogen of elements of groups VI and VII. Therefore, hydrogen compounds of these elements do not eliminate hydrogen ions H in an aqueous solution, and thus do not have acidic properties.

The physical and chemical properties of the elements of the nitrogen subgroup change with increasing atomic number in the same sequence that was observed in the previously considered groups. But since non-metallic properties are less pronounced than in oxygen and, especially, fluorine, these properties weaken when moving to the next elements entails the appearance and increase of metallic properties. The latter are already noticeable in arsenic, antimony has both properties approximately equally, and in bismuth the metallic properties predominate over the nonmetallic ones.

DESCRIPTION OF ELEMENTS.

NITROGEN(from Greek ázōos - lifeless, lat. Nitrogenium), N, chemical element of group V of the Mendeleev periodic system, atomic number 7, atomic mass 14.0067; colorless gas, odorless and tasteless.

Historical reference. Nitrogen compounds - saltpeter, nitric acid, ammonia - were known long before nitrogen was obtained in a free state. In 1772, D. Rutherford, burning phosphorus and other substances in a glass bell, showed that the gas remaining after combustion, which he called “suffocating air,” does not support respiration and combustion. In 1787, A. Lavoisier established that the “vital” and “asphyxiating” gases that make up the air are simple substances, and proposed the name “nitrogen”. In 1784, G. Cavendish showed that nitrogen is part of the saltpeter; This is where the Latin name nitrogen comes from (from the Late Latin nitrum - saltpeter and the Greek gennao - I give birth, I produce), proposed in 1790 by J. A. Chaptal. By the beginning of the 19th century. The chemical inertness of nitrogen in the free state and its exclusive role in compounds with other elements as bound nitrogen were clarified. Since then, the “binding” of air nitrogen has become one of the most important technical problems of chemistry.

Prevalence in nature. Nitrogen is one of the most common elements on Earth, and the bulk of it (about 4´1015 tons) is concentrated in a free state in the atmosphere. In the air, free nitrogen (in the form of N2 molecules) is 78.09% by volume (or 75.6% by mass), not counting its minor impurities in the form of ammonia and oxides. The average nitrogen content in the lithosphere is 1.9´10-3% by mass.

Natural nitrogen compounds. - ammonium chloride NH4Cl and various nitrates (see Saltpeter.) Large accumulations of saltpeter are characteristic of a dry desert climate (Chile, Central Asia). For a long time, nitrate was the main supplier of nitrogen for industry (now the industrial synthesis of ammonia from air nitrogen and hydrogen is of primary importance for nitrogen fixation). Small amounts of fixed nitrogen are found in coal (1-2.5%) and oil (0.02-1.5%), as well as in the waters of rivers, seas and oceans. Nitrogen accumulates in soils (0.1%) and in living organisms (0.3%).

Although the name "nitrogen" means "non-life-sustaining", it is actually an essential element for life. Animal and human protein contains 16 - 17% nitrogen. In the organisms of carnivorous animals, protein is formed due to the consumed protein substances present in the organisms of herbivorous animals and in plants. Plants synthesize protein by assimilating nitrogenous substances contained in the soil, mainly inorganic. Significant amounts of nitrogen enter the soil thanks to nitrogen-fixing microorganisms that are capable of converting free nitrogen from the air into nitrogen compounds.

In nature, there is a nitrogen cycle, in which the main role is played by microorganisms - nitrophying, denitrophying, nitrogen-fixing, etc. However, as a result of the extraction of huge amounts of fixed nitrogen from the soil by plants (especially with intensive farming), the soils become depleted of nitrogen. Nitrogen deficiency is typical for agriculture in almost all countries; nitrogen deficiency is also observed in animal husbandry (“protein starvation”). On soils poor in available nitrogen, plants develop poorly. Nitrogen fertilizers and protein feeding of animals are the most important means of boosting agriculture. Human economic activities disrupt the nitrogen cycle. Thus, burning fuel enriches the atmosphere with nitrogen, and factories producing fertilizers bind nitrogen from the air. Transporting fertilizers and agricultural products redistributes nitrogen to the surface of the earth.

Nitrogen is the fourth most abundant element in the solar system (after hydrogen, helium and oxygen).

Isotopes, atom, molecule. Natural nitrogen consists of two stable isotopes: 14N (99.635%) and 15N (0.365%). The 15N isotope is used in chemical and biochemical research as a labeled atom. Of the artificial radioactive isotopes of nitrogen, 13N has the longest half-life (T1/2 - 10.08 min), the rest are very short-lived. In the upper layers of the atmosphere, under the influence of neutrons from cosmic radiation, 14N turns into the radioactive carbon isotope 14C. This process is also used in nuclear reactions to produce 14C. The outer electron shell of the nitrogen atom. consists of 5 electrons (one lone pair and three unpaired - configuration 2s22p3). Most often nitrogen. in compounds it is 3-covalent due to unpaired electrons (as in ammonia NH3). The presence of a lone pair of electrons can lead to the formation of another covalent bond, and nitrogen becomes 4-covalent (as in the ammonium ion NH4+). Nitrogen oxidation states vary from +5 (in N205) to -3 (in NH3). Under normal conditions, in a free state, nitrogen forms an N2 molecule, where the N atoms are connected by three covalent bonds. The nitrogen molecule is very stable: its dissociation energy into atoms is 942.9 kJ/mol (225.2 kcal/mol), therefore, even at a temperature of about 3300°C, the degree of dissociation is nitrogen. is only about 0.1%.

Physical and chemical properties. Nitrogen is slightly lighter than air; density 1.2506 kg/m3 (at 0°C and 101325 n/m2 or 760 mm Hg), melting point -209.86°C, boiling point -195.8?C. A. liquefies with difficulty: its critical temperature is quite low (-147.1 °C), and its critical pressure is high 3.39 Mn/m2 (34.6 kgf/cm2); The density of liquid nitrogen is 808 kg (m3. Nitrogen is less soluble in water than oxygen: at 0°C, 23.3 g of nitrogen dissolves in 1 m3 of H2O. Nitrogen is soluble in some hydrocarbons better than in water.

Nitrogen interacts only with such active metals as lithium, calcium, magnesium when heated to relatively low temperatures. Nitrogen reacts with most other elements at high temperatures and in the presence of catalysts. The compounds of nitrogen with oxygen N2O, NO, N2O3, NO2 and N2O5 have been well studied. From these, during the direct interaction of elements (4000°C), NO oxide is formed, which, upon cooling, is easily oxidized further to NO2 dioxide. In the air, nitrogen oxides are formed during atmospheric discharges. They can also be obtained by exposing a mixture of nitrogen and oxygen to ionizing radiation. When nitrous anhydrides N2O3 and nitric anhydrides N2O5 are dissolved in water, nitrous acid HNO2 and nitric acid HNO3 are obtained, respectively, forming salts - nitrites and nitrates. Nitrogen combines with hydrogen only at high temperatures and in the presence of catalysts, and ammonia NH3 is formed. In addition to ammonia, numerous other compounds of nitrogen with hydrogen are known, for example, hydrazine H2N-NH2, diimide HN-NH, hydronitric acid HN3(H-N-NºN), octazone N8H14, etc.; Most nitrogen and hydrogen compounds are isolated only in the form of organic derivatives. Nitrogen does not directly interact with halogens, therefore all nitrogen halides are obtained only indirectly, for example, nitrogen fluoride NF3- when fluorine reacts with ammonia. As a rule, nitrogen halides are low-stable compounds (with the exception of NF3); Nitrogen oxyhalides are more stable - NOF, NOCI, NOBr, N02F and NO2CI. Nitrogen does not combine directly with sulfur either; nitrogenous sulfur N4S4 is obtained as a result of the reaction of liquid sulfur with ammonia. When hot coke reacts with nitrogen, cyanogen (CN) is formed.;. By heating nitrogen with acetylene C2H2 to 1500°C, hydrogen cyanide HCN can be obtained. The interaction of nitrogen with metals at high temperatures leads to the formation of nitrides (for example, Mg3N2).

16. Which of the gases taken with the same mass occupies the largest volume under the same conditions:

17. Determine the molar mass equivalent (g/mol) of sulfur in sulfur oxide (VI):

18. What is the mass fraction (%) of the metal in the oxide if the molar mass of the trivalent metal equivalent is 15 g/mol:

19. What is the relative molecular mass of a gas if this gas is 2.2 times heavier than air:

20. Which of the following equations is called the Mendeleev–Clapeyron equation:

3) PV = RT

21. List 3 gases that have the same density as any other gas:

1) CH 4, SO 2, Cl 2

2) C 2 H 4, CH 4, F 2

3) CO, Cl 2, H 2

4) CO, C 2 H 4, N 2

5)N 2, CH 4, H 2

22. How many moles of oxygen are formed from 3 moles of potassium chlorate during its complete thermal decomposition:

23. What amount (mol) of FeS 2 will be required to obtain 64 g of SO 2 according to the equation:

4 FeS 2 + 11O 2 = 2Fe 2 O 3 + 8SO 2;

24. What mass (g) of calcium carbonate will be consumed to produce 44.8 liters of carbon dioxide, measured at ambient conditions:

1) 200,0;

25. The equivalent of aluminum is:

1) aluminum atom;

2) 1/2 part aluminum atom;

3) 1/3 part aluminum atom;

4) two aluminum atoms;

5) 1 mole of aluminum atoms.

26. The law of constancy of the composition of substances is valid for substances:

1) with a molecular structure;

2) with a non-molecular structure;

3) with an ionic crystal lattice;

4) with an atomic crystal lattice;

5) for oxides and salts.

27. The equivalent of magnesium is:

1) magnesium atom;

2) 1/2 part of a magnesium atom;

3) 1/3 part of a magnesium atom;

4) two magnesium atoms;

5) 1 mole of magnesium atoms.

28. To neutralize 2.45 g of acid, 2.80 g of potassium hydroxide is consumed. Define

molar mass of acid equivalent:

1) 98 g/mol;

2) 36.5 g/mol;

3) 63 g/mol;

4) 40 g/mol;

G/mol.

Classification and nomenclature of inorganic compounds

1) Na 2 O; CaO; CO2

2) SO 3; CuO; CrO3

3)Mn 2 O 7; CuO; CrO3

4) SO 3; CO2; P2O5

5) Na 2 O; H2O; CO2

30. Only acid oxides series:

1) CO 2; SiO2; MnO; CrO3

2) V 2 O 5; CrO3; TeO3; Mn2O7

3) CuO; SO2; NiO; MnO

4) CaO; P 2 O 3 ; Mn 2 O 7; Cr2O3

5) Na 2 O; H2O; CuO; Mn2O7

31. Cannot be used to neutralize sulfuric acid:

1) sodium bicarbonate;

2) magnesium oxide;

3) hydroxomagnesium chloride;

4) sodium hydrogen sulfate;

5) sodium oxide

32. To neutralize sulfuric acid you can use:

2) Mg(OH) 2

33. Using a glass tube, carbon dioxide is exhaled into solutions. The change will be in the solution:

3) Ca(OH) 2;

34. By dissolving the corresponding oxide in water you can obtain:

35. Under certain conditions, salt is formed in the case of:

1) N 2 O 5 +SO 3;

4) H 2 SO 4 +NH 3;

36. Can form acid salts:

1) H 3 PO 4;

37. Can form basic salts:

2) Ba(OH)2;

38. Mass of limestone required to produce 112 kg of quicklime:

39. Reacts with water:

2) CaO;

40. Soluble in water:

3) Ba(OH)2;

41. To obtain potassium phosphate, potassium hydrogen phosphate must be affected by:

42. Acid oxide:

3) Mn 2 O 7;

43. Will interact directly in an aqueous solution:

2) Cu(OH) 2 and ZnO;

3) AI 2 O 3 and HCI;

4) Rb 2 O and NaOH;

5) CaO and K 2 O.

44. All salts are acidic in the group:

1) KCI, CuOHCI, NaHSO 4;

2) KAI(SO 4) 2, Na, Ca(HCO 3) 2;

3) CuS, NaHSO 3, Cu(HS) 2;

4) NaHCO 3, Na 2 HPO 4, NaH 2 PO 4;

5) AIOHCI 2, NaHCO 3, NaCN.

45. Does not form acid salts:

4) HPO 3;

46. ​​The title is spelled incorrectly:

1) ferrous sulfate;

2) potassium sulfate;

3) iron (II) hydrochloride;

4) copper (I) chloride;

5) ammonium sulfate.

47. When water is separated from a monobasic acid weighing 16.0 g, formed by an element in the oxidation state +5, an oxide weighing 14.56 g is obtained. The acid was taken:

1) nitrogen;

2) metavanadium;

3) orthophosphoric;

4) arsenic;

5) chloric.

48. When calcining metal (III) weighing 10.8 g in air, a metal oxide weighing 20.4 g was obtained. For calcination the following was taken:

2) aluminum AI;

3) iron Fe;

4) scandium Sc;

5) sodium Na.

49. Sign characterizing hydrochloric acid:

1) dibasic;

2) weak;

3) volatile;

4) oxygen-containing;

5) acid – oxidizing agent.

50. Dibasic acid:

1) nitrogen;

2) salt;

3) vinegar;

4) cyanide;

Selenium.

51. Monoprotic acid:

1) selenium;

2) phosphorous;

3) tellurium;

4) boric;

5) prussic

52. Two types of acid salts are formed:

1) sulfuric acid;

2) orthophosphoric acid;

3) metaphosphoric acid;

4) selenic acid;

5) sulfurous acid.

53. Does not form acid salts:

1) sulfuric acid;

2) orthophosphoric acid;

3) metaphosphoric acid;

4) selenic acid;

5) sulfurous acid.

54. Specify the cationic complex:

1) Na 3;

3) K 3;

4) CI 3;

5) K 2.

55. Complex non-electrolyte:

1) Na 3;

2) ;

3) K 3;

4) CI 3;

5) K 2.

56. Anion complex:

1) potassium hexacyanoferrate(III);

2) tetrachlorodiammineplatinum (IV);

3) diammine silver chloride;

57. Complex non-electrolyte:

1) potassium hexacyanoferrate (III);

2) tetrachlorodiammineplatinum (IV);

3) diammine silver chloride;

4) tetraammine copper (II) sulfate;

5) hexaaquachrome (III) chloride.

58. Formula of hexaaquachrome (III) chloride:

1) Na 3;

2) CI

3) CI 2;

4) CI 3;

5)K 2 Cr 2 O 7 .

59. Formula of hexaaquachrome (II) chloride:

1) Na 3;

2) CI

3) CI 2; 3bl

4) CI 3;

5)K 2 Cr 2 O 7 .

60. Yellow blood salt refers to:

1) To aqua complexes;

2) Hydrates;

3) To acidocomplexes;

4) To ammonia;

5) To chelates.

61. Copper sulfate refers to:

1) To aqua complexes;

2) Hydrates;

3) To acidocomplexes;

4) To ammonia;

5) To chelates.

62. To obtain CaCO 3, the following should be added to a solution of Ca(HCO 3) 2:

1) Ca(OH) 2;

“The structure of matter and the periodic law of D.I. Mendeleev"

63. In the nucleus of the most common lead isotope 207 Pb neutrons:

2) 125

64. Maximum number of electrons at level n = 3:

65. At an energy level with n = 4 sublevels:

66. Number of energy levels in a tungsten atom:

67. In the nucleus of the osmium atom there are protons:

68. The nucleus of a krypton atom contains:

P and 44n

69. Number of electrons in chromium ion:

70. An ion containing 18 electrons and 16 protons has a nuclear charge:

71. The maximum number of electrons that can occupy a 3s orbital:

72. The atom has the electronic configuration 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1:

73. The designations of orbitals are incorrect:

3) 1p, 2d

74. The particle has the same electronic configuration as the argon atom:

1) Ca 2+

75. Electron affinity is called:

1) the energy required to remove an electron from an unexcited atom;

2) the ability of an atom of a given element to attract electron density;

3) transition of an electron to a higher energy level;

4) release of energy when an electron attaches to an atom or ion;

5) chemical bond energy.

76. As a result of a nuclear reaction isotope is formed:

77. In a hydrogen atom, absorption of a photon with minimal energy requires the transition of an electron:

78. The particle-wave nature of the electron is characterized by the equation:

79. For the valence electron of a potassium atom, the possible values ​​of quantum numbers (n, l, m l , m s):

1) 4, 1, -1, - :

2) 4, 1, +1, + : 3bm

3) 4, 0, 0, + :

4) 5, 0, +1, + :

80. Charge of the nucleus of an atom whose configuration of valence electrons in the ground state is ...4d 2 5s 2:

81. The principal quantum number n determines:

1) the shape of the electron cloud;

2) electron energy;

82. Orbital quantum number l determines:

1) the shape of the electron cloud;

2) electron energy;

3) orientation of the electron cloud in space;

4) rotation of the electron around its own axis;

5) hybridization of the electron cloud.

83. Magnetic quantum number m determines:

1) the shape of the electron cloud;

2) electron energy;

3) orientation of the electron cloud in space;

4) rotation of the electron around its own axis;

5) hybridization of the electron cloud.

84. The spin quantum number m s determines:

1) the shape of the electron cloud;

2) electron energy;

3) orientation of the electron cloud in space;

4) rotation of the electron around its own axis;

5) hybridization of the electron cloud.

85. During - decay, the nucleus of an atom of a radioactive element emits:

1) electron;

2) positron;

4) two protons;

5) two neutrons.

86. During - - decay, the nucleus of an atom of a radioactive element emits:

1) electron;

2) positron;

3) two protons and two neutrons combined into the nucleus of a helium atom;

4) two protons;

5) two neutrons.

87. During + - decay, the nucleus of an atom of a radioactive element emits:

1) electron;

2) positron;

3) two protons and two neutrons combined into the nucleus of a helium atom;

4) two protons;

5) two neutrons.

88. The atomic orbital has the smallest value of the sum (n + l):

89. The atomic orbital has the largest value of the sum (n + l)

90. The nitrogen atom will be more stable if at the 2p sublevel three electrons are distributed, one in each orbital. This matches the content:

2) Pauli principle;

3) Hund's Rules;

4) 1st Klechkovsky rule;

5) 2nd Klechkovsky rule.

91. The twenty-first electron of the scandium atom is located on the 3d sublevel, and not on the 4p sublevel. This matches the content:

1) The principle of least energy;

2) Pauli principle;

3) Hund's Rules;

4) 1st Klechkovsky rule;

5) 2nd Klechkovsky rule.

92. The nineteenth electron of the potassium atom is located on the 4s sublevel, and not on the 3d sublevel. This matches the content:

1) The principle of least energy;

2) Pauli principle;

3) Hund's Rules;

4) 1st Klechkovsky rule;

5) 2nd Klechkovsky rule.

93. The only electron of the hydrogen atom in the ground state is located at the first energy level. This matches the content:

1) The principle of least energy;

2) Pauli principle;

3) Hund's Rules;

4) 1st Klechkovsky rule;

5) 2nd Klechkovsky rule.

94. The maximum number of electrons at the second energy level of atoms of elements

equals 8. This corresponds to the content:

1) The principle of least energy;

2) Pauli principle;

3) Hund's Rules;

4) 1st Klechkovsky rule;

5) 2nd Klechkovsky rule.

95. One of the mechanisms for the formation of a covalent bond:

1) radical;

2) exchange;

3) molecular;

4) ionic;

5) chain.

96. An example of a non-polar molecule having a polar covalent bond would be:

4) CCl 4

97. Non-polar molecule:

98. In the series of molecules LiF - BeF 2 - BF 3 - CF 4 - NF 3 - OF 2 - F 2:

1) the nature of the connection does not change;

2) the ionic nature of the bond is enhanced;

3) the covalent nature of the bond weakens;

4) the covalent nature of the bond is enhanced;

5) there is no correct answer.

99. A covalent bond is formed in a molecule by a donor-acceptor mechanism:

2) CCl 4;
3) NH 4 C1;
4) NH 3;

100. In a nitrogen molecule the following are formed:

1) only - connections;

2) only - connections;

3) both - and - connections;

4) single bond;

5) double bond.

101. The methane molecule has the structure:

1) flat;

2) tetrahedral;

3) pyramidal;

4) square;

102. The formation of an ionic lattice is characteristic of:

1) cesium iodide;

2) graphite;

3) naphthalene;

4) diamond;

103. Which of the following substances is characterized by the formation of an atomic lattice:

1) ammonium nitrate;

2) diamond;

4) sodium chloride;

5) sodium.

104. Chemical elements are arranged in order of increasing electronegativity in

1) Si, P, Se, Br, Cl, O;

2) Si, P, Br, Se, Cl, O;

3) P, Si, Br, Se, Cl, O;

4) Br, P, Cl, Si, Se;

5) Si, P, Se, Cl, O, Br

105. The valence orbitals of the beryllium atom in the beryllium hydride molecule ... are hybridized

106. The beryllium hydride molecule has the structure:

1) square

Flat

3) tetrahedral

5) spherical.

107. The valence orbitals of the boron atom in the BF 3 molecule are hybridized as follows:

108. Which molecule is the strongest?

109. Which of the following molecules has the largest dipole?

110. What spatial configuration does the molecule have during sp 2 hybridization of AO:

1) linear

2) tetrahedron

3) flat square

Flat trigonal

111. The molecule has an octahedral structure if the following hybridization occurs

3) d 2 sp 3

112. The modern theory of atomic structure is based on the concepts of:

1) classical mechanics;

2) quantum mechanics;

3) Bohr's theory;

4) electrodynamics;

5) chemical kinetics.

113. Of the following, the characteristics of atoms of elements change periodically:

1) charge of the atomic nucleus

2) relative atomic mass;

3) the number of energy levels in an atom;

4) number of electrons at the outer energy level;

5) total number of electrons.

114. Within a period, an increase in the serial number of an element is usually accompanied by:

1) a decrease in the atomic radius and an increase in the electronegativity of the atom;

2) an increase in the atomic radius and a decrease in the electronegativity of the atom;

3) a decrease in atomic radius and a decrease in the electronegativity of the atom

4) an increase in atomic radius and an increase in the electronegativity of the atom

5) a decrease in electronegativity.

115. The atom of which element gives up one electron most easily:

1) sodium, serial number 11;

2) magnesium, serial number 12;

3) aluminum, serial number 13;

4) silicon, serial number 14;

5) sulfur, serial number 16.

116. Atoms of elements of group IA of the periodic system of elements have the same number:

1) electrons at the outer electronic level;

2) neutrons;

3) all electrons;

4) electronic shells;

5) protons.

117. Which of the following elements is named after the country:

118. Which series includes only transitional elements:

1) elements 11, 14, 22, 42;

2) elements 13, 33, 54, 83;

3) elements 24, 39, 74, 80;

4) elements 19, 32, 51, 101;

5) elements 19, 20, 21, 22.

119. The atom of which of the VA group elements has the maximum radius:

2) phosphorus;

3) arsenic;

4) bismuth;

5) antimony.

120. Which series of elements is presented in order of increasing atomic radius:

1) O, S, Se, Te;

3) Na, Mg, AI, Si;

4) J, Br, CI, F;

5) Sc, Te, V, Cr.

121. Metallic character of the properties of elements in the series Mg – Ca – Sr – Ba

1) decreases;

2) increases;

3) does not change;

4) decreases and then increases;

5) increases and then decreases.

122. Basic properties of hydroxides of elements of the JA group as the atomic number increases

1) decrease,

2) increase,

3) remain unchanged,

4) decrease and then increase,

5) increase and then decrease.

123. Simple substances of which elements have the greatest similarity of physical and chemical properties:

3) F, CI;

124. The existence of which of the following elements was predicted by D.I. Mendeleev:

3) Sc, Ga, Ge;

125. What distinguishes large periods from small ones:

1) the presence of alkali metals;

2) absence of inert gases;

3) the presence of d- and f-elements;

4) the presence of non-metals;

5) the presence of elements with metallic properties.

126.How to determine the period in which a given element is located using the electronic formula of an element:

1) by the value of the principal quantum number of the external energy level;

2) by the number of valence electrons;

3) by the number of electrons in the external energy level;

4) by the number of sublevels in the external energy level;

5) by the value of the sublevel where the last valence electron is located.

127. Which element has the lowest ionization potential:

128. A chemical element of the third period forms a higher oxide of composition E 2 O 3. How are electrons distributed in an atom of a given element?

1) 1s 2 2s 2 2p 1

2) 1s 2 2s 2 2p 6 3s 1

3) 1s 2 2s 2 2p 6 3s 2 3p 1

4) 1s 2 2s 2 2p 6 3s 2 3p 6

5) 1s 2 2s 2 2p 3

129.Which chemical element forms the base with the most pronounced properties

1) calcium

3) aluminum

Potassium

5) beryllium

130. A chemical element has the following distribution of electrons across the electron layers in the atom 2.8.6. What position does it occupy in the periodic table of chemical elements D.I. Mendeleev:

1) 6th period 6th group

Period 6 group

3) 2nd period 6th group

4) 3rd period 2nd group

5) 2nd period 8th group

131. The quantum numbers of the last electron in an element’s atom are n = 5, l = 1, m = -1, m s = - . Where is this element located on the periodic table?

1) 5th period, first group

2) 5th period, main subgroup of 4th group

3) 4th period, sixth group

period, sixth group main subgroup

5) 5th period, sixth group - secondary subgroup.

132. Formula of the highest oxide of the chemical element EO 2. Which group of the main subgroup of the periodic system of chemical elements belongs to D.I. Does this element belong to Mendeleev?

Fourth

5) sixth.

133. From the given list of elements - Li, Na, Ag, Au, Ca, Ba - alkali metals include:

1) all metals;

2) Li, Na;

3) Li, Na, Ag, Au;

134. In the series from Li to Fr:

1) metallic properties are enhanced;

2) metallic properties decrease;

3) the atomic radius decreases;

4) the connection of valence electrons with the nucleus is enhanced;

5) activity towards water decreases

135. The sequence of elements does not apply to metals:

3) B, As, Te;

136. With increasing atomic number of the element, the acidic properties of the oxides N 2 O 3 - P 2 O 3 - As 2 O 3

Sb 2 O 3 - Bi 2 O 3

1) intensify;

2) weaken;

3) remain unchanged;

4) strengthen, then weaken;

5) weaken, then strengthen.

137. The ammonia molecule has the form:

1) curved;

2) linear;

3) planar;

4) pyramidal;

138. In the series C-Si-Ge-Sn-Pb, non-metallic characteristics of the elements:

1) increase;

2) weaken;

3) do not change;

4) increase and then decrease;

5) weaken and then increase.

139. The valence orbitals of the carbon atom in the CH4 methane molecule can be described based on

ideas about hybridization type (sp; sp 2; sp 3; d 2 sp 3; dsp 2).

In this case, the methane molecule has the form:

1) linear;

2) flat;

3) tetrahedral;

5) square.

140. The valence orbitals of the silicon atom in the SiH 4 silane molecule can be described based on the concept of hybridization of the type (sp; sp 2 ; sp 3 ; d 2 sp 3 ; dsp 2).

Therefore, the silane molecule has the form:

1) linear;

2) flat;

3) tetrahedral;

5) square.

141.What is the maximum number of covalent bonds that a nitrogen atom can form:

142. The nitrogen atom of an ammonia molecule with a hydrogen ion forms:

1) ionic bond;

2) covalent bond by exchange mechanism;

3) non-polar covalent bond;

4) covalent bond through the donor-acceptor mechanism;

5) hydrogen bond.

143. Which statement is false:

4) Ionic bonding is saturable;

144. Which statement is false:

1) The covalent bond is saturable;

2) The covalent bond has directionality;

3) The ionic bond is unsaturable;

4) The ionic bond is directional;

5) The ionic bond is non-directional.

“Regularities of chemical processes and their energy”

145. What changes in temperature T and pressure P contribute to the formation of CO according to the reaction C(solid) + CO 2 (g) 2CO (g) -119.8 kJ:

1) increase in T and increase in P;

2) increase in T and decrease in P;

3) decrease in T and increase in P;

4) decrease in T and decrease in P;

5) increase in R.

146. How many times will the rate of a chemical reaction increase when the temperature increases by 30 0, if the temperature coefficient of the rate is 2?

147. How many degrees must the temperature be lowered so that the reaction rate decreases by 27 times, if the temperature coefficient of the rate is 3?

148. How many times will the reaction rate X+ 2Y = Z increase with increasing concentration

Y 3 times?

149. How many times will the rate of the forward reaction increase compared to the rate of the reverse reaction in the 2NO + O 2 2NO 2 system when the pressure doubles?

150. Specify the correct expression for the speed for the system: 2Cr+3Cl 2 = 2CrCl 3

5) v= k[A][C].

154. A catalyst speeds up a chemical reaction due to:

1) decrease in activation energy;

2) increasing the activation energy;

3) reducing the heat of reaction;

4) increasing concentration;

5) all answers are incorrect.

155. The equilibrium of the reaction Fe 3 O 4 +4CO «3Fe +4CO 2 -43.7 kJ shifts to the left:

1) when the temperature drops;

2) with increasing temperature;

3) with increasing pressure;

4) with increasing concentration of starting substances;

5) when adding a catalyst.

156. How many times will the rate of a chemical reaction increase when the temperature increases by 30 0, if the temperature coefficient of the rate is 3?

157. How many degrees must the temperature be increased for the reaction rate to increase 27 times, if the temperature coefficient of the rate is 3?

158. How many times does the rate of reaction X+2Y=Z increase when the concentration of X increases by 3 times?

159. How many times will the rate of the forward reaction increase compared to the rate of the reverse reaction in the 2CO+O 2 2CO 2 system when the pressure doubles?

160. How will the rate of the gas reaction 2NO 2 =N 2 O 4 increase with an increase in the concentration of NO 2 by 5 times?

161. How many times will the rate of the gas reaction 2NO+O 2 =2NO 2 decrease when the mixture of reacting gases is diluted 3 times?

162. How many degrees must the temperature be lowered so that the reaction rate decreases by 81 times at a temperature coefficient of 3?

163. How many times will the rate of reaction 2NO+O 2 =2NO 2 increase when the pressure in the system increases by 4 times?

164. How many times will the rate of the forward reaction increase compared to the rate of the reverse reaction in the 2NO+O 2 2NO 2 system when the pressure in the system increases by 5 times?

165. How will the rate of the reaction 2SO 2.g + O 2.g 2SO 3.g change with increasing concentration

1) will increase 3 times;

2) will increase 9 times;

3) will decrease by 3 times;

4) will decrease by 9 times;

5) will not change.

166. How will the rate of the reaction 2O 3.g 3O 2.g change when the pressure doubles?

1) will decrease by 2 times;

2) will decrease by 8 times;

3) will increase 4 times;

4) will decrease by 4 times;

5) will increase by 2 times.

167. How will the rate of the reaction 2NO g + O 2.g 2NO 2.g change with a simultaneous decrease

concentration of NO and O 2 2 times?

1) will increase by 2 times;

2) will decrease by 2 times;

3) will increase by 2 4 times;

4) will decrease by 2-4 times;

Will decrease by 8 times.

168. How will the rate of the direct reaction H 2 O, g H 2, g + O 2, g change if the pressure in the system increases 4 times?

1) will increase by 2 times;

2) will decrease by 2 times;

3) will not change;

4) will increase 4 times;

5) will decrease by 4 times.

169. The law of mass action was discovered:

1) M.V. Lomonosov

2) G.I. Hessom

3) J.W. Gibbs

K. Guldberg and P. Waage

5) Van't - Hoff

170. Which of the following systems is homogeneous

Sodium chloride solution

2) ice water

3) saturated solution with sediment

4) coal and sulfur in the air atmosphere

5) a mixture of gasoline and water

171. The value of the rate constant of a chemical reaction does not depend

1) from the nature of the reacting substances

2) on temperature

3) from the presence of catalysts

From the concentration of substances

5) from any factors

172. Activation energy is

1) the energy required to remove an electron from an atom

2) the excess energy that molecules must have per 1 mole in order for their collision to lead to the formation of a new substance

3) ionization potential

4) energy that is released as a result of the reaction

5) energy that is released when an electron attaches to an atom.

173. The increase in reaction rate with increasing temperature is usually characterized by:

1) rate constant of a chemical reaction

2) chemical equilibrium constant

Test papers:

OPTION 1

Part 1

A1. The element of the third period of the main subgroup of the III group of PSHE is:

A2. Designation of an isotope whose nucleus contains 8 protons and 10 neutrons:

A3. An atom of a chemical element whose electron shell contains 17 electrons:

A4. An atom has two electronic layers (energy levels):

A5. A pair of chemical elements that have 5 electrons at the outer electronic level:

A6.

A. In a period, the metallic properties of atoms of elements increase with increasing atomic number.

B. In a period, the metallic properties of atoms of elements weaken with increasing atomic number.

Part 2

IN 1.

Particle:	Electron distribution:
	1) 2e, 8e, 8e, 2e
	2) 2e, 8e, 2e
	4) 2e, 8e, 3e
	5) 2e, 8e, 18e, 4e

AT 2. Compounds with ionic bonds are:

AT 3. The relative molecular weight of barium chloride BaCl2 is __________.

Part 3

C1. Give the characteristics of the element with Z = 11 (Appendix 3, points I (1-5), II (1-4)). Write down the structure diagram of its Na+ ion.

Dear eighth grader!

40 minutes are allotted to complete the test. The work consists of 3 parts and includes 10 tasks.

Part 1 includes 6 basic level tasks (A1-A6). For each task there are 4 possible answers, of which only one is correct. For completing each task - 1 point.

Part 2 consists of 3 advanced level tasks (B1-B3), to which you must give a short answer in the form of a number or a sequence of numbers. For completing each task - 2 points.

Part 3 contains 1 of the most complex voluminous tasks C1, which requires a complete answer. For completing the task you can get 3 points.

Points received for completed tasks are summed up. You can score a maximum of 15 points. I wish you success!

Performance evaluation system:

OPTION-2

Part 1

A1. The element of the second period of the main subgroup of the III group of PSHE is:

A2. Designation of an isotope whose nucleus contains 26 protons and 30 neutrons:

A3. An atom of a chemical element whose nucleus contains 14 protons is:

A4. An atom has three electronic layers (energy levels):

A5. A pair of chemical elements that have 3 electrons at the outer electronic level:

A6. Are the following statements true?

A. In the main subgroup, the nonmetallic properties of atoms of elements increase with increasing atomic number.

B. In the main subgroup, the nonmetallic properties of atoms of elements weaken with increasing atomic number.

Part 2

IN 1. Establish a correspondence between the particle and the distribution of electrons over energy levels:

Particle:	Electron distribution:
	1) 2e, 8e, 7e
	2) 2e, 8e, 2e
	4) 2e, 8e, 8e
	6) 2e, 8e, 8e, 1e

AT 2. Compounds with a covalent polar bond are:

AT 3. The relative molecular weight of aluminum oxide Al2O3 is _______.

Part 3

C1. Give the characteristics of the element with Z = 16 (Appendix 3, points I (1-5), II (1-4)). Write down the structure diagram of its ion S2-.

Answers.

Part 1

Option 1
Option 2

Part 2

Option 1
Option 2

Part 3

Characteristics plan	Option 1	Option 2
I. Position element in periodic system:
1. serial number, name
(big, small)
4. group, subgroup	1, main	6, main
5. relative atomic mass
II. Structure atom of an element
1. charge of the nucleus of an atom
2. formula atomic composition (number p; n; e -)	Na (11p;12n;) 11 e-	S (16p; 16n;) 16 e-
atomic structure
4. formula electronic configurations		1s2 2s2 2p6 3s23p4
5. number e - at the last level, metal or non-metal		6, non-metal
III. Comparison metallic and non-metallic properties with neighbors:
1. by period
2. by group (metal with non-metal do not compare)
Structure diagram and she

Test No. 2

Topic: SECOND GROUP OF THE PERIODIC SYSTEM

1 Characteristics. Atoms of elements of group 2 of the periodic table in the outer electron layer have 2 electrons located at a considerable distance from the nucleus. Therefore, these 2 electrons are relatively easily split off from atoms, which turn into positive doubly charged ions.

The difference in the structure of the second outer layer of a number of elements of the second group determines the existence of two subgroups: the main one, including alkaline earth metals (beryllium, magnesium, calcium, strontium, barium, radium) and a secondary subgroup, including the elements: zinc, cadmium and mercury.

All elements included in the main subgroup, except beryllium and radium, have pronounced metallic properties.

The higher the atomic mass, the more electropositive the metal. Thus, barium is as strong a reducing agent as the alkali metals. With water, oxides of alkaline earth metals form hydroxides, the solubility of which increases from beryllium hydroxide to barium hydroxide. The basic character of these compounds increases in the same sequence.

The elements of the secondary subgroup (Zn, Cd, Hg), as well as the elements of the main subgroup, exhibit an oxidation state of +2, but there is also a difference between them due to the different sizes of the radii of their atoms and ionization potentials.

The metallic properties of elements of the secondary subgroup weaken from zinc to mercury. Their hydroxides are insoluble in water and have weakly basic properties.

The elements of interest for medicine are Mg, Ca, Ba, Zn and Hg. All these elements are part of the structure of the most important drugs.

The most toxic of the group II elements is barium in its soluble compounds, which are used only as reagents and poisons for insects and rodents. In medicine, barium sulfate, a salt practically insoluble in any solvent, is used mainly.

2. MAGNESIUM COMPOUNDS

Magnesium is widely distributed in nature. It is not found in free form, but only in the form of carbonate compounds, forming minerals dolomite MgC0 3 *CaCO 3 and magnesite MgC0 3.. Magnesium is part of silicates - talc 3MgO*4Si0 2 *H 2 0, etc.

Magnesium salts are also found in soil, natural waters, especially sea waters, and many mineral springs.

The value of magnesium is great. It is part of the green plant pigment - chlorophyll, participating in the process of plant photosynthesis.

Magnesium compounds play a significant role in the activity of the central nervous system of living organisms.

According to its physiological action, magnesium is a calcium antagonist. So, if magnesium salts cause anesthesia and paralysis, then calcium compounds relieve this phenomenon. On the contrary, the effect exerted by calcium compounds is removed by magnesium.

Pharmacopoeial preparations of magnesium are: magnesium oxide, burnt magnesia, basic magnesium carbonate, white magnesium, Magnesium trisilicate, magnesium sulfate.

The first three drugs exhibit an antacid effect, i.e. they are used for increased acidity of gastric juice. They act in the same way as very mild laxatives. Magnesium sulfate is used as a sedative, antispasmodic and laxative.

Magnesium sulfate Magnesii sulfas

MgS0 4- 7H 2 0 M. m. 246.50

Magnesium sulfate as a remedy was first used in England - Epsom or bitter salt.

A) Receipt. Magnesium sulfate is distributed in nature in the form of kieserite MgS0 4 *7H 2 0. Magnesium sulfate is found in large quantities in sea water.

A preparation is obtained from magnesite MgC0 3 by treating it with sulfuric acid.

MgCO 3 + H 2 S0 4 > MgS0 4 + C0 2 + H 2 O

The resulting solution is concentrated by evaporation until crystallization, resulting in MgS0 4 *7H 2 0.

B) Properties. Colorless prismatic crystals, weathering in air, salty-bitter taste, odorless. It is highly soluble in water, practically insoluble in alcohol.

B) Authenticity

GF - for Mg 2+ , the formation of a precipitate of double ammonium and magnesium phosphate when

interaction of the drug with dibasic sodium phosphate in an ammonia solution in the presence of ammonium chloride.

MgS0 4 + Na H P0 4 + NH 4 OH = MgNH 4 P0 4  + Na 2 S0 4 + H 2 0

White

If this reaction is carried out using the drop method on a glass slide, crystals of a certain shape are formed, which can serve as confirmation of the authenticity of the drug.

GF - With organic rheu-8-hydroxyquinoline, in the presence of an ammonia solution with the addition of ammonium chloride NH 4 C1 produces magnesium hydroxyquinolate, colored greenish-yellow.

GF - The sulfate ion opens with a solution of barium chloride a white milky precipitate of barium sulfate precipitates. Insoluble in acids and alkalis.

MgS 0 4 + BaС1 2 = Mg С1 2 + BaS 0 4 

D) Cleanliness . Arsenic no more than 0.0002%, chlorides, heavy metals, moisture are allowed.

The preparation used for injection Solutio Magnesii sulfatis 20% aut 25% pro injectionibus is tested for manganese.

GF complexometric titration method. An ammonia buffer solution and a special acid chromium black special indicator are added to the drug solution and titrated with Trilon B until the red color turns blue. D.b. 99% -102%

E) Application. Myotropic antispasmodic, laxative.

Used as a laxative, 15 x 30 g per oral dose.

When administered parenterally, magnesium sulfate has a calming effect on the central nervous system.

As an antispasmodic for hypertension in the form of a 25% solution (subcutaneous);

For labor pain relief, 10 x 20 ml of 25% solution intramuscularly;

As an anticonvulsant in the same doses “as for pain relief in childbirth;

As a choleretic agent, 20 x 25% solution (orally).

In case of respiratory depression associated with overdose (curarepod), a 10% calcium chloride solution is used intravenously.

Release: powder, 10%, 20%, 25% solution in ampoules, 2.5, 10 and 20 ml.

Powder in bags 10.0-50.0. Cormagnesin, 32% magnesium-Diasporal forte

g) Storage: dry, cool place.

3. CALCIUM COMPOUNDS

Calcium is widely distributed in nature. Due to its high chemical activity, it is found in nature only in a bound state. It occurs in the form of numerous deposits of limestone, chalk and marble - these are natural varieties of calcium carbonate CaCO3. Gypsum CaS0 is also found 4 -2H 2 0, phosphorite Ca 3 (P0 4) 2 and silicates.

All natural calcium compounds, especially carbonates, serve as sources for medical calcium preparations; marble is often used as the purest.

Calcium plays an important role in the functioning of the body. It is part of dental tissue, bones, nervous tissue, muscles, and blood. Calcium ions enhance the vital activity of cells, promote the contraction of skeletal muscles and heart muscles, and are necessary for the formation of bone tissue and the blood clotting process.

With a decrease in the concentration of calcium ions in the blood, muscle excitability increases, which often leads to cramps. Solutions of calcium salts relieve itching caused by an allergic condition, so they are classified as antiallergic drugs.

Of the calcium compounds used in medicine, calcium oxide (burnt lime), burnt calcium sulfate (burnt gypsum), precipitated calcium carbonate (precipitated chalk), calcium chloride and salts of organic acids (calcium glycerophosphate, calcium gluconate, etc.). The pharmacopoeial drug is calcium chloride.

Calcium chloride Calcii chloridum

CaC1 2 -6H 2 0 M. m. 219.08

A) Receipt. Calcium chloride, intended for medical purposes, is obtained by treating calcium carbonate (marble) with hydrochloric acid.

CaCO 3 + 2HC1 = CaCl 2 + C0 2 + H 2 O

Pure calcium chloride CaCl crystallizes from water 2 -6Н 2 0.

B) Properties. They are colorless, odorless prismatic crystals with a bitter-salty taste; very easily soluble in water, causing a strong cooling of the solution. Easily soluble in 95% alcohol.

The drug is very hygroscopic and dissolves in air. At a temperature of 94°C it melts in its water of crystallization. Aqueous solutions have a neutral reaction. When heated to 200°C, it loses part of its water of crystallization and turns into calcium chloride dihydrate CaCl1 2 -2H 2 0, The hygroscopicity of the drug and its ability to dissolve under the influence of moisture make the composition of this salt inconsistent, which can lead to inaccurate dosage when preparing drugs with calcium chloride. Taking this into account, pharmacies prepare a 50% solution of it (Calcium chloratum solutum 50%) and the necessary drugs are prepared from this concentrate.

B) Authenticity:

GF - reaction to Ca 2+ reaction with ammonium oxalate,

(NH 4 ) 2 C 2 0 4 + CaC 1 2 = CaC 2 0 4  + 2NH 4 Cl

White sediment

The precipitate is soluble in mineral acids and insoluble in acetic acid.

Formation of a white precipitate when the drug interacts with sulfuric acid or alkali metal sulfates.

CaCl 2 + H 2 S0 4 = CaS0 4  + 2HC1

White sediment

The calcium sulfate precipitate dissolves in ammonium sulfate to form a colorless complex.

GF- calcium salts color the burner flame brick red.

GF for chlorides with silver nitrate

CaCl 2 + Ag N O 3 = Ag Cl  + Ca (N O 3 ) 2

White curdled sediment

D) Cleanliness . Impurities of soluble salts of barium, iron, aluminum, and phosphates are not allowed in the preparation.

Sulfates, heavy metals, arsenic, and magnesium salts according to standards are allowed.

D) Quantitative definition

GF - determined complexometrically with the indicator acid chromium dark blue. When titrating with Trilon B, when adding an ammonia buffer solution, the color of the solution changes from cherry-red to bluish-lilac (indica eriochrome black special T). Must be at least 98.0%.

Photometric, - argentometry (Mora)

Fluorometric, - refractometry

By weight (oxalate).

E) Application. Antiallergic

As a hemostatic agent for pulmonary, gastrointestinal, nasal and uterine bleeding;

In surgical practice to increase blood clotting;

For allergic diseases (bronchial asthma, urticaria) to relieve itching;

As an antidote for poisoning with magnesium salts.

Anti-inflammatory, for colds

The drug is prescribed orally in the form of a 5 x 10% solution, intravenously as a 10% solution. It cannot be administered subcutaneously or intramuscularly, as in this case necrosis may occur.

Release form: powder, 10% solution in ampoules.

g) Storage. In small, well-sealed glass jars with a stopper, filled with paraffin, in a dry place.

4. ZINC COMPOUNDS

In nature, zinc occurs in the form of minerals: gallite ZnCO 2 and zinc blende ZnS. Zinc is found in muscle, dental and nervous tissue of the human body. The use of zinc compounds in medicine is based on the fact that zinc produces compounds with proteins - albuminates, soluble albuminates have an effect ranging from weakly astringent to cauterizing. Insoluble albuminates usually form a film on the tissue surface and thus promote tissue healing (drying effect).

Zinc compounds are toxic in large doses; when applied topically, they can be used as astringents and cauterizing agents. When administered orally, zinc compounds cause vomiting.

Pharmacopoeial preparations of zinc are zinc oxide and zinc sulfate.

Zinc sulfate Zinci sulfas

ZnSO4 *7H 2 0 M. m. 287.54

Zinc sulfate has been used in medicine since ancient times under the name of white sulfate, in contrast to colored copper and iron sulfate.

A) Receipt. From natural ore zinc blende ZnS, which is roasted. In this case, zinc sulfide is converted into an oxide, which is then treated with dilute sulfuric acid, resulting in the formation of zinc sulfate in solution. 2 ZnS + ZO 2 = 2 ZnO + 2 SO 2

ZnO + Ha 2 S 0 4 = ZnS 0 4 + 4 H 2 O

A solution containing zinc sulfate is evaporated until the salt crystallizes in the form of heptahydrate (ZnS0 4 -7H 2 0).

B) Properties. Colorless transparent crystals or finely crystalline powder, having an astringent metallic taste, odorless, very easily soluble in water, slowly in glycerin, insoluble in alcohol. It erodes in the air.

B) Authenticity.

GF - Sulfate ion is determined by the formation of a white precipitate.

ZnS0 4 + Ba Cl 2 = Ba S0 4  + Zn Cl 2

White milky, insoluble in acids and alkalis

GF- on Zn 2+ reaction with a solution of sodium sulfide produces white zinc sulfide ZnS (different from other heavy metal salts).

ZnS0 4 +Na 2 S = ZnS 4  + Na 2 S0 4

White sediment

GF - Zn 2+ reaction with a solution of potassium ferrocyanide a white-yellowish crystalline precipitate of a double salt is formed, insoluble in acids, but soluble in alkalis. 3 ZnS 0 4 + 2 K 2 [Fe (CN) 6] = K 2 Zn 3 [Fe (CN) 6] 2 + 3 K 2 SO 4

White-yellowish

A specific reaction to zinc is the formation of Rinman green. ZnS 0 4 drop onto filter paper and add cobalt nitrate on top, calcinate, this produces a characteristic green color – Rinman green: CoZnO2

With dithizone ions Zn 2+ in an alkaline environment they form a red color.

D) Cleanliness . Not d.b. impurities of iron, copper, aluminum, magnesium, calcium and other heavy metals.

Arsenic admixture is allowed

D) Quantification

GF complexometry. In the presence of an ammonia buffer solution and an acid indicator, special black chrome (or eriochrome black T). Titrate with Trilon B until the color of the solution changes from cherry-red to bluish-lilac.

E) Application externally as an antiseptic and astringent

In ophthalmic practice in the form of 0.1; 0.25; 0.5% solutions. In eye drops, zinc sulfate is often prescribed along with boric acid.

In gynecological practice for douching in the form of a 0.1 x 0.5% solution.

For skin diseases: acne, acne, dermatoses.

Rarely prescribed orally as an emetic.

Release forms: powder, eye drops 0.1; 0.25; 0.5%, drops of zinc sulfate with boric acid. Combined: Zinkin, Zincteral

g) Storage. With caution in well-sealed jars. List B.

Zinc oxide Zinci oxydum

It is a white amorphous powder with a yellowish tint that easily absorbs carbon dioxide from the air. A characteristic property of zinc oxide is that when heated, it becomes yellow, and when cooled, it becomes white.

Zinc oxide is used externally in the form of powders, ointments, linements as an astringent, drying and disinfectant for skin diseases: dermatitis, prickly heat, bedsores, diaper rash, ulcers, wounds, burns.

5. MERCURY COMPOUNDS

Mercury is a liquid metal. The distribution of mercury in nature is low. It is found in native form, disseminated in rocks, but mainly in the form of mercury sulfide HgS (cinnabar) of a bright red color.

Pharmacopoeial drugs are mercury compounds having an oxidation state of +2: mercury yellow oxide, mercury dichloride, mercury amide chloride, mercury oxycyanide and mercury cyanide.

Inorganic mercury preparations are used as antiseptic, diuretic and laxatives.

The antiseptic effect of mercury compounds is based on the ability of the mercury ion to precipitate proteins. The diuretic effect of some mercury salts is associated with

in that, when excreted through the kidneys, they irritate the renal epithelium and promote urination.

Similarly, mercury compounds, released through the intestines and irritating it, exhibit a laxative effect.

Soluble mercury salts are very toxic and belong to list A.

Mercury oxide yellow Hydrargyri oxydum flavum

HgO M. m. 216.59

A) Receipt . Precipitation reactions from soluble mercury salts are used. For this purpose, mercuric dichloride or nitrate is most often used. A concentrated solution of mercury (II) salt is slowly poured into a dilute alkali solution.

Hg(NO 3 ) 2 + 2NaOH = 2NaNO 3 + HgO + H2O

Bright yellow sediment

After the mercury oxide precipitate has settled, the liquid is drained, the precipitate is washed until there is no alkaline reaction and dried. All operations should be carried out in the dark, otherwise mercuric oxide Hg may be formed 2 0 black.

B) Properties. Heavy fine powder of yellow or orange-yellow color. Insoluble in water, but easily soluble in hydrochloric, nitric and acetic acids. The light gradually darkens.

B) Authenticity for Hg2+.

To do this, it is treated with diluted hydrochloric acid to obtain a soluble mercury (II) salt, in which the Hg cation is determined 2+

HgO + 2HC1 = HgCl 2 + H.0

GF - reaction with alkali solutions, a precipitate of yellow mercury oxide precipitates.

HgCl 2 + 2KOH > HgO  + 2KS + H 2 0

Bright yellow sediment

GF - reaction with potassium iodide solution; A bright red precipitate of mercury diiodide is formed, which dissolves in excess potassium iodide.

HgCl 2 + 2Kl = HgJ 2  + 2KCl HgJ 2 + 2KI > K 2

Bright red colorless solution

A solution of this complex salt is known as Nessler's reagent and is used as a very sensitive reagent for NH 4+;

GF - reaction with sodium sulfide solution; a brown-black precipitate is formed, insoluble in dilute nitric acid.

HgCl 2 + NaS = HgS  + 2NaCl

Brownish-black sediment

D) Quantitative content

GF - neutralization indirectly through interaction with potassium iodide. When yellow oxide is exposed to mercury with a solution of potassium iodide, a soluble complex salt and alkali are formed, which is titrated with acid against methyl orange HgO + 4 KI + H 2 O > K 2 [Hgl 4 ] + 2KOH

KON +NS1 = KS1 + N 2 0

Rhodanometric method: yellow mercuric oxide is dissolved in nitric acid, and the resulting salt is titrated with ammonium thiocyanate in the presence of ferroammonium alum until it turns red.

G) Application as a gentle antiseptic for the preparation of eye ointments 2%.

E) Store should be taken with caution in well-sealed dark glass jars, since mercuric oxide may form in the light, which is detected by the darkening of the preparation. List B.

Topic FIRST GROUP OF THE PERIODIC SYSTEM

1.Characteristics.All elements that make up the first group of the periodic table have only an electron in their outer electron layer, which they easily give up, turning into singly charged positive ions. This explains their very high reactivity towards electronegative elements such as halogens.

The main subgroup includes lithium, sodium, potassium, rubidium, cesium and francium. The side group consists of copper, silver and gold.

The elements of the main subgroup are called alkali metals, since their oxides, when interacting with water, form strong alkalis. Alkali metal salts are used in medicine.

The most widely used in medicine are sodium and potassium salts, described above in preparations derived from halogens.

2. COMPOUNDS OF COPPER AND SILVER

A secondary subgroup of elements of the first group consists of copper, silver and gold. They have a tendency to form complexes, especially copper, and also the ability to be reduced from compounds to free metal, with silver being reduced more easily than copper.

Of the inorganic copper compounds, copper sulfate is used in medicine. When taken orally, it has an emetic effect; as an external remedy it is used for catarrh of the mucous membranes and ulcers due to its astringent, irritating and cauterizing effect.

Silver belongs to the “noble” metals. In nature, it occurs mainly in the form of compounds with sulfur (Ag 2 S).

The use of silver preparations in medicine is based on its bactericidal properties. It has been proven that silver ions kill gram-positive and gram-negative microorganisms, as well as viruses. Silver preparations are used in medicine internally and externally as astringent, antiseptic and cauterizing agents in the treatment of skin, urological and eye diseases.

Of the silver compounds, the most widely used is silver nitrate (AgN03), as a good astringent and cauterizing agent. In medicine, colloidal preparations are also used, where silver is bound to protein and is only partially ionized. In colloidal silver preparations, only the disinfecting properties of silver are retained and its cauterizing effect disappears.

All soluble copper and silver compounds are poisonous.

3. Silver nitrate Argenti nitras

AgN0 3

A) Receipt by dissolving a copper-silver alloy in nitric acid when heated. To clean the resulting silver nitrate from impurities, it is precipitated with hydrochloric acid in the form of silver chloride. The latter is reduced with zinc, and silver, freed from impurities, is again dissolved in nitric acid.

The resulting silver nitrate is treated with a small amount of water, and crystals crystallize when standing. The isolated crystals are filtered, washed with water and dried in the dark.

B) Properties colorless transparent crystals in the form of plates or cylindrical rods of a radiant-crystalline structure in a fracture. Easily soluble in water, difficult in alcohol. Crystals darken in light.

B) Authenticity

GF - Ag+ : with hydrochloric acid or its salts a white precipitate of silver chloride precipitates, insoluble in nitric acid and highly soluble in ammonia solution AgNO 3 + HCl = AgCI  + HNO 3

White

AgCl + 2NH 4 0H = Cl + 2H 2 O

GF - Ag+ reduction to free silver (reaction of formation of a silver mirror). A solution of formaldehyde is added to the ammonia solution of silver oxide and the liquid is heated. After some time, a coating of metallic silver in the form of a mirror forms on the walls of the vessel.

[ Ag (NH 3 ) 2 ] OH + HSON = 2Ag  + HCOOH + 4 NH 3 + 2 H 2 O

Black sediment

Ag+ with potassium chromate, a brownish-red precipitate of silver chromate precipitates. 2AgNO 3 + K 2 Cr0 4 = Ag Cr0 4  + 2KNO 3

Brownish-red precipitate

The precipitate is soluble in nitric acid, ammonium hydroxide, and sparingly soluble in acetic acid.

GF - Nitrate ion determined with diphenylamine in con. Sulfuric acid produces a blue color

Formation of a brown ring when silver nitrate reacts with ferrous sulfate in concentrated sulfuric acid.

Nitrate ion Potassium permanganate does not discolor in an acidic environment, unlike nitrite.

D) Cleanliness allowable acidity limit

Salts of heavy metals (lead, copper, bismuth) are not allowed.

D) Quantitativecontent - Volhard precipitation method, titrated with ammonium thiocyanate (rodanide)

AgNO 3 + NH 4 SCN = AgSCN + NH 4 NO,

White sediment

3NH 4 SCN + (NH 4 )Fe(S0 4 )= Fe(SCN) 3 + 2(NH 4 ) 2 S0 4

The indicator is ferroammonium alum until it turns red. D.b. less than 99.75%.

G) Application antiseptic and cauterizing. The latter is due to the ability of silver nitrate to coagulate proteins, turning them into insoluble compounds, which is used to cauterize wounds and ulcers. For this purpose, silver nitrate in the form of sticks (Stilus Argenti nitrici) is used.

In small concentrations it has an astringent and anti-inflammatory effect. Used externally for erosions, ulcers, acute conjunctivitis, trachoma in the form of 2510% aqueous solutions, as well as ointments (12%). It is prescribed orally in the form of a 0.05 x 0.06% solution for gastric ulcers and chronic gastritis. Release form: powder, lapis sticks.

IRR orally 0.03 g, IRR 0.1

E) Storage in well-sealed dark glass jars, since it can decompose in the light, which is detected by the darkening of the drug. List A.

4. Protargol Protargolum, Argentum proteinicum Silver proteinate

A) Receipt from silver nitrate and protein (casein, gelatin, egg white, peptone)

Protected colloid: contains silver oxide (7.8 x 8.3%) and albumin hydrolysis products.

B) Properties Light amorphous powder of yellow-brown color, odorless, slightly bitter, slightly astringent taste. Easily soluble in cold water, insoluble in alcohol.

B) Authenticity

GF- Protein is determined by the appearance of the smell of burnt horn and the charring of the preparation when heated.

GF- the residue from combustion (it is white) is dissolved in HNO 3 and carry out reactions on Ag+ with chlorides.

- (biuret re-I) the drug is boiled with dil. HCl, a precipitate forms, it is filtered and NaOH and C are added to the clear filtrate uS O 4, A violet color appears (on protein).

D) Cleanliness not d.b. impurities of silver compounds, protein decomposition products.

D) Quantitativedefinition: after ashing the preparation with sulfuric acid. Argentometry method, Volhard version. D.b. 7.88.3%

G) Application

Antibacterial, anti-inflammatory agent. Used externally in ophthalmology 1-2% solution (conjunctevitis, blenorrhea, blepharitis), urology 0.1-1% (bladder lavage), otorhinolaryngology (ears, nose), gynecology. Orally for stomach ulcers and intestinal diseases.

Release form: powder and dosage form in pharmacies.

E) Storage : according to list B. In well-closed dark glass jars

5. Collargol (Collargolum, Argentum colloidale, Silver colloid)

Colloidal system with 70-75% content of highly dispersed metallic silver and protective proteins (hydrolysates of casein and gelatin).

Greenish-black or bluish-black plates with a metallic sheen, soluble in water to form a colloidal solution. When treated with water, it swells and forms alkaline, negatively charged sols.

Antibacterial agent. Apply:

0.2 1% solutions for washing purulent wounds;

1 2% solutions for washing the bladder for chronic cystitis and urethritis,

2 x 5% solutions in the form of eye drops for the treatment of purulent conjunctivitis and blenorrhea.

For erysipelas and chancre, 15% ointment is sometimes prescribed.

Rarely in septic conditions: intravenous administration.

Storage: according to list B. In well-closed dark glass jars