The p-elements of the periodic system include elements with a valence p-sublevel. These elements are located in III, IV, V, VI, VII, VIII groups, main subgroups. In a 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 rather than 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, a sharp transition from non-metal to metals can be traced. The properties and behavior of boron are similar, which is the result of the diagonal affinity of elements in the periodic system, according to which a shift in the period to the right causes an increase in the non-metallic character, and down the group - a metallic one, therefore elements similar in properties turn out to be located diagonally side by side, for example Li and Mg, Ber and Al, B and Si.
The electronic structure of the valence sublevels of Group III p-element atoms 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 constitute the main subgroup. The atoms of these elements have complete outer electron layers, so the electronic configuration of the valence sublevels of their atoms in the ground state has the form 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 the free state, they exist in the form of atoms (monatomic molecules). Helium (1s 2), neon (2s 2 2p 6) and argon (3s 2 3p 6) atoms have a particularly stable electronic structure, so valence-type compounds are unknown to them.
Krypton (4s 2 4p 6), xenon (5s 2 5p 6) and radon (6s 2 6p 6) differ from the previous noble gases in larger atomic sizes and, accordingly, lower ionization energies. They are able to form compounds that often have low resistance.
p-elements are:
P-elements include non-transition metals and most non-metals. P-elements have different properties, both physical and mechanical. P-non-metals are highly reactive, as a rule, substances with a strong electronegativity, P-metals are moderately active metals, and their activity increases towards the bottom of the PSCE.
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Elements in Mendeleev's periodic system are divided into s-, p-, d-elements. This subdivision is carried out on the basis of how many levels the electron shell of the element atom has and what level the filling of the shell with electrons ends with.
TO s-elements refer 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. With an increase in the radius in the Li-Fr group, the bond of the valence electron with the nucleus weakens and the ionization energy decreases. Atoms of alkaline elements easily donate their valence electron, which characterizes them as strong reducing agents.
Restorative properties are enhanced with increasing serial number.
TO p-elements include 30 items IIIA-VIIIA-groups periodic system; p-elements are located in the second and third small periods, as well as in the fourth to sixth large periods. Elements IIIA-groups have one electron in the p orbital. V IVA-VIIIA-groups the filling of the p-sublevel up to 6 electrons is observed. General electronic formula of p-elements ns2np6. In periods with an increase in the nuclear charge, the atomic radii and ionic radii of p-elements decrease, the 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 serial number.
TO d-elements includes 32 elements of the periodic system IV–VII big periods. V IIIB-group the atoms have the first electron in the d-orbital, in subsequent B-groups the d-sublevel is filled up to 10 electrons. General formula of the outer electron shell (n-1)dansb, where a=1?10, b=1?2. With an increase in the ordinal N, the properties of the d-elements change slightly. For d-elements, the atomic radius slowly increases, and they also have a variable valence associated with the incompleteness of the pre-external d-electron sublevel. In the lower oxidation states, d-elements show metallic. St. Islands, with an increase in the order. N in groups B they decrease. In solutions, d-elements with the highest degree of oxidation show acidic and oxidizing properties, and vice versa at lower degrees of oxidation. Elements with int. step. oxidation show amphoteric. St. Islands.
covalent bond.
The chemical bond carried out by common electron pairs arising in the shells of the bound atoms having antiparallel spins is called atomic or covalent bond. Covalent bond is two-electron and two-center (holds nuclei). An atom at its outer energy level can contain from one to eight electrons. Valence electrons– electrons of the pre-outer, outer electron layers involved in chemical bonding. Valence- the property of the atoms of an element to form a chemical bond.
The p-block contains the last six elements of the main subgroup, excluding helium (which is in the s-block). This block contains all non-metals (excluding hydrogen and helium) and semi-metals, as well as some metals.
The P-block contains elements that have different properties, both physical and mechanical. P-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.
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General characteristics of p-elements
The general electronic formula of p-elements is ns 2 np 1 ¸6 , where n is the main quantum number. Most p-elements are non-metals. Elements such as Al, Ga, In, Tl, Sn, Pb, Sb, Bi, Po are conditionally considered as metallic, although they retain many of the properties of non-metals. All valence electrons of p-elements are in the outer level, so they belong to the main subgroups.
Atoms of p-elements are capable of exhibiting both positive and negative oxidation states. As a rule, atoms of p-elements exhibit variable valency, moreover, in even groups it is even, and in odd groups it is odd.
In a period, as the number of p-electrons at the outer level in the atoms of elements increases, the radius of the atoms decreases, the ionization energy and the energy of electron affinity increase, i.e. the oxidizing properties (the ability to accept electrons) of the atom are enhanced. p-elements, being oxidizing agents, can also exhibit reducing properties; therefore, most p-elements are capable of disproportionation reactions. For instance:
CaO + 3C = CaC 2 + CO
2As + 3NaOH = AsH 3 + Na 3 AsO 3
3S + 6KOH = 2K 2 S + K 2 SO 3 + 3H 2 O
Within the subgroup, from top to bottom, as the atomic number of the element increases, the non-metallic properties of p-elements weaken and metallic properties increase, so the most characteristic positive oxidation state decreases. For example, the characteristic oxidation state of elements:
in period III Al 3+, Si 4+, P 5+, S 6+
in the VI period Tl 1+, Pb 2+, Bi 3+, Po 4+
from this we can conclude that the compounds Tl 3+ , Pb 4+ , Bi 5+ are strong oxidizing agents, and the compounds Ga 1+ , Ge 2+ , As 3+ are reducing agents.
The strength of hydrogen compounds in the main subgroups decreases from top to bottom due to an increase in the radius of the atom. For instance:
CH 4 ® SiH 4 ® GeH 4 ® SnH 4 ® PbH 4 ; NH 3 ® PH 3 ® SbH 3 ® BiH 3 .
Almost all p-elements are acid formers, and the stability and strength of oxygen-containing acids increases as the degree of oxidation of the p-element increases. For example, the strength of acids increases in the series:
HClO ® HClO 2 ® HClO 3 ® HClO 4 ; H 2 SO 3 ® H 2 SO 4; HNO 2 ® HNO 3 .
The redox properties of compounds of p-elements depend, as a rule, on the degree of oxidation of their atoms that make up these compounds. Compounds in which the p-element atom is in an intermediate oxidation state can exhibit both oxidizing and reducing properties (H 2 O 2, N 2 H 4, NH 2 OH, HNO 2, H 3 PO 2, H 2 SO 3 etc.).
p-elementsGroup VII (halogens)
Topic work plan:
1. General characteristics of the properties of p-elements of group VII, being in nature, obtaining. Physical and chemical properties of simple substances.
2. Compounds in the lowest oxidation state: hydrogen halides, hydrohalic acids and their salts. Receipt. restorative properties.
3. Compounds in positive oxidation states: oxygen-containing acids, their production, stability, acid strength and redox properties. Salts of oxygen-containing acids, preparation, chemical properties.
Exercise 1
1. Why are odd valences more characteristic of halogens than even ones? Justify your answer in terms of the theory of the structure of the atom.
2. How to explain the existence of hydrofluorides? Why don't chlorine, bromine and iodine form analogous compounds? Justify the answer.
3. What substances does hydroiodic acid interact with: a) Ca; b) P 2 O 3; c) NaOH. Justify your answer by writing the equations for the corresponding reactions.
4. How should the concentrations of reactants be changed in order to increase the yield of chlorine: O 2 + 4HCl 2Cl 2 + 2H 2 O? Justify your answer using Le Chatelier's principle.
5. Determine the oxidizing agent in the following reactions: a) I 2 + H 2 O 2 → HIO 3 + H 2 O
b) HIO 3 + H 2 O 2 → O 2 + I 2 + H 2 O.
Arrange the coefficients using the electron-ion balance method. Determine the equivalent of the oxidizer and calculate the molar mass of the equivalent of the oxidizer.
Task 2
1. Why can't halogen molecules contain more than two atoms? Justify your answer in terms of the valence bond method (MVS).
2. Why is a hydrofluoric acid molecule written H 2 F 2? Justify the answer.
3. How many σ-bonds are in the molecules of oxygen-containing chlorine acids? Justify your answer in terms of the valence bond method (MVS).
4. Is it possible to prepare solutions containing the following salts: a) NaCl and KNO 3 ; b) NaCl and AgNO 3 ; c) NaCl and AgF. Justify your answer by writing the equations for the corresponding reactions.
5. Write the reaction equation for the interaction of potassium bromide with potassium dichromate at pH< 7. Расставьте коэффициенты в уравнении методом электронно-ионного баланса. Определите эквивалент восстановителя, рассчитайте его молярную массу.
Task 4
1. The molecule of which of the following compounds is more polar: a) HF; b) HCl; c) HI; d) HBr? Why? Justify your answer in terms of the method of valence bonds.
2. Which of the following substances convert bromine into a soluble state: a) H 2 O; b) H 2 SO 4 solution; c) NaOH solution; d) benzene. Justify the answer.
3. Which of the following substances will interact with hydrochloric acid: a) Cu; b) P; c) MgO; d) NaOH. Justify your answer by writing the equations of the corresponding chemical reactions.
4. Write a mathematical expression for the chemical equilibrium constant for the reaction: H 2 (g.) + I 2 (g.) Û 2HI (g.) . How should the concentrations of reactants be changed in order to reduce the yield of hydrogen iodine? Justify your answer in terms of Le Chatelier's principle.
5. Complete the equation for the following redox reaction: NaHSO 3 + NaIO 3 + H 2 O → NaHSO 4 + I 2 +. . . Arrange the coefficients using the electron-ion balance method. Determine the equivalent of the oxidizing agent and reducing agent, calculate the molar mass of the equivalent of the oxidizing agent and reducing agent.
Task 5
1. Which of the hydrogen halides in an aqueous solution has the highest degree of dissociation: a) HF; b) HCl; c) HI; d) HBr. Why? Justify the answer.
2. Compose the electronic formulas of the chlorine atom and the ion Cl -. Explain from the point of view of the theory of the structure of the atom, why, under normal conditions, the chlorine atom does not exist in a free state, but the Cl ion exists (in an aqueous solution, in a crystal lattice)?
3. In which of the following transformations is the oxidation process indicated: a) Cl - → Cl 0 ; b) Cl 5+ → Cl; c) I 0 → I 5+ ? Justify your answer by writing the electronic reaction equations.
4. How do the following substances react with each other: a) Cu and F 2 ; b) Fe and Cl 2; c) Ca and Br 2 ; d) Zn and I 2 . Write the equations for the corresponding reactions and give the names of the reaction products.
5. Complete the chemical reaction equation: Na 2 S + NaBrO + H 2 SO 4 →. . . Arrange the coefficients in the equation using the electron-ion balance method. Determine the equivalent of the oxidizing agent and reducing agent, calculate the molar mass of the equivalent of the oxidizing agent and reducing agent.
Task 6
1. Why does fluorine never show a positive oxidation state? Justify the answer.
2. How many σ-bonds are in the molecules of oxygen-containing acids of halogens in the +5 oxidation state?
3. In which of the following transformations is the recovery process indicated: a) I־ → I 0 ; b) Cl 3+ → Cl 5+; c) Cl 3+ → Cl־. Justify your answer by writing the electronic reaction equations.
4. Which of the given examples of chemical reactions correspond to the brief ionic equation Ag + + Cl־ = AgCl:
a) AgNO 3 + HCl → ...; b) Ag 2 SO 4 + NaCl → ...; c) Ag 2 O + HCl → ....
5. Complete the equation for the reaction at pH > 7: MnCl 2 + KClO + . . . →
If the color is known to turn green as a result of the reaction. Arrange the coefficients in the equation using the electron-ion balance method. Calculate the mass of manganese(II) chloride required to react with 5 mol equivalents of KClO under these conditions.
p-elementsGroup VI
Topic work plan:
1. General characteristics of the properties of p-elements of group VI.
2. Oxygen. allotropic modifications. The structure of oxygen and ozone molecules. Oxides, peroxides, superoxides, ozonides. Getting and properties.
3.Water. Anomaly in the physical properties of water. Chemical properties of water. Hydrogen peroxide, production methods, molecular structure, chemical properties (acid-base and redox).
4. Sulfur. Allotropic modifications, physical and chemical properties of a simple substance.
5. Hydrogen sulfide. The structure of molecules, obtaining, physical and chemical properties. Hydrosulphuric acid, sulfides and persulfides, their properties, preparation and application. Reducing properties of sulfur compounds in the lowest oxidation state.
6. Sulfur oxides, halides and oxohalides. Oxygen-containing sulfur acids, characterization of acidic and redox properties of acids and their derivatives. Sulfuric acid: obtaining, molecular structure, chemical properties. The interaction of sulfuric acid with metals. sulfates. Polythionic acids and their salts. Thiosulfuric acid and sodium thiosulfate: preparation, molecular structure, chemical properties. Sulfur peroxoacids (peracids), peroxosulfates: preparation, molecular structure, properties.
7. Elements of the selenium subgroup. Finding in nature. Properties of simple substances. Comparative characteristics of the compounds of the elements of the selenium subgroup: acid-base, redox properties.
Individual tasks
Exercise 1
1. How many milliliters (N.O.) of sulfur dioxide will be required to react with 50 ml of 0.1 N potassium hydroxide solution?
2. In which of the following transformations is the oxidation process indicated: a) S +4 → S 2 ־; b) S 2 ־→ S 0 ; c) Se +4 → Se 0 . Justify your answer by writing the electronic equations of the corresponding reactions.
3. Elemental selenium can be obtained from selenic acid by reduction with strong reducing agents. Write electron-ionic and molecular equations for the reaction of selenic acid with hydrazine, which is oxidized to nitrogen.
4. One of the common natural sulfur compounds is the mineral pyrite, the main component of which is FeS 2 sulfide, and also contains other impurities. Determine what volume of sulfur oxide (IV) can be obtained (n.o.) by firing 600 g of pyrite if the mass fraction of impurities in it is 20%.
5. Calculate the mass fraction of salt in the solution obtained after complete neutralization of a 40% sulfuric acid solution with a 15% sodium hydroxide solution.
Task 2
1. What volumes (n.o.) of hydrogen sulfide and sulfur oxide (IV) must react with each other so that the mass of sulfur formed is 100 kg?
2. What is the spatial configuration (geometry) of the sulfate ion: a) square; b) a quadrangular pyramid; c) tetrahedron. Why? Justify your answer in terms of the theory of the structure of the atom.
3. Why can hydrogen peroxide exhibit both oxidizing and reducing properties? Compose the electron-ionic and molecular equations for the reactions of hydrogen peroxide: a) with a solution of potassium permanganate acidified with sulfuric acid; b) with a solution of potassium iodide.
4. What mass of a solution with a mass fraction of sulfuric acid of 70% can be obtained from pyrite weighing 200 kg containing FeS 2 and impurities? The mass fraction of impurities in pyrite is 10%, and the yield of sulfuric acid is 80%.
5. 30 g of hydrogen sulfide was passed through a solution containing 10 g of sodium hydroxide. What salt was formed in this case? Determine its mass.
Task 3
1. How many liters of sulfur dioxide (N.O.) can be obtained by reacting 6.5 g of copper with concentrated sulfuric acid?
2. What salts undergo hydrolysis in an aqueous solution: a) K 2 SO 4; b) Al 2 (SO 4) 3; c) Al 2 S 3; d) K 2 S. Justify your answer by writing molecular and ion-molecular reaction equations.
3. What is the degree of oxidation of oxygen in compounds: O 2; O 3 ; Na 2 O; H 2 O 2 ; KO2; KO 3? Sodium peroxide absorbs ammonia, oxidizing it as much as possible. Write the molecular and electronic equations for the reaction.
4. The mass fraction of ozone in a mixture with oxygen is 10%. Calculate the mass of hydrogen required to react with 8 g of this mixture. Note that when hydrogen reacts with both allotropic modifications of oxygen, water is formed.
5. Calculate the mass of a sulfuric acid solution with a mass fraction of H 2 SO 4 96%, which can be obtained from pyrite weighing 3.6 kg.
Task 4
1. How much of a 10% (by mass) solution of sulfuric acid will be required to obtain 33.6 liters of hydrogen (N.O.) when interacting with zinc?
2. What salts undergo hydrolysis in an aqueous solution: a) Na 2 SO 4; b) Na 2 S 2 O 3; c) Na 2 S; d) Na 2 SO 3. Write molecular and ion-molecular reaction equations and determine the pH of the medium.
3. What properties of hydrogen peroxide are more pronounced: oxidizing or reducing? Motivate your answer with the values of the corresponding potentials. Sodium peroxide absorbs hydrogen sulfide, oxidizing it as much as possible. Write the molecular and electron-ionic equations for this reaction.
4. What volume of air and what mass of water must be taken to convert sulfur oxide (IV) with a volume of 10 liters (normal conditions) into sulfuric acid? The volume fraction of oxygen in the air is 20.95%.
5. In which case will more oxygen be obtained: with the decomposition of 5 g of potassium permanganate or with the decomposition of 5 g of potassium chlorate? Justify your answer by writing the equations of the corresponding reactions and making the necessary calculations.
Task 5
1. Determine the mass of SeO 2, upon hydration of which 3 moles of the corresponding acid are obtained.
2. What compounds can exhibit oxidizing properties: a) H 2 S; b) H 2 SO 3; c) H 2 SO 4 (razb); d) H 2 SO 4 (conc)? Why? Justify your answer in terms of the OVR theory.
3. Elemental tellurium can be obtained from H 6 TeO 6 by reduction with strong reducing agents. Write the electronic and molecular equations for the reaction of orthotelluric acid with sulfur oxide (IV).
4. Oxygen was obtained from potassium permanganate weighing 7.9 g, which reacted with magnesium. What mass of magnesium oxide will be obtained in this case?
5. Based on the structure of the oxygen atom, indicate its valence capabilities. What are the oxidation states of oxygen in compounds? Justify your answer with relevant examples.
Task 6
1. How many moles of sodium selenite are needed to react with 33.6 liters of chlorine (n.o.) according to the equation: Na 2 SeO 3 + Cl 2 + H 2 O →. . . ?
2. What compounds can exhibit reducing properties: a) H 2 S; b) H 2 SO 3; c) H 2 SO 4 (razb); d) H 2 SO 4 (conc). Why? Justify your answer in terms of the OVR theory.
3. Make electronic formulas of sulfur and selenium atoms. Are they full electronic counterparts? Justify your answer in terms of the theory of the structure of the atom.
4. List the laboratory and industrial methods for obtaining oxygen, give the equations of the corresponding reactions. Name the most important areas of practical application of oxygen.
5. How and why do acid properties change in the series: H 2 S, H 2 Se, H 2 Te?
p-elementsGroup V
Topic work plan:
1. General characteristics of the properties of p-elements of the V group, being in nature, obtaining. Physical and chemical properties of simple substances.
2. Nitrogen. Obtaining, properties and application of nitrogen in engineering. Ammonia, hydrazine, hydroxylamine, nitric acid. Their obtaining, properties, application. Liquid ammonia as an ionizing solvent. ammonia as a ligand. metal nitrides. Ammonium salts, obtaining, properties.
3. Nitrogen oxides. Obtaining, structure of molecules, properties. Oxygen-containing acids of nitrogen, properties. Salts of these acids, behavior in solution and when heated, in redox reactions. Interaction of nitric acid with metals and non-metals. Aqua regia.
4. Phosphorus, obtaining, properties, application. Phosphides and phosphines. Phosphorous acid and hypophosphites. Phosphorous anhydride and phosphorous acid. Phosphoric anhydride and phosphoric acids. Halides, oxohalides.
5. Subgroup of arsenic. Structure and properties of simple substances. Compounds with hydrogen and with metals. Oxides, sulfides, halides and oxohalides of elements - As, Sb, Bi. Thioacids and their salts. Acid-base properties of hydroxides and redox properties of arsenic, antimony and bismuth compounds in various oxidation states. Application.
Exercise 1
1. Give a comparative description of the atoms of the elements of the nitrogen subgroup, indicating: a) electronic configurations; b) valence possibilities; c) the most characteristic oxidation states.
2. What is the mass of potassium nitrite that can be oxidized in the presence of sulfuric acid with 30 ml of 0.09 N potassium permanganate solution?
3. What mass of ammonia will be required to obtain nitric acid with a mass of 12.6 tons, given that losses in production are 5%.
4. Using the electronic balance method, select the coefficients in the schemes of the following redox reactions:
a) Ca + N 2 → Ca 3 N 2
b) R 4 + O 2 → R 4 O 6
c) NO 2 + O 2 + H 2 O → HNO 3
5. Calculate the pH of a 0.1 N sodium nitrite solution and the degree of salt hydrolysis in this solution.
Task 2
1. Write the equations of chemical reactions that must be carried out to carry out the following transformations:
Pb(NO 3) 2 → NO 2 → N 2 O 4 → HNO 3 → NH 4 NO 3 → NH 3
2. What volume of 0.05 N potassium permanganate solution will be required to oxidize 20 ml of sodium arsenite solution containing 0.02 g of NaAsO 2?
3. Complete the reaction equation: Cu 2 S + HNO 3 (conc.) → .... Arrange the coefficients in the equation using the electron-ion balance method.
4. Describe the electronic structure of NH 3 , NH 4 + , HNO 3 in terms of the method of valence bonds. What is the oxidation state of nitrogen in each of these compounds?
5. Determine the mass of nitrogen, which at a temperature of 20 ° C and a pressure of 1.4 ∙ 10 5 Pa occupies a volume of 10 liters.
Task 3
1. Give examples of nitrogen compounds, in the molecules of which there are bonds formed by the donor-acceptor mechanism.
2. What volume of a 0.25 N solution of potassium permanganate will be required for the oxidation of 0.05 l of a 0.2 M solution of sodium nitrite in an acidic environment.
4. Describe the electronic structure of the N 2 molecule in terms of the VS method. What chemical properties does nitrogen exhibit as a simple substance?
5. Write the equations of the reactions that must be carried out to carry out the following transformations:
Ca 3 (RO 4) 2 → P → P 4 O 10 → H 3 RO 4 → CaHRO 4 ∙ 2H 2 O.
Task 4
1. A mixture of sulfides As 2 S 3 , Sb 2 S 3 , Bi 2 S 3 treated with a solution of sodium sulfide. What sulfide remained undissolved? Justify your answer by writing the equations for the reactions of dissolution of sulfides.
2. How many moles of gaseous products are obtained by decomposition of 10 moles of nickel(II) nitrate?
3. What nitrogen compounds are obtained by direct binding (fixation) of atmospheric nitrogen? Write the reaction equations for their preparation and indicate the conditions for the reactions.
4. What volume of ammonia (normal conditions) can be obtained by acting with two liters of 0.5 n alkali solution on the ammonium salt?
5. What mass of phosphorus (V) oxide is formed during the complete combustion of phosphine PH 3 obtained from calcium phosphide Ca 3 P 2 weighing 18.2 g?
Task 5
1. Give examples of addition, hydrogen substitution and oxidation reactions characteristic of ammonia. Write the equations for the corresponding reactions.
2. Calculate the volume (N.O.) of nitrogen dioxide required to react with 50 ml of 0.1 N potassium hydroxide solution?
3. What mass of ammonium chloride is formed during the interaction of hydrogen chloride weighing 7.3 g with ammonia weighing 5.1 g? What gas will be left in excess? Determine the mass of the excess.
4. Arrange the coefficients using the electronic balance method in the equation: Ca 3 (RO 4) 2 + SiO 2 + C → CaSiO 3 + P + CO. Determine the molar mass equivalent of the oxidizing agent and reducing agent.
5. Suggest a method by which sparingly soluble Sb(OH) 3 and Bi(OH) 3 can be separated from each other? Justify your answer by writing the equations of the corresponding reactions.
Task 6
1. How many tons of calcium cyanamide can be obtained from 3600 m 3 of nitrogen (20 ° C, normal atmospheric pressure) by interacting with calcium carbide if nitrogen losses are 40%?
2. Write the reaction equation for the interaction of bismuth with concentrated nitric acid. Arrange the coefficients in the equation using the ion-electron balance method. Determine the equivalent and molar mass of the equivalent of the reducing agent and oxidizing agent.
3. What products are obtained by calcining nitrates: sodium, calcium, copper, lead, mercury and silver? Write the equations of the corresponding reactions, arrange the coefficients using the electronic balance method.
4. Ammonium nitrate can decompose in two ways: 1) NH 4 NO 3 (c) \u003d N 2 O (g) + 2H 2 O (g); 2) NH 4 NO 3 (c.) \u003d N 2 (g) + ½O 2 (g) + 2H 2 O (g). Which of the following reactions is most likely and which is more exothermic at 25°C? Confirm your answer by calculating ∆G° 298 and ∆Н° 298 . How will the probability of these reactions change with increasing temperature?
5. What factors determine the composition of nitric acid reduction products? Justify your answer by giving the equations of the corresponding reactions.
p-elementsIV group
Topic work plan:
1. General characteristics of p-elements of group IV, finding in nature, obtaining. Physical and chemical properties of simple substances.
2. Carbon: natural compounds, production, application, physical properties, chemical properties. Allotropic modifications of carbon. Carbon monoxide (II) and metal carbides. Carbon monoxide (IV). Carbonic acid, carbonates, thiocarbonates.
3. Compounds of carbon with non-metals: cyanide, carbon disulfide; rhodanic acid and thiocyanates.
4. Silicon: natural compounds, production, application, physical and chemical properties. Oxygen compounds of silicon. Silicic acid, silicates.
5. Germanium subgroup elements: natural compounds, production, application, physical properties, chemical properties. Oxygen compounds of germanium subgroup elements: acid-base and redox properties.
Exercise 1
1. Describe the physical and chemical properties of the silicon element. Write the equations for the corresponding reactions.
2. How can the oxidizing properties of lead (IV) oxide be explained? Complete the reaction equation: PbO 2 + HCl → ... Arrange the coefficients in the equation using the electron-ion balance method. Determine the mass of salt and the volume of gas (n.o.) that are obtained as a result of the reaction of 0.2 mol PbO 2 with hydrochloric acid.
3. Make the equations for the reactions of obtaining chloride and silicon nitride and indicate the conditions for their occurrence. Why do silicon halides "smoke" in humid air? Justify your answer by writing the equations for the corresponding reactions.
4. What volume of acetylene (normal conditions) can be obtained by the interaction of water with 0.80 kg of CaC 2.
5. Prove the amphoteric character of Sn(OH) 2 . Give the equations of the corresponding reactions.
Task 2
1. Describe the physical and chemical properties of the element carbon. Write the equations for the corresponding reactions.
2. Without calculating, determine the reaction of the medium (pH = 7, pH< 7, рН >7) an aqueous solution of sodium silicate. Justify your answer by giving the equations of the corresponding reactions.
3. When burning 3.00 g of anthracite, 5.30 liters of CO 2 were obtained, measured at n.o. Calculate what percentage of carbon (by mass) anthracite contains.
4. Complete the reaction equation: C + HNO 3 (conc.) CO 2 + ... Arrange the coefficients in the equation using the electron-ion balance method. Determine the equivalent and molar mass of the equivalent of the reducing agent and oxidizing agent.
5. How many grams of NaCl can be obtained from 265 g of Na 2 CO 3.
Task 3
1. Describe the physical and chemical properties of the elements of the germanium subgroup. Write the equations for the corresponding reactions.
2. What class of compounds do Pb 2 O 3 and Pb 3 O 4 (minium) belong to? Give their graphic formulas. Write an equation for the interaction of red lead with a solution of potassium iodide in a sulfuric acid medium.
3. How many grams of CaCO 3 precipitates if an excess of soda solution is added to 400 ml of a 0.5 n CaCl 2 solution.
4. Given the values of the dissociation constants of hydrocyanic and carbonic acids: 5*10 -10, 4*10 -7, respectively, consider how atmospheric carbon dioxide affects aqueous solutions of alkaline cyanides. Why should cyanide be stored in tightly closed containers?
5. What are the acid-base properties of lead (II) oxide and hydroxide? Justify your answer by giving the equations of the corresponding reactions.
Task 4
1. Describe the physical and chemical properties of carbon monoxide (IV) and carbonic acid. Write the equations for the corresponding reactions.
2. Why does germanium not react with dilute sulfuric acid, while it dissolves in concentrated acid? Write an equation for the interaction of germanium with concentrated sulfuric acid. Arrange the coefficients in the equation using the electron-ion balance method.
3. When passing water vapor over hot coal, water gas is obtained, consisting of equal volumes of CO and H 2. What volume of water gas (normal conditions) can be obtained from 3.0 kg of coal.
4. What transformations do sodium and potassium cyanides undergo during long-term storage of their aqueous solutions? Write the equations for the corresponding reactions.
5. What color will litmus be colored in aqueous solutions of KCN, Na 2 CO 3. Justify your answer by writing the equations of the corresponding reactions.
Task 5
1. Describe the physical and chemical properties of silicon oxide (IV) and silicic acid. Write the equations for the corresponding reactions.
2. What is the difference between interactions of germanium and lead with concentrated nitric acid? Why? Write the equations for the corresponding reactions. Arrange the coefficients in the equations using the electron-ion balance method.
3. Calcium carbonate decomposes when heated into CaO and CO 2. What mass of natural limestone containing 90% (wt.) CaCO 3 will be required to obtain 7.0 tons of quicklime.
4. Complete the reaction equation: PbS + HNO 3 (conc.) PbSO 4 +NO 2 + .... Arrange the coefficients in the equation using the electron-ion balance method. Determine the equivalent of the oxidizing agent and reducing agent.
5. Determine the pH of 0.02 N soda solution Na 2 CO 3, taking into account only the first stage of hydrolysis.
Task 6
1. Give the electronic formulas of tin in the oxidation states (+2) and (+4). What properties (oxidizing or reducing) can tin compounds exhibit in these oxidation states? Justify your answer by writing the equations for the corresponding reactions.
2. When 0.5 g of limestone is dissolved in hydrochloric acid, 75 ml of carbon dioxide (n.o.) is obtained. Calculate the percentage of calcium carbonate in limestone.
3. Calculate the weight loss (in percent) that occurs when the sodium bicarbonate is ignited.
4. Compare the degree of hydrolysis of the salt and the pH of the medium in 0.1 M and 0.001 M solutions of potassium cyanide. Justify your answer by doing the appropriate calculations.
5. Complete the reaction equation: SnCl 2 + HgCl 2 Hg 2 Cl 2 + ... Arrange the coefficients in the equation using the electron balance method. Determine the equivalent, calculate the molar mass of the oxidizing agent and reducing agent.
general characteristicsd-elements
d-elements include elements in whose atoms the d-sublevel of the preexternal energy level is filled. They are also called transitional, located in the periodic system in large periods in side subgroups of all groups between s- and p-elements. The general electronic formula of valence electrons of atoms of d-elements (n-1)d 1-10 ns 2, where n is the principal quantum number, i.e. valence electrons are at different energy levels, so d-elements are located in side subgroups.
At the outer level, d-elements have 1-2 electrons (n s-state), the remaining valence electrons are located on the (n-1) d sublevel (pre-outer layer). This structure of the electron shells of atoms of d-elements determines a number of their general properties:
1. All d-elements are metals, which are characterized by high hardness, refractoriness, and significant electrical conductivity.
2. For each decade of d-elements, the most stable electronic configurations are: d 0 ,d 5 ,d 10 .
: (so, Sc, Y, La, unlike other d-elements, exhibit a constant oxidation state of +3) (n-1) d 1 ns 2
: (Mn, Fe, Re) - (n-1)d 5 ns 2
electron slip 24 Cr: …3d 4 4s 2 →…3d 5 4s 1 .
: (Zn, Cd, Hg) – (n-1)d 10 ns 2
electron slip: 29 Cu: …3d 10 4s 1 ; 47 Ag:…4d 10 5s 1 ; 79 Au:…5d 10 6s 1 ; 46 Pd:…4d 10 5s 0 .
3. The increased stability of unfilled, half-filled, and fully filled d-shells determines the most characteristic oxidation states of these elements and the stability of their compounds. Thus, the Fe 3+ (d 5), Zn 2+ (d 10) compounds are stable, while the Cr 2+ and Mn 3+ compounds with the d 4 configuration are unstable.
4. In the formation of compounds, s-electrons and part or all of the d-electrons are used. Moreover, at first, s-electrons take part in the formation of bonds, and then - d-electrons. The exceptions are the elements of the Zn subgroup, in whose atoms there are no unpaired d-electrons - [(n-1)d 10 ns 2 ] and Pd - (4d 10 5s 0), whose atom in the unexcited state has no external s-electrons. In this regard, the features of d-elements are:
– a large set of valence states;
- wide ranges of changes in the redox and acid-base properties of their compounds.
5. In each subgroup, the properties of the first elements (elements of period IV) differ markedly from the properties of the remaining elements. The similarity of the elements of periods V and VI is due to lanthanide compression.
6. Unlike p-elements, d-elements do not show negative oxidation states. They do not form gaseous compounds with hydrogen. If p-elements in the group from top to bottom tend to show the highest degree of oxidation, then for d-elements, on the contrary, this tendency increases. An increase in the stability of higher oxidation states is due to the fact that all valence electrons in heavy atoms are located at a large distance from the nucleus and are more effectively shielded from it. So, for the d-elements of group VI Mo and W, the oxidation state is +6, while Cr is stable in compounds where its oxidation state is +3. The consequence of this 
is a decrease in the oxidative capacity of compounds in the highest oxidation state of d-elements in a group from top to bottom.
increased stability,
weakening of oxidizing properties is observed.
So, for example, oxide Mn (VII) is unstable and decomposes with an explosion: 2Mn 2 O 7 \u003d 4MnO 2 + 3O 2,
while the corresponding oxides of technetium and rhenium are stable crystalline substances. For the same reason, Mn and Re interact differently with nitric acid:
Mn + 4HNO 3 \u003d 4Mn (NO 3) 2 + 2NO 2 + 4H 2 O,
Re + 7HNO 3 = HReO 4 + 7NO 2 + 3H 2 O
7. The acid-base properties of hydroxides of d-elements depend on their degree of oxidation: with an increase in the degree of oxidation, the chemical properties of hydroxides change from basic through amphoteric to acidic. For instance:
Fe (OH) 2 Fe (OH) 3 H 2 FeO 4
Cr(OH) 2 Cr(OH) 3 H 2 CrO 4
basic amphoteric acid
MnO Mn 2 O 3 MnO 2 MnO 3 Mn 2 O 7
basic amphoteric acid
8. In a group from top to bottom, the acidic properties of hydroxides, when manifested by elements of the same oxidation state, fall. For example: H 2 MnO 4 -H 2 TcO 4 -H 2 ReO 4
weakening of acidic properties
9. For d-elements, the formation of various coordination compounds is characteristic (especially 4d- and 5d-elements). Most d-element compounds are colored.
10. d-elements are good catalysts and are used in many catalytic processes.
d-elementsVI,VII,VIII groups
Topic work plan:
1. d-elements of group VIII. Iron family: natural compounds, production, application, physical properties, chemical properties.
2. Oxygen compounds of elements of the iron subgroup: acid-base and redox properties.
3. Complex compounds of elements of the iron subgroup.
4. d-elements of the chromium subgroup: natural compounds, production, application, physical properties, chemical properties.
5. Oxygen compounds of elements of the chromium subgroup: acid-base and redox properties.
6. d-elements of the manganese subgroup: natural compounds, production, application, physical properties, chemical properties.
7. Oxygen compounds of manganese subgroup elements: acid-base and redox properties.
Exercise 1
1. Describe the physical properties of the elements of the iron family.
2. Determine what mass of lead dioxide can be reduced by 0.15 l of a 0.2 N solution of potassium chromite in an alkaline medium.
3. Determine how much volume the nickel tetracarbonyl will take, formed in accordance with the chemical reaction equation: Ni (tv) + 4CO (g) \u003d (g), if 23.48 kg of nickel entered the reaction, and production losses were 10%?
4. Complete the chemical reaction equation: KMnO 4 + HBr = Br 2 + ... Arrange the coefficients in the equation using the electron-ion balance method. Determine the equivalent and molar mass of the oxidizing agent and reducing agent.
5. In what two ways can nickel (II) chloride be obtained from metallic nickel? Write the equations for the corresponding reactions.
Task 2
1. Describe the chemical properties of the elements of the iron family, compare their chemical activity. Give the equations of the corresponding reactions.
2. An alloy of copper and nickel weighing 1.5 g was affected by an excess of hydrochloric acid solution. This collected a gas volume of 114 ml (N.O.). Calculate the mass fraction of metals in the mixture.
3. Compose molecular and ion-molecular equations for the formation of nickel (II) hydroxide and its dissolution in nitric acid.
4. Complete the chemical reaction equation: H 2 O 2 + K 2 Cr 2 O 7 + HCl \u003d O 2 + ... Arrange the coefficients in the equation using the electron-ion balance method.
5. Write the reaction equations for the production of cobalt (II) hydroxide and its oxidation with atmospheric oxygen.
Task 3
1. d-elements of the iron family: natural compounds, production, application.
2. How can iron (III) chloride be obtained from iron (II) chloride, and vice versa? Write equations for the corresponding reactions.
3. The most common ore from which chromium is obtained is chromium iron ore FeCr 2 O 4 . Calculate the percentage of impurities contained in the ore if 240 kg of ferrochrome (an alloy of iron and chromium) containing 65% chromium was obtained from 1 ton of it during smelting.
4. Complete the chemical reaction equation: KMnO 4 + KBr + H 2 SO 4 \u003d Br 2 + ... Arrange the coefficients in the equation using the electron-ion balance method. Determine the equivalent and molar mass of the oxidizing agent and reducing agent.
5. In natural waters, iron is present mainly in the form of bicarbonate, which, under the action of water and atmospheric oxygen, gradually turns into iron (III) hydroxide. Write an equation for this reaction, indicate which element donates electrons and which one adds them. Arrange the coefficients in the equation using the electron-ion balance method
Task 4
1. Oxygen compounds of iron: characterize their acid-base and redox properties.
2. What volume of chlorine (n.o.) will be released when 1 mole of potassium dichromate reacts with an excess of hydrochloric acid?
3. Indicate the characteristic valence states of the Ni atom. Which of them are sustainable? Write the formulas for oxides and hydroxides of nickel. Give a brief description of the acid-base properties of these compounds. Give the equations of the corresponding reactions.
4. Iron pentacarbonyl decomposes in the light according to the reaction equation: 2=+CO. Calculate how much of the substance decomposed if this formed 5.6 liters of carbon monoxide (II) (n.o.).
5. Complete the chemical reaction equation: PbO 2 + MnSO 4 + HNO 3 = PbSO 4 + Pb(NO 3) 2 + ... Arrange the coefficients in the equation using the electron-ion balance method.
Task 5
1. Describe the ratio of elements of the iron family to air, water, acids. How does the chemical activity of elements change in the series: Fe → Co → Ni? Why? Give the equations of the corresponding reactions.
2. Write the equations of chemical reactions with which you can carry out the following transformations: Co 2 O 3 → Co → Co(NO 3) 2 ®Co(OH) 2 → Co(OH) 3 → CoCl 2 → CoCl 3.
3. Specify the characteristic valence states of the Fe atom. Which of them are sustainable? Write the formulas for oxides and hydroxides of iron. Give a brief description of the acid-base properties of these compounds. Give the equations of the corresponding reactions.
4. What volume of a sulfuric acid solution with a mass fraction of H 2 SO 4 20% (p = 1.143 g / ml) should be taken to dissolve iron, the mass fraction of impurities in which is 12.5%?
5. Complete the chemical reaction equation: K 2 Cr 2 O 7 + SO 2 + H 2 SO 4 = K 2 SO 4 + ... Arrange the coefficients in the equation using the electron-ion balance method.
Determine the equivalent and molar mass of the oxidizing agent and reducing agent.
Task 6
1. d-elements of the chromium subgroup: natural compounds, production, application.
2. Iron filings weighing 16.8 g were burned in an atmosphere of chlorine. The resulting product was dissolved in 400 ml of water. Determine the mass fraction (%) of the solute in the resulting solution.
3. Write the equations of chemical reactions with which you can carry out the following transformations: NiO → Ni → Ni(NO 3) 2 → Ni(NO 3) 3 → NiCl 2 .
4. Indicate the characteristic valence states of the Co atom. Which of them are sustainable? Write the formulas for oxides and hydroxides of cobalt. Give a brief description of the acid-base properties of these compounds. Give the equations of the corresponding reactions.
5. Complete the chemical reaction equation: Na 2 SO 3 + KMnO 4 + H 2 SO 4 = Na 2 SO 4 + ... Arrange the coefficients in the equation using the electron-ion balance method. Determine the equivalent and molar mass of the oxidizing agent and reducing agent.
general characteristicss-elements
The s-elements include elements of the main subgroups of groups I and II (IA and IIA - subgroups) of the periodic system. The general electronic formula of the valence layer of s-elements is ns 1-2, where n is the main quantum number.
Elements IA - subgroups Li, Na, K, Rb, Cs, and Fr - are called alkali metals, and for elements IIA - subgroups - Be, Mg, Ca, Sr, Ba, Ra - the last four elements are called alkaline earth metals.
Alkali metal atoms for the formation of chemical bonds have only one electron located in ns - atomic orbital (AO). A relatively small value of the ionization energy decreases from Li (I = 520 kJ/mol) to Cs (I = 342 kJ/mol), which facilitates the detachment of an electron from the AO. Therefore, alkali metal atoms in various chemical reactions are easily converted into singly charged cations with a stable eight-electron (n-1)s 2 (n-1)p 6 configuration of the corresponding noble gas. For example: K (4s 1) - e \u003d K + ().
Thus, in their many ionic compounds, alkali metals have only one oxidation state (+1).
Elements of the IIA - subgroup already contain two electrons at the external energy level, capable of separation before the formation of ionic chemical bonds with the transition of one of them to np AO: ns 2 → ns 1 np 1. The oxidation state of elements IIA - subgroups in their various compounds is (+2).
Beryllium in its physicochemical properties stands out sharply among the IIA - subgroup. The atoms of this element have the highest value of the first ionization energy among all s-elements (I=901 kJ/mol) and the greatest difference in ns and np-AO. Therefore, beryllium with other elements forms predominantly covalent chemical bonds, which are usually considered from the standpoint of the method of valence bonds. The atomic orbitals of beryllium undergo sp-hybridization, which corresponds to the formation of linear molecules BeCl 2, BeI 2, etc. Beryllium (+ II) is characterized by a tendency to form complex compounds:
Be(OH) 2 + 2OH - → 2-
BeCl 2 + 2Cl - → 2-
Oxides and hydroxides of s-elements have basic properties. Among all s-elements, only Be, its oxide and hydroxide exhibit amphoteric properties.
The chemical behavior of Li and Mg, as well as Be and Al, is largely similar due to the diagonal periodicity.
Alkali metals with oxygen form not only oxides Me 2 [O], but also compounds of the Me 2 type - peroxides; Me, superoxides; Me - ozonides. The oxidation state of oxygen in these compounds is, respectively, –1; -1/2; -1/3.
Peroxides of alkaline earth metals are known. Of these, barium peroxide BaO 2 has the greatest practical value.
Of interest are also compounds of s-elements with hydrogen - hydrides, in which hydrogen has an oxidation state of -1.
Topic work plan:
1. General characteristics of s-elements of groups I and II of the periodic system D.I. Mendeleev.
2. Properties of simple substances.
3. Finding in nature and obtaining simple substances.
4. The most important compounds of s-elements: oxides, peroxides, hydroxides, salts.
Exercise 1
1. What chemical properties of alkali metals characterize them as the most typical metals? Justify your answer by giving the equations of the corresponding reactions.
2. At 25 0 C, the solubility of NaCl is 36.0 g in 100 g of water. Find the mass fraction of NaCl in the saturated solution.
3. Determine the percentage of impurities in technical calcium carbide, if the complete decomposition of 1.8 kg of the sample with water produced 560 liters of acetylene (n.o.).
4. What s-elements of group II are complete electronic analogues? Why?
5. What amount of calcium hydroxide should be added to 162 g of a 5% solution of calcium bicarbonate to obtain an average salt?
Task 2
1. Describe the properties of oxides of group I s-elements. Give ways to get them. Write the equations for the corresponding reactions.
- sodium dihydrogen phosphate and caustic potash;
− calcium carbonate and hydrochloric acid;
− tin (II) hydroxide and caustic soda.
3. Write the equations of chemical reactions, as a result of which the following transformations can be carried out: Be → BeCl 2 → Be(OH) 2 → Na 2 → BeSO 4.
4. Complete the equation for the following chemical reaction: BaO 2 + Cr 2 (SO 4) 3 + NaOH → .... Arrange the coefficients in the equation using the electron-ion balance method. Calculate the equivalent of the oxidizing agent. Give the names of the initial substances and reaction products in accordance with the international nomenclature.
5. The density of a 26% KOH solution is 1.24 g/ml. How many moles of KOH equivalent are there in 5 liters of solution?
Task 3
1. Describe the properties of oxides of group II s-elements. Give ways to get them. Write the equations for the corresponding reactions.
2. What substances are formed during the combustion of calcium in air? Why, when the resulting product is wetted with water, a significant amount of heat is released and the smell of ammonia is felt. Justify your answer by writing the equations for the corresponding reactions.
3. What volume of SO 2 (at n.o.) can be obtained by treating a solution of potassium sulfite with a 0.085 N solution of sulfuric acid with a volume of 0.05 l?
4. Determine the type of chemical bond between atoms in the CaCl 2 molecule. What is the geometric shape of a molecule? Are the bonds in the molecule polar, is the molecule polar?
5. Why cannot alkali metals be used to restore substances dissolved in water? Justify the answer.
Task 4
1. Describe the properties of hydroxides of group I s-elements. Give ways to get them. Write the equations for the corresponding reactions.
2. Why is sodium chlorite solution neutral and sodium hypochlorite alkaline? Justify your answer by writing the equations for the corresponding reactions.
3. To prepare a 5% solution of MgSO 4, 400 g of MgSO 4 * 7H 2 O were taken. Find the mass of the resulting solution.
4. What volume of 0.25 n H 2 SO 4 can be neutralized by adding 0.6 l of 0.15 n Ca (OH) 2? Justify your answer with appropriate calculations.
5. 25 g of baking soda was calcined, the residue was dissolved in 200 g of water. Calculate the mass fraction of salt in the solution.
Task 5
1. Describe the properties of hydroxides of group II s-elements. Give ways to get them. Write the equations for the corresponding reactions.
2. Compose molecular and ion-molecular equations for reactions occurring in solutions between the following substances:
Potassium hydrogen phosphate and caustic soda;
Calcium bicarbonate and carbon monoxide (IV);
Lead hydroxide (II) and caustic potash.
3. What reaction is the production of alkali metal hydrides based on? Write the reaction equations for the hydrolysis of sodium hydride and the electrolysis of a melt of lithium hydride.
4. To dissolve 4 g of the oxide of a divalent element, 25 g of 29.2% hydrochloric acid was required. Determine the oxide of which element was taken?
5. How can barium hydride and nitride be obtained? Write the reaction equations for the interaction of these compounds with water.
Task 6
1. Sodium oxide and peroxide. Preparation, physical and chemical properties. Write the equations for the corresponding reactions.
2. Why does magnesium dissolve well in water containing ammonium salts? Justify your answer by writing the equations for the corresponding reactions.
3. One of the industrial methods for obtaining potassium is the interaction of molten KOH with liquid sodium (440˚С): Na + KOH → NaOH + K. Prove that the above reaction is possible.
4. How many grams of CaCO 3 precipitates if an excess of soda solution is added to 400 ml of a 0.5 n CaCl 2 solution?
5. Complete the equation for the following chemical reaction: BaO 2 + FeSO 4 + H 2 SO 4 → .... Arrange the coefficients using the electron-ion balance method. Calculate the molar mass of the oxidizing agent equivalent. Give the names of the initial substances and reaction products in accordance with the international nomenclature.
