Chemistry- the science of substances, the laws of their transformations (physical and chemical properties) and application.
Currently, more than 100 thousand inorganic and more than 4 million are known. organic compounds.
Chemical phenomena: some substances are transformed into others that differ from the original ones in composition and properties, while the composition of the atomic nuclei does not change.
Physical phenomena: varies physical state substances (vaporization, melting, electrical conductivity, radiation of heat and light, malleability, etc.) or new substances are formed with a change in the composition of atomic nuclei.
Atomic-molecular science.
1. All substances are made up of molecules.
Molecule - the smallest particle of a substance that has its chemical properties.
2. Molecules are made up of atoms.
Atom - the smallest particle of a chemical element that retains all its chemical properties. Different elements have different atoms.
3. Molecules and atoms are in continuous motion; there are forces of attraction and repulsion between them.
Chemical element - this is a type of atoms characterized by certain nuclear charges and the structure of electronic shells. Currently, 118 elements are known: 89 of them are found in nature (on Earth), the rest are obtained artificially. Atoms exist in a free state, in compounds with atoms of the same or other elements, forming molecules. The ability of atoms to interact with other atoms and form chemical compounds is determined by its structure. Atoms consist of a positively charged nucleus and negatively charged electrons moving around it, forming an electrically neutral system that obeys the laws characteristic of microsystems.
Atomic nucleus - the central part of the atom, consisting of Zprotons and N neutrons, in which the bulk of the atoms are concentrated.
Core charge - positive, equal in value to the number of protons in the nucleus or electrons in a neutral atom and coincides with the atomic number of the element in the periodic table.
The sum of the protons and neutrons of an atomic nucleus is called the mass number A = Z+N.
Isotopes
- chemical elements with identical nuclear charges, but different mass numbers due to different numbers of neutrons in the nucleus.
|
Mass |
A |
63 |
Cu and |
65 |
35 |
Cl and |
37 |
Chemical formula - this is a conventional notation of the composition of a substance using chemical symbols (proposed in 1814 by J. Berzelius) and indices (index is the number at the bottom right of the symbol. Indicates the number of atoms in the molecule). The chemical formula shows which atoms of which elements and in what ratio are connected to each other in a molecule.
Allotropy - the phenomenon of the formation by a chemical element of several simple substances that differ in structure and properties. Simple substances - molecules, consist of atoms of the same element.
Cfalse substances - molecules are made up of different atoms chemical elements.
Atomic mass constant equal to 1/12 of the mass of isotope 12 C - the main isotope of natural carbon.
m u = 1 / 12 m (12 C ) =1 a.u.m = 1.66057 10 -24 g
Relative atomic mass (A r) - dimensionless quantity equal to the ratio of the average mass of an atom of an element (taking into account the percentage of isotopes in nature) to 1/12 of the mass of an atom 12 C.
Average absolute atomic mass (m) equal to the relative atomic mass times the amu.
Ar(Mg) = 24.312
m(Mg) = 24.312 1.66057 10 -24 = 4.037 10 -23 g
Relative molecular weight (M r) - a dimensionless quantity showing how many times the mass of a molecule of a given substance is greater than 1/12 the mass of a carbon atom 12 C.
M g = m g / (1/12 m a (12 C))
m r - mass of a molecule of a given substance;
m a (12 C) - mass of a carbon atom 12 C.
M g = S A g (e). The relative molecular mass of a substance is equal to the sum of the relative atomic masses of all elements, taking into account the indices.
Examples.
M g (B 2 O 3) = 2 A r (B) + 3 A r (O) = 2 11 + 3 16 = 70
M g (KAl(SO 4) 2) = 1 A r (K) + 1 A r (Al) + 1 2 A r (S) + 2 4 A r (O) =
= 1 39 + 1 27 + 1 2 32 + 2 4
16 = 258
Absolute molecular mass equal to the relative molecular mass multiplied by the amu. The number of atoms and molecules in ordinary samples of substances is very large, therefore, when characterizing the amount of a substance, a special unit of measurement is used - the mole.
Amount of substance, mol . Means a certain number of structural elements (molecules, atoms, ions). Designatedn , measured in moles. A mole is the amount of a substance containing as many particles as there are atoms in 12 g of carbon.
Avogadro's number (N A ). The number of particles in 1 mole of any substance is the same and equals 6.02 10 23. (Avogadro's constant has the dimension - mol -1).
Example.
How many molecules are there in 6.4 g of sulfur?
The molecular weight of sulfur is 32 g/mol. We determine the amount of g/mol of substance in 6.4 g of sulfur:
n (s) = m(s)/M(s ) = 6.4 g / 32 g/mol = 0.2 mol
Let's determine the number of structural units (molecules) using the constant Avogadro N A
N(s) = n (s)N A = 0.2 6.02 10 23 = 1.2 10 23
Molar mass shows the mass of 1 mole of a substance (denotedM).
M = m / n
The molar mass of a substance is equal to the ratio of the mass of the substance to the corresponding amount of the substance.
The molar mass of a substance is numerically equal to its relative molecular mass, however, the first quantity has the dimension g/mol, and the second is dimensionless.
M = N A m (1 molecule) = N A M g 1 amu = (N A 1 amu) M g = M g
This means that if the mass of a certain molecule is, for example, 80 amu. ( SO 3 ), then the mass of one mole of molecules is equal to 80 g. Avogadro’s constant is a proportionality coefficient that ensures the transition from molecular relationships to molar ones. All statements regarding molecules remain valid for moles (with replacement, if necessary, of amu by g). For example, the reaction equation: 2 Na + Cl 2 2 NaCl , means that two sodium atoms react with one chlorine molecule or, which is the same thing, two moles of sodium react with one mole of chlorine.
M.: Higher School, 2002. - 415 p.
The manual is intended for schoolchildren, applicants and teachers. The manual outlines in a brief but informative and clear form modern basics chemistry. These are the basics that every high school graduate must understand and absolutely must know for anyone who sees himself as a chemistry, medical, or biologist student of the 21st century.
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TABLE OF CONTENTS
Preface ...................... 3
PART 1. THEORETICAL CHEMISTRY ............ 5
CHAPTER 1. Basic concepts and laws of chemistry.......... 5
§ 1.1. Chemistry subject.................. 5
§1.2. Atomic-molecular theory............ 7
§ 1.3. Law of conservation of mass and energy......... 10
§ 1.4. Periodic law............... 12
§ 1.5. Basic concepts of chemistry......................... 14
§ 1.6. Stoichiometric ratios in chemistry........ 18
§ 1.7. Gas laws........................ 19
CHAPTER 2. Atomic structure................. 22
§ 2.1. Development of ideas about complex structure atom.... 22
§ 2.2. Quantum numbers of electrons............ 25
§ 2.3. Distribution of electrons in atoms.......... 28
§ 2.4. Radioactive transformations............ 33
§ 2.5. Periodicity of properties of atoms of elements....... 37
CHAPTER 3. Chemical bonding and molecular structure......... 41
§ 3.1. The nature of the chemical bond............... 41
§ 3.2. Covalent bond................... 44
§ 3.3. Ionic bonding................... 48
§ 3.4. Metal bond............... 50
§ 3.5. Intermolecular chemical bonds......... 51
§ 3.6. Valency and oxidation state........... 55
§ 3.7. Spatial structure of molecules......... 58
CHAPTER 4. States of matter................ 63
§ 4.1. Characteristic properties of gases, liquids and solids... 63
§ 4.2. Phase diagrams of substances............ 66
§ 4.3. Gases...................... 68
§ 4.4. Liquids................... 70
§ 4.5. Crystalline substances............... 73
§ 4.6. Various shapes existence of substances...... 80
CHAPTER 5. Energy effects of chemical reactions...... 81
§ 5.1. Release and absorption of energy in chemical reactions... 81
§ 5.2. Exothermic and endothermic reactions. Thermochemical
Hess's law................... 87
CHAPTER 6. Kinetics of chemical reactions............ 93
§ 6.1. Basic concepts and postulates of chemical kinetics... 93
§ 6.2. The influence of temperature on the reaction rate........ 97
§ 6.3. Catalysis......................... 99
CHAPTER 7. Chemical equilibrium............... 103
§ 7.1. Determination of the state of equilibrium......... 103
§ 7.2. Chemical equilibrium constant.......... 105
§ 7.3. Shift in chemical equilibrium. Le Chatelier's principle.. 108
§ 7.4. On optimal conditions for obtaining substances in industrial
scale................... 111
CHAPTER 8. Solutions.................... 114
§ 8.1. Dissolution as a physical and chemical process....... 114
§ 8.2. Factors affecting the solubility of substances...... 117
§ 8.3. Methods of expressing the concentration of solutions...... 121
CHAPTER 9. Electrolytic dissociation and ionic reactions in solutions . 122
§ 9.1. Electrolytes and electrolytic dissociation...... 122
§ 9.2. Degree of dissociation. Strong and weak electrolytes. Dissociation constant....... 123
§ 9.3. Ionic reaction equations......................... 126
§ 9.4. Hydrolysis of salts................... 128
CHAPTER 10. Basic types of chemical reactions.......... 129
§ 10.1. Symbolism and classification characteristics of reactions.... 129
§ 10.2. Classification by the number and composition of reagents and reaction products.................................. 131
§ 10.3. Classification of reactions according to phase characteristics..... 136
§ 10.4. Classification of reactions according to the type of particles transferred... 137
§ 10.5. Reversible and irreversible chemical reactions..... 138
CHAPTER 11. Redox processes....... 140
§ 11.1. Redox reactions....... 140
§ 11.2. Selection of stoichiometric coefficients in OVR.... 144
§ 11.3. Standard OVR potentials.......... 148
§ 11.4. Electrolysis of solutions and melts of electrolytes..... 152
PART II. INORGANIC CHEMISTRY........... 154
CHAPTER 12. General characteristics of inorganic compounds, their classification and nomenclature.154
§ 12.1. Oxides......................... 155
§ 12.2. Bases (metal hydroxides) ......... 158
§ 12.3. Acids................... 160
§ 12.4. Salts........................ 165
CHAPTER 13. Hydrogen.................... 168
§ 13.1. Atomic structure and position in the periodic table D.I.
Mendeleev................... 168
§ 13.2. Chemical properties of hydrogen........... 171
§ 13.3. Production of hydrogen and its use........ 173
§ 13.4. Hydrogen oxides............... 174
CHAPTER 14. Halogens................... 178
§ 14.1. Physical properties of halogens........... 178
§ 14.2. Chemical properties and production of halogens....... 180
§ 14.3. Hydrogen halides, hydrohalic acids and their salts 185
§ 14.4. Oxygen-containing halogen compounds...... 187
CHAPTER 15. Chalcogens .................. 190
§ 15.1. General characteristics............... 190
§ 15.2. Simple substances............ 191
§ 15.3. Sulfur compounds................ 196
CHAPTER 16. Nitrogen subgroup................. 204
§ 16.1. General characteristics......................... 204
§ 16.2. Properties of simple substances......................... 205
§ 16.3. Ammonia. Phosphine. Phosphorus halides........ 207
§ 16.4. Nitrogen oxides. Nitric and nitrous acids...... 210
§ 16.5. Phosphorus oxides and acids............ 214
CHAPTER 17. Carbon subgroup............... 218
§ 17.1. General characteristics......................... 218
§ 17.2. Carbon......................... 219
§ 17.3. Carbon oxides................... 223
§ 17.4. Carbonic acid and its salts............ 226
§ 17.5. Silicon................... 228
§ 17.6. Silicon compounds with oxidation state +4..... 230
§ 17.7. Silicon compounds with oxidation state -4..... 233
CHAPTER 18. Properties s-metals and their compounds .......... 234
§ 18.1. General characteristics......................... 234
§ 18.2. Chemical properties of metals.......... 236
§ 18.3. Connections s-metals............... 239
CHAPTER 19. Aluminum and boron.................. 240
§ 19.1. General characteristics............... 240
§ 19.2. Properties and preparation of simple substances........ 242
§ 19.3. Compounds of boron and aluminum............ 247
CHAPTER 20. Major transition metals............ 249
§ 20.1. General characteristics......................... 249
§ 20.2. Chromium and its compounds......................... 251
§ 20.3. Manganese and its compounds......................... 253
§ 20.4. Iron triad................... 255
§ 20.5. Iron and steel production............ 258
§ 20.6. Copper and its compounds................... 261
§ 20.7. Zinc and its compounds......................... 263
§ 20.8. Silver and its compounds......................... 264
CHAPTER 21. Noble gases ................ 265
§ 21.1. General characteristics......................... 265
§ 21.2. Chemical compounds of noble gases....... 267
§ 21.3. Application of noble gases........... 269
PART III. ORGANIC CHEMISTRY............ 271
CHAPTER 22. Basic concepts and principles in organic chemistry.. 271
§ 22.1. Subject of organic chemistry............ 271
§ 22.2. Classification of organic compounds........ 272
§ 22.3. Nomenclature of organic compounds........ 274
§ 22.4. Isomerism of organic compounds......... 278
§ 22.5. Electronic effects and reactivity of organic compounds....... 279
§ 22.6. General characteristics......................... 281
CHAPTER 23. Saturated hydrocarbons............. 283
§ 23.1. Alkanes............................. 283
§ 23.2. Cycloalkanes................... 286
CHAPTER 24. Alkenes and alkadienes............... 289
§ 24.1. Alkenes......................... 289
§ 24.2. Diene hydrocarbons......................... 293
CHAPTER 25. Alkynes.................... 295
§ 25.1. General characteristics......................... 295
§ 25.2. Preparation and chemical properties.......... 296
CHAPTER 26. Arenas .................... 300
§ 26.1. General characteristics............... 300
§ 26.2. Preparation and chemical properties.......... 303
§ 26.3. Orientants (deputies) of the first and second kind.... 308
CHAPTER 27. Alcohol and phenols................. 310
§ 27.1. General characteristics......................... 310
§ 27.2. Monohydric alcohols...................311
§ 27.3. Polyhydric alcohols......................... 315
§ 27.4. Phenols......................... 316
CHAPTER 28. Aldehydes and ketones............... 321
§ 28.1. General characteristics......................... 321
§ 28.2. Methods of obtaining ............... 323
§ 28.3. Chemical properties............... 324
CHAPTER 29. Carboxylic acids............... 327
§ 29.1. Classification, nomenclature and isomerism....... 327
§ 29.2. Monobasic saturated carboxylic acids..... 334
§ 29.3. Monobasic unsaturated carboxylic acids.... 339
§ 29.4. Aromatic carboxylic acids......... 342
§ 29.5. Dibasic carboxylic acids.......... 343
CHAPTER 30. Functional derivatives of carboxylic acids..... 345
§ 30.1. Classification of functional derivatives...... 345
§ 30.2. Anhydrides of carboxylic acids........... 346
§ 30.3. Carboxylic acid halides........ 348
§ 30.4. Amides of carboxylic acids.............. 350
§ 30.5. Esters............................. 352
§ 30.6. Fats......................... 353
CHAPTER 31. Carbohydrates (sugars)................ 357
§ 31.1. Monosaccharides......................... 357
§ 31.2. Selected representatives of monosaccharides....... 363
§ 31.3. Oligosaccharides................... 366
§ 31.4. Polysaccharides................... 368
CHAPTER 32. Amines.................... 371
§ 32.1. Saturated aliphatic amines.......... 371
§ 32.2. Aniline................... 375
CHAPTER 33. Amino acids. Peptides. Squirrels............ 377
§ 33.1. Amino acids................... 377
§ 33.2. Peptides................... 381
§ 33.3. Proteins................... 383
CHAPTER 34. Nitrogen-containing heterocyclic compounds...... 387
§ 34.1. Six-membered heterocycles............ 387
§ 34.2. Compounds with a five-membered ring.......... 390
CHAPTER 35. Nucleic acids............... 393
§ 35.1. Nucleotides and nucleosides.............. 393
§ 35.2. Structure nucleic acids............ 395
§ 35.3. Biological role nucleic acids........ 398
CHAPTER 36. Synthetic high molecular weight compounds (polymers).
Copyright JSC Central Design Bureau BIBKOM & LLC Book-Service Agency Ministry Agriculture Russian Federation Federal State Budgetary Educational Institution of Higher Professional Education “Saratov State Agrarian University named after N.I. Vavilov" GENERAL CHEMISTRY short course of lectures for first-year students Direction of training 110400.62 Agronomy Profile of training Agronomy Saratov 2011 Copyright JSC Central Design Bureau BIBKOM & LLC Book-Service Agency UDC 54 BBK 24 R99 Reviewers: Doctor of Chemical Sciences, Professor of the Department of Ecology » Federal State Budgetary Educational Institution of Higher Professional Education "Saratov State Technical University" T.I. Gubina Doctor of Technical Sciences, Professor of the Department of Biotechnology and Chemistry of the Federal State Budgetary Educational Institution of Higher Professional Education "Saratov State Agrarian University" L.A. Fomenko R99 General chemistry: a short course of lectures for first-year students in the field of preparation 110400.62 “Agronomy” / Compiled by: G.E. Ryazanova // Federal State Budgetary Educational Institution of Higher Professional Education "Saratov State Agrarian University". – Saratov, 2011. – 97 p. Short course lectures on the discipline “General Chemistry” are compiled in accordance with the discipline program and are intended for students in the training direction 110400.62 “Agronomy”. A short course of lectures contains theoretical material on basic issues of general chemistry. Aimed at developing in students knowledge about the basic laws of chemical phenomena, at using this knowledge to understand the processes occurring in nature and to solve environmental problems. Question-oriented material professional competence future agricultural specialists. UDC 54 BBK 24 Ryazanova G.E., 2011 Federal State Budgetary Educational Institution of Higher Professional Education "Saratov State Agrarian University", 2011 Copyright JSC Central Design Bureau BIBKOM & LLC Book-Service Agency Introduction Chemistry is one of the natural science disciplines. She studies the structure, properties and transformations of substances as a result of chemical reactions. Modern chemistry consists of many sections, the boundaries between which are arbitrary. The basis of general (theoretical) chemistry is the atomic-molecular theory, the theory of the structure of atoms and molecules, the theory of periodicity, the theory of chemical bonds, the theory of solutions, the theory of redox reactions, the theory of complex compounds, chemical kinetics, and the thermodynamics of chemical processes. Knowledge of general chemistry is basic for the study of other chemical disciplines, as well as for the subsequent study of agricultural chemistry, soil science, plant physiology, and chemical plant protection. Chemistry is a science inextricably linked with human production activities. The quality of knowledge in general chemistry allows us to theoretically understand the problems associated with organizing effective chemicalization of agriculture. 3 Copyright OJSC Central Design Bureau BIBKOM & LLC Kniga-Service Agency Lecture 1 BASIC CONCEPTS AND BASIC LAWS OF CHEMISTRY 1.1. Objectives of studying the discipline Chemistry is one of the natural sciences, that is, the sciences about nature, about the world around us. Chemistry is a general education and not a special discipline, but it is of great importance for agricultural workers. It forms an important part of the ideological and professional baggage of any agricultural specialist. This determines the goals of studying chemistry: To obtain chemical knowledge, which is important integral part universal science and culture; Master the methods of studying chemistry. Using the example of chemical phenomena, develop intellectual abilities and the ability to think logically; Understand the importance of chemistry for agriculture and ecology, and receive professional guidance. 1.2. Importance of Chemistry for Agriculture Chemistry is of great importance for agriculture. The most important factor in the intensification of agricultural production is the chemicalization of agriculture (this term was introduced by D.N. Pryanishnikov in 1924). Chemicalization of agriculture is the use chemicals and processes to increase soil fertility, increase production efficiency and labor productivity in agriculture. It includes: The use of mineral fertilizers containing plant nutrients. These are: – The most important macroelements – nitrogen, phosphorus, potassium (NPK); – Microelements – magnesium, iron, copper, zinc, molybdenum, sulfur, boron, etc. Chemical reclamation, creating an environment favorable for plants. These are: – liming of acidic soils (CaCO3), etc.; – gypsuming of alkaline saline soils (gypsum – CaSO4∙2H2O), etc. Chemical plant protection products. These are: – pesticides (for pests and diseases); – herbicides (for weed control), etc. Plant growth regulators. Biotechnology products: microbiological fertilizers, enzymes, vitamins, etc. The effectiveness of chemicalization in agriculture depends on the farming culture. D.N. Pryanishnikov said: “The absence necessary knowledge cannot be replaced even by excess fertilizer.” Thus, if the doses and timing of fertilizer application are violated, environmental problems arise and plant metabolism is disrupted. Improper implementation of chemical reclamation worsens soil fertility. 4 Copyright JSC Central Design Bureau BIBKOM & LLC Book-Service Agency 1.3. Subject of Chemistry Studying chemistry should begin with an understanding of what chemistry is. Currently, there are several dozen definitions of the subject of chemistry proposed by different authors. They differ from each other in the accuracy of the definition and aspects of the approach to the issue. This question sometimes even becomes controversial. Let's give some definitions. Chemistry is a science that studies the processes of transformation of substances, accompanied by changes chemical composition and structures. Chemistry is the science of chemical compounds of atoms (chemical substances and their transformations. ... The presence of a chemical bond in a substance is the main criterion for whether it is chemical. (O.S. Sirotkin) Chemistry is the science that studies the processes of changing the composition and structure of matter forms of matter, the material carriers of which are atoms (V.E. Komarov) Assignment for independent work: get acquainted with the definition of the subject of chemistry by different authors (F. Engels, D. I. Mendeleev, N. L. Glinka, O. S. Zaitsev, N. N. Semenov, D. N. Knyazev, etc.) and choose the most corresponding to your understanding of this issue. DI. Mendeleev believed: “Atoms are chemical units of matter that are indecomposable by chemical means.” The material carrier of a chemical substance (its smallest particle) is an atom. 1.4. Dialectics of basic concepts and laws of chemistry Chemistry is based on experimental data obtained by many generations of scientists different countries. The basis of general (theoretical) chemistry is the atomic-molecular theory, the doctrine of periodicity, the theory of the structure of atoms and molecules, the theory of chemical bonds, the theory of solutions, the theory of redox reactions, the theory of complex compounds, chemical kinetics, and the thermodynamics of chemical processes. Atomic-molecular theory, which is common basis not only chemistry, but also all natural sciences, has been created since the 18th century and continues to develop at the present time. It is based on the basic laws of chemistry, the laws of stoichiometry (from the Greek stoicheion - element). Stoichiometry is a branch of chemistry that studies the relationship between the amount of reactants that enter into a reaction and the amount of reactants formed as a result of the reaction. The coefficients in front of the formulas of substances in the equations of chemical reactions are called stoichiometric coefficients. People's knowledge is not something frozen. They correspond to the state of science at a given stage of development and can change, since science is an open system striving for an increasingly true, deep and complete reflection of natural phenomena. An example of this is the change in ideas about the basic laws of chemistry at different stages of the development of science. The law of conservation of mass was discovered by the great Russian scientist M.V. Lomonosov (1748-1756): 5 Copyright JSC Central Design Bureau BIBKOM & LLC Kniga-Service Agency The mass of substances that entered into the reaction is equal to the mass of the substances formed as a result of the reaction. Regardless of him, this law was confirmed in 1789 by the French chemist A.-L. Lavoisier. The law of conservation of mass makes it possible to draw up reaction equations and put an equal sign between the formulas of the starting substances and reaction products: 2NaOH + H2SO4 = Na2SO4 + 2H2O This is of great importance for solutions practical issues, since it allows one to calculate required amount starting materials to obtain the desired products in industrial scale . In the twentieth century, this law was clarified. The general law of nature is the law of conservation of energy: Energy does not arise from nothing and does not disappear without a trace, but only passes from one form to another in strictly defined quantities. The creator of the theory of relativity, A. Einstein (1905), proved the relationship between mass and energy: E = mc2 Therefore, during exothermic reactions when heat is released or endothermic reactions when heat is absorbed, the mass of the reaction products will be slightly less or greater than the mass of the starting substances according to the equation: ∆E = ∆mc2 Based on this, the following formulation of the Law of Conservation of Mass is more accurate: For exothermic reactions, the mass of substances that entered into the reaction is equal to the sum of the mass of the reaction products and the mass equivalent to the energy released. For an endothermic reaction, the mass of substances that enter into the reaction is equal to the difference between the mass of the reaction products and the mass equivalent to the absorbed energy. Since in chemical reactions the change in mass caused by the release or absorption of heat is very small, it can be argued that the law of conservation of mass is satisfied with high accuracy. The content of the law of constancy of composition has also changed. Law of constancy of composition (J. Proust (1801-1808), France): Every chemical compound has a constant composition, regardless of the method of its preparation. Proust's Law is fundamental. He confirmed the existence of molecules and the indivisibility of atoms. Berthollet (France) became Proust's opponent. Berthollet argued that the composition of a substance depends on the method of its preparation. Proust was supported by the great English chemist John Dalton and Berthollet's idea was rejected. It has now been established that the law of constancy of composition is valid only for substances with a molecular structure. 6 Copyright JSC Central Design Bureau BIBKOM & LLC Kniga-Service Agency The composition of substances with a non-molecular structure depends on the method of their preparation (oxides, sulfides of transition metals, feldspar, etc.). Berthollet's idea is currently the basis of chemical materials science, which uses precisely the dependence of the composition of the structure and properties of the material on the method of its preparation. Now substances of constant composition are called daltonides, and substances of variable composition are called berthollides in honor of scientists for whom the formulation of the law of constancy of composition was the subject of debate. In 1811, Amadeo Avogadro (Italy) discovered a law that received recognition 50 years later (1860). Avogadro's Law: Equal volumes of different gases under the same conditions (p and t) contain the same number of molecules. The first corollary of Avogadro's law: One mole of any gas under normal conditions occupies 22.4 liters. The second corollary of Avogadro’s law: The molar mass of a gas is equal to the product of the relative density of one gas over another (D) by the molar mass of another gas: Mgas = 2DN 2, Mgas = 29Dair Avogadro established: - molecules of gaseous simple substances consist of two atoms (H2, O2, N2, Cl2); – 1 mole of any substance contains 6.023 1023 atoms or molecules (Avogadro’s number). The law of equivalents was formulated by W. Richter (1793) and W. Wollaston (1807) independently of each other: Substances react with each other in quantities proportional to their equivalents. The law of equivalents is based on the concept of “chemical equivalent”. A chemical equivalent is a real or conditional particle of a substance that is equivalent to one hydrogen cation in a given acid-base reaction or one electron in a given redox reaction. The equivalence factor f'eq is calculated based on the stoichiometric coefficients of a particular reaction. The equivalence factor is a number indicating what fraction of a real particle is equivalent to one hydrogen cation in a given acid-base reaction or one electron in a given redox reaction. The equivalence factor is a dimensionless quantity. The equivalence factor may be equal to one or less than one. For example: a) H2SO4 + 2KOH = K2SO4 + 2H2O; feq(H 2 SO 4) = b) H2SO4 + KOH = KHSO4 + H2O; feq(H 2 SO 4) = 1 1 2 Stoichiometric calculations are widely used to carry out processes involving chemical substances. 7 Copyright JSC Central Design Bureau BIBKOM & LLC Kniga-Service Agency Basic provisions of atomic-molecular teaching: – substances consist of molecules; – molecules consist of atoms; – molecules and atoms are in a state of continuous movement. An atom is the smallest electrically neutral particle of a chemical element. A molecule is the smallest electrically neutral particle of a chemical substance. Currently, the unit of measurement for the amount of matter is the MOL. A mole is the amount of a substance containing as many structural units as there are atoms contained in 0.012 kg of carbon 12C. A mole is an amount of substance containing 6.02 1023 (Avogadro's number) formula units (molecules or atoms). For calculations, the following ratios are used: m nM n m M M m , n where m is the mass of the substance, g; M – molar mass, g/mol; n – amount of substance, mol. From a modern point of view, the structure of chemical substances can be different: molecular (methane CH4, ammonia NH3), atomic (diamond), ionic (salts), radical (Cl, H). A substance may consist of macromolecules or be a complex combination of various particles. Modern chemistry is based on the principle of infinite qualitative diversity of the material world. Substances with different structures obey different laws. 1.5. Methods for studying chemistry A deep understanding of chemical phenomena is possible only from the standpoint of philosophy, when applying its theories and categories. Philosophy is the love of wisdom (philosophia - Greek). The most important theory of philosophy is dialectics. Dialectics - Categories of philosophy, the doctrine of the most general laws of Matter, movement, contradiction, nature, society and thinking, quantity and quality, cause and effect Matter is a philosophical category to designate objective reality. Types of matter Matter is a type of matter that is characterized by a rest mass m ≠ 0; A chemical substance is a type of substance whose material carrier is atoms; Field – a type of matter that has zero rest mass (magnetic, electromagnetic, gravitational field) m = 0; Vacuum is a special state of matter. There are no particles in it (emptiness), but “virtual particles” arise from short-term energy fluctuations. In strong fields, real particles appear. Plasma is a special state of matter at very high temperatures(>7000). A kind of gas is formed from elementary particles, nuclei, ions. 8 Copyright JSC Central Design Bureau BIBKOM & LLC Kniga-Service Agency BEC (gaseous Bose-Einstein condensate) is a new state of matter at ultra-low temperatures. Special properties: light slows down to V = 17 m/sec, as near a rotating black hole. Figure 1.1. Forms of motion of matter (f.m.) Figure 1.2. Types of movement Drawing. 1.3. Evolution of inorganic matter 9 Copyright JSC Central Design Bureau BIBKOM & LLC Kniga-Service Agency Cause-and-effect relationships in chemical phenomena The principle of determinism: real phenomena arise, develop and are destroyed naturally, as a result of the action of certain causes. Figure 1.4. Simple scheme causality Figure 1.5. Cause-and-effect relationships in the properties of some hydroxides Methods for studying chemical phenomena Experiment Observation Modeling chemical reactions Induction is a method of reasoning from the particular to the general. Deduction is a method of reasoning from the general to the specific. General rules should be applied to solving specific problems. Analogy is a method of cognition in which, based on the similarity of a characteristic of chemical objects, a conclusion is drawn about their similarity in other characteristics. So, knowing the formula phosphoric acid, you can write the formula of orthoarsenic acid. Analysis is the separation of a chemical substance or chemical phenomenon into its component parts (properties, characteristics). 10 Copyright OJSC Central Design Bureau BIBKOM & LLC Kniga-Service Agency Synthesis is the combination of aspects, properties, characteristics of a chemical substance or chemical phenomenon into a single whole. A concept is a thought about an object or its characteristics. A judgment is the operation of concepts. Inference is the operation of judgments. Scheme 1.1. Logical forms of thinking Let's try to apply the acquired knowledge to solve a professional issue. To do this, we apply: – concepts about the properties of chemical substances; – method of deduction; – “cause-effect” categories; – logical forms of thinking: concept → judgment → inference. Problem: Is it effective to use CuSO4 microfertilizer on alkaline soils? CuSO4 + 2NaOH = Cu(OH)2↓ + Na2SO4 soluble insoluble 1. Concepts: CuSO4 is a soluble salt. Cu(OH)2↓ is a weak insoluble base. 2. Judgment: Only soluble substances are available for absorption by plants. On alkaline soils, soluble CuSO4 turns into insoluble Cu(OH)2↓ and becomes inaccessible to plants. 3. Conclusion: The use of CuSO4 on alkaline soils is inappropriate. Questions for self-control 1. Give a definition of chemistry that matches your understanding of this issue. 2. Compare the understanding of the law of conservation of mass in the 18th and 21st centuries. 3. Determine the type of development (progress or regression in chemical reactions): H2CO3 = CO2 + H2O CaO + CO2 = CaCO3 4. Based on the analogy method, write the formulas for arsenic and selenic acid. 5. Use deduction to explain the ineffectiveness of FeSO4 microfertilizer on alkaline soils. 6. How do you understand the words of D.N. Pryanishnikova: “The lack of necessary knowledge cannot be replaced even by an excess of fertilizers”? 11 Copyright JSC Central Design Bureau BIBKOM & LLC Book-Service Agency REFERENCES Main 1. Glinka, N.L. General chemistry / N.L. Glinka – M.: KNORUS, 2009. – 752 p. 2. Knyazev, D.A. Inorganic chemistry/D.A. Knyazev, S.N. Smarygin. – M.: Bustard, 2004. – 592 p. 3. Ryazanova, G.E. General and inorganic chemistry. Tables and diagrams / G.E. Ryazanov. – Saratov: Federal State Educational Institution of Higher Professional Education “Saratov State Agrarian University”, 2006 – 284 p. Additional 1. Egorov, V.V. Environmental chemistry. /V.V. Egorov. – St. Petersburg: Lan Publishing House, 2009. – 192 p. 2. Naydysh, V.M. Concepts modern natural science. /V.M. Naydysh. – M.: AlfaM; INFRA-M, 2004, – 622 p. 3. Sirotkin, O.S. Chemistry is in place. / O.S. Sirotkin // Chemistry and life. – 2003. - No. 5. – P. 26. 4. Mineev V.G., In defense of nitrates and phosphates. / V.G. Mineev // Chemistry and life. – 2008. No. 5. – P. 20. 5. http://www.xumuk.ru/encyclopedia/2/2994.html 12 Copyright JSC Central Design Bureau BIBKOM & LLC Book-Service Agency Lecture 2 MODERN TEACHING ABOUT THE STRUCTURE OF THE ATOM 2.1. Dialectics of ideas about the structure of the atom In classical chemistry there were ideas about two types of microparticles - atoms and molecules. Until the 19th century, science had the idea that the atom was indivisible and did not contain component parts. At the end of the 19th century, physicists discovered facts proving the complexity of the structure of the atom (the discovery of the electron, cathode rays, the phenomenon of radioactivity). In 1911 laureate Nobel Prize Ernest Rutherford, based on a series of experiments on the scattering of particles by thin metal plates proposed a nuclear (planetary) model of the structure of the atom. Rutherford's Planetary Model of the Atom An atom consists of a positively charged nucleus and negatively charged electrons that orbit around the nucleus. Contradictions of Rutherford's model 1. Why is the atom stable? According to the laws of classical physics, an electron, when moving around a nucleus, must emit energy and exhaust it in 10-8 seconds. and fall onto the core. 2. If an electron emits energy in a continuous stream, a continuous spectrum must correspond to this. However, the spectrum is lined. Why? Addressing these problems, Niels Bohr proposed a theory in 1913 that combined Rutherford's nuclear model of the atom with Planck's quantum theory of light. Bohr's main idea (very bold): The movement of an electron in an atom does not obey the laws of classical physics. Different laws apply in the macrocosm and microcosm. Bohr's postulates 1. An electron in an atom moves not in any, but in a certain stationary orbit and does not emit energy. 2. The emission of energy in portions (quanta) occurs only when an electron moves from one orbit to another. Bohr's theory could not explain some important properties of the atom; there were contradictions in it. In the 20s of the 20th century, a quantum mechanical theory of the structure of the atom was created. The theory is based on the idea of Louis de Broglie (1924): The electron has dual particle-wave properties, i.e. properties of both particles and waves simultaneously. 13 Copyright JSC Central Design Bureau BIBKOM & LLC Kniga-Service Agency An electron as a particle has mass and manifests itself as a single whole. The wave properties of an electron are manifested in the peculiarities of its movement, in the interference and diffraction of the electron. The de Broglie equation reflects this dualism: λ = h/m, where h is Planck’s constant; m – particle mass; – particle speed. E. Schrödinger (Austria, 1926) formulated an equation whose solution allows one to find the electron wave function Ψ(x, y, z), corresponding to the amplitude of the three-dimensional wave process. Schrödinger equation (Austria, 1926) ψ x2 2 ψ у2 ψ z2 2 2 8π 2 m (E V)ψ h2 0, where is the wave function; x, y, z – coordinates of three-dimensional space. The value 2 corresponds to the probability of finding an electron at various points in the circumnuclear space. Solving the Schrödinger equation allows us to find the electron orbital. Van den Broek (Holland, 1912) suggested: The charge of the nucleus of an atom of any element is numerically equal to the atomic number of the element in the periodic table. This brilliant intuitive guess was confirmed experimentally in 1913 by Moseley (England). 2.2. Modern quantum mechanical model of the structure of the atom The atom has a complex structure. It consists of a positively charged nucleus, which contains protons and neutrons, and negatively charged electrons, which move around the nucleus. An electron in an atom moves with enormous variable speed. The trajectory of his movement is unknown. The electron moves chaotically, randomly, forming an electron cloud (electron orbital). An electron orbital is a region of space around the nucleus of an atom in which electron movement is most likely. The behavior of an electron in an atom is characterized by the Heisenberg uncertainty principle: for a microparticle on the atomic scale it is impossible to simultaneously indicate the coordinates and angular momentum with equal accuracy. Quantum numbers The description of the state of an electron in an atom is determined by four quantum numbers. The main quantum number n characterizes the energy reserve of an electron at a given level and its distance from the nucleus. Corresponds to the electronic level number. Accepts the values of integers: 1, 2, 3, 4, 5, 6, 7, ... 14 Copyright JSC Central Design Bureau BIBKOM & LLC Kniga-Service Agency The secondary (orbital) quantum number l characterizes the orbital angular momentum electron. Determines the spatial shape of the electron cloud (orbital). Accepts integer values from 0 to n–1. Numeric values (0, 1, 2, 3) have letter designations(s, p, d, f). Table 2.1 – Orbital quantum numbers and shape of orbitals 0 1 2 3 Orbital shape Spherical Dumbbell Two crossed dumbbells Even more complex shape l (s) (p) (d) (f) Electrons with the same shape of electron orbitals form sublevels at a given energy level. The number of sublevels at a level corresponds to the level number. Table 2.2 – Dependence of the number of sublevels on the number of levels n (level) 1 2 3 4 l (sublevels) S s, p s, p, d s, p, d, f Magnetic quantum number m Characterizes the magnetic properties of the electron, which depend on the direction of the electron clouds in space. Accepts integer values within the magnitude of the side quantum number, both positive and negative, including 0 (zero). Table 2.3 - m Relationship between the values of secondary and magnetic quantum numbers Values m 0 (one) 1, 0, –1 (three) 2, 1, 0, –1, –2 (five) 3, 2, 1, 0, –1 , –2, –3 (seven) l 0 1 2 3 The number of values of the magnetic quantum number Nm can be calculated using the formula: Nm = 2l +1, where l is the value of the side quantum number. Table 2.4 – Number of orbitals at sublevels Sublevel s p D f Number of orbitals – (one) – – – (three) – – – – – (five) – – – – – – – (seven) 15 Copyright OJSC “CDB “BIBKOM” & Kniga-Service Agency LLC The spin quantum number characterizes the electron’s own angular momentum. Takes the values h 1 1 + and – (in units). 2 2 2 The Pauli principle An atom cannot have two electrons with the same values of all four quantum numbers. Corollary: One orbital can contain a maximum of two electrons with opposite spins. Table 2.5 – Maximum number of electrons at level n N 1 2 2 8 3 18 4 32 Nn = 2n2, where n is the level number (principal quantum number). Table 2.6 – Maximum number of electrons at the sublevel l N1 0(s) 2 where l is a side quantum number. 1(p) 6 2(d) 10 3(f) 14 N1 = 2(2l + 1), Principle of minimum energy An electron in an atom occupies a position with a minimum amount of energy at the level and sublevel; it is energetically favorable and stable for it. 1s< 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s < 4f < 5d < 6p < 7s < 5f < 6d < 7p Рисунок 2.1. Схема расположения энергетических подуровней в многоэлектронном атоме 16 Copyright ОАО «ЦКБ «БИБКОМ» & ООО «Aгентство Kнига-Cервис» Правило Хунда Устойчивому состоянию атома соответствует такое распределение электронов в пределах энергетического подуровня, при котором величина суммарного спина электронов максимальна. Это условие выполняется, если электроны заполняют все свободные орбитали сначала по одному, а затем происходит пополнение каждой орбитали вторым электроном. Пример: Рисунок 2.2. Последовательность заполнения электронами р – подуровня Правила В.М. Клечковского 1. Уровни и подуровни атомов заполняются электронами в порядке возрастания суммы главного и орбитального квантовых чисел (n+l). 2. При одном и том же значении суммы (n+l) электроны заполняют орбитали с меньшим п, но с большим l. 2.3. Электронные формулы элементов На основе принципа Паули, принципа минимальной энергии, правила Хунда, представлений об энергетических уровнях и подуровнях составляются электронные формулы элементов. Электронные формулы показывают распределение электронов на энергетических уровнях и подуровнях. При составлении электронной формулы сначала пишут цифру, соответствующую номеру энергетического уровня, на который поступает электрон. Затем пишут букву, соответствующую подуровню и справа наверху, в виде показателя, обозначают число электронов на данном подуровне. Примеры: I период (один энергетический уровень) Электронная формула: Н 1S1 Электронно-графическая формула: Водород Н №1 +1 (заряд ядра) 1ē 1-й S - элемент 1 S Степень окисления водорода: +1 (Н2О, NH3, HCl) Исключение: степень окисления водорода – 1 в соединениях с очень активными металлами (NaH, KH). 17 Copyright ОАО «ЦКБ «БИБКОМ» & ООО «Aгентство Kнига-Cервис» Водород Водородное пламя создает t = 2800o Восстановитель Источник, переносчик и концентратор энергии Получение аммиака, жидкого топлива, твердых жиров Экологически чистое топливо будущего Схема 2.1. Значение и применение водорода Элементы, у которых электроны заполняют s - подуровень, называются s– элементами. Первые два элемента любого периода являются s–элементами. Не Гелий Не №2 +2 (заряд ядра) 1S2 1 S первый уровень завершен 2å 2-й S - элемент Гелий – инертный газ. Гелий Дирижабли с гелиевым наполнением Теплоноситель ядерных реакторов В медицине: гелиевый воздух лечит бронхиальную астму, применяется для водолазов В электронике: жидкий гелий – среда для сверхпроводников Схема 2.2. Значение и применение гелия (инертный, легкий) Литий Li №3 +3 (заряд ядра) 3ē 1-й S - элемент II период (два энергетических уровня) Li 1S2 2S1 (валентный уровень) Литий – активный щелочной металл. Степень окисления: +1. Формулы соединений: Li2O → LiOH. Литий В атомной технике: теплоноситель, растворитель соединений урана Для кондиционирования воздуха применяют соли лития В медицине: для лечения психических расстройств Схема 2.3. Значение и применение лития 18 Copyright ОАО «ЦКБ «БИБКОМ» & ООО «Aгентство Kнига-Cервис» Be 1S2 2S2 Бериллий Ве №4 +4 (заряд ядра) 4ē 2-й S - элемент (валентный уровень) Бериллий – амфотерный металл. Соединения: ВеО → Ве(ОН)2 Н2ВеО2 Бериллий Материалы космической, ракетной и авиационной техники Бериллий токсичен ПДК в воздухе – 0,001 мг/м3 Для человека, животных вреден: рахит, ослабление костной ткани, отек легких Для растений безвреден Схема 2.4. Значение и применение бериллия При получении энергии извне (при нагревании) атомы переходят в возбужденное состояние. При этом электронная пара распаривается и электрон переходит в свободную орбиталь подуровня в пределах своего номера уровня. В 1S22S22р1 Электронно-графическая формула Бор (валентный уровень) №5 +5 (заряд ядра) 5ē 1-й р - элемент Бор Бороводородное топливо, военная техника, легирование стали Важнейший микроэлемент, влияет на белковый и углеводный обмен Схема 2.5. Значение и применение бора Бор – первый р – элемент второго периода. Элементы, у которых электроны заполняют р – подуровень, называются р – элементами. Последние 6 элементов любого периода (кроме первого и седьмого) являются р – элементами. 19 Copyright ОАО «ЦКБ «БИБКОМ» & ООО «Aгентство Kнига-Cервис» Свойства s – и р – элементов У s– и р– элементов электроны заполняют внешний уровень от 1 до 8. Chemical properties elements change rapidly, from metallic to non-metallic to inert gas. The number of electrons in the outer level of s– and p– elements corresponds to the number of the group in which the element is located. In s– and p– elements, the valence electrons are the electrons of the outer level. The main subgroups are formed by s– and p– elements. Rules for finding the oxidation state of an element 1. The negative oxidation state of an element (non-metal) is equal to the number of electrons it accepts to complete the outer level to 8 electrons. 2. Positive oxidation states of an element are equal (usually, there are exceptions) to the number of unpaired electrons in the normal and excited states of the atom that it gives up. Questions for self-control 1. Define the concept of “atom”. 2. What does an atom consist of? 3. What particles are included in the nucleus of an atom? 4. What are the same and what are different about the electrons of the carbon atom? 5. Formulate the main idea of Niels Bohr, Louis de Broglie. 6. What is an electron orbital? 7. How many orbitals are there in the s –, p –, d –, f – sublevels? 8. How is the principle of minimum energy manifested in the structure of atoms of elements? 9. How do you understand the Pauli principle? 10. Explain how you understand Hund's rule. REFERENCES Main 1. Glinka, N.L. General chemistry / N.L. Glinka – M.: KNORUS, 2009. – 752 p. 2. Knyazev, D.A. Inorganic chemistry/D.A. Knyazev, S.N. Smarygin – M.: Bustard, 2004. – 592 p. 3. Ryazanova, G.E. General and inorganic chemistry. Tables and diagrams / G.E. Ryazanov. – Saratov: Federal State Educational Institution of Higher Professional Education “Saratov State Agrarian University”, 2006 – 284 p. Additional 1. Klinsky, G.D. Inorganic chemistry /G.D. Klinsky, V.D. Skopintsev. – M: Publishing house MCHA, 2001. – 384 p. 2. Gelfman, M.I. Inorganic chemistry / M.I. Gelfman, V.P. Yustratov. – St. Petersburg: Publishing house “Lan”, 2009. – 528 p. 3. Inorganic chemistry (biogenic and abiogenic elements): Textbook / ed. V.V. Egorova. – St. Petersburg: Lan Publishing House, 2009. – 320 p. 4. ru.wikipedia.org 20 Copyright JSC Central Design Bureau BIBKOM & LLC Kniga-Service Agency Lecture 3 PERIODIC LAW AND PERIODIC SYSTEM D.I. MENDELEEV 3.1. Periodic law and periodic system of elements D.I. Mendeleev - the basis of modern chemistry Periodic law and periodic system D.I. Mendeleev is the greatest achievement of chemical science. The discovery of the law and the creation of the periodic system are the result of long and hard work of the great Russian scientist D.I. Mendeleev. Periodic law: The properties of simple bodies, as well as the forms and properties of compounds of elements, are periodically dependent on the atomic weights of the elements D. I. Mendeleev (1869) Mendeleev believed that “The future does not threaten the periodic law with destruction, but only promises superstructures and development.” The periodic law was discovered in the 19th century. In the twentieth century S.A. Shchukarev said: “Comprehension of the full meaning of the system is a task of bottomless depth, the solution of which a person will eternally strive for, as one of the unattainable and completely incomprehensible truths.” Periodic law (modern formulation) The properties of elements, as well as the forms and properties of compounds of elements, are periodically dependent on the magnitude of the charges of the nuclei of their atoms. The periodic system concentrates extensive chemical knowledge in an extremely concise form. The periodic table of elements should be considered in specific conditions of space and time. At ultrahigh temperatures, the atom is deprived of electrons, and high pressure changes their arrangement in the atoms. Under these conditions, atoms cease to obey the periodic law. The periodic system of elements is a graphic (tabular) expression of the periodic law. A system is a strict mathematical category meaning an ordered set. To establish connections between sets of functions, the idea of a matrix is used. A matrix is a rectangular table consisting of rows and columns. 21 Copyright JSC Central Design Bureau BIBKOM & LLC Book-Service Agency Table 3.1. Periodic system of chemical elements by D.I. Mendeleev 22 Copyright JSC Central Design Bureau BIBKOM & LLC Kniga-Service Agency The periodic system is a kind of matrix. Its elements are chemical elements. Lines are periods. Columns are groups. The periodic table reflects the fact that individual chemical elements represent an ordered set of elements, subject to internal relationships. Each chemical element is a natural step in the development of a substance, and not a random formation. Classification of chemical elements Obtaining substances with given properties Confirmation of the laws of philosophy Tools for foreseeing further research Periodic law Agrochemistry Geological exploration of minerals Space research Medicine, pharmaceutical industry Scheme 3.1. The meaning of the periodic law “Living matter in its average composition is connected with the soil and builds its cells from elements according to the same laws by which matter is built from materials or the atmosphere of the Sun and stars.” A.E. Fersman The periodic system of elements includes 118 currently discovered elements, which are located in 7 periods and 8 groups. 3.2. Structure of the Periodic Table Each chemical element has a serial number. The atomic number is equal to: the charge of the atomic nucleus; number of protons in the nucleus; the total number of electrons in a neutral atom. A period is a horizontal series of elements that begins with s-elements (alkali metals) and ends with p-elements (noble gases). During the period, the external electronic level is built up from 1 to 8 electrons. The completed outer level contains 8 electrons. The period number corresponds to the number of energy levels in the atom. Period I contains only two elements (one row). The second and third periods contain 8 elements each. Periods I, II and III are called “small periods” (one row). Periods IV, V and VI (major periods) consist of two rows. 23 Copyright JSC Central Design Bureau BIBKOM & LLC Kniga-Service Agency A group is a vertical sequence of elements with the same type of electronic configuration. The group of elements contains two subgroups: main (A) and secondary (B). The main subgroups are formed by s - and p - elements. Side subgroups form d - elements. The number of electrons (ē) at the outer energy level of elements of the main subgroups is equal to the group number. The number of electrons in the outer energy level of elements of side subgroups is 2, regardless of the group number. Periodic changes in the electronic structure of atoms of elements cause periodic changes in the chemical properties of elements, as well as the complex compounds that they form. The importance of the periodic table for solving educational problems Using the periodic table, you can characterize the chemical properties of an element. To do this, you need to use the numbers of the periodic system as a code that reveals the structure of the atom and the properties of the element. 3.3. Algorithm for general characteristics of an element atom 1. Ordinal number of the element (No.). 2. Charge of the nucleus of an atom (Z). 3. The number of protons in the nucleus of an atom (p+). 4. Relative atomic mass (Ar). 5. Number of neutrons in the nucleus of an atom (no). 6. Total number of electrons in an atom (ē). 7. Period number → number of energy levels. 8. Group number, subgroup (main or side) → number ē at the external level of the atom. 8.1. Main subgroup → number ē at the external level is equal to the group number. 8.2. Side subgroup number ē at the external level is 2, regardless of the group number (or 1). 9. Evaluate whether an element is a metal or non-metal (number ē at the external level). 10. Determine which electronic family the element belongs to (s, p or d) and which rank among the elements of this electronic family it is. 11. Create an electronic formula, highlight the valence electrons. 11.1. For an element of the main subgroup - electrons of the outer level. 11.2. For an element of the side subgroup - electrons of the outer level and d - the sublevel of the penultimate level. 12. Create electron graphic formulas (based on Hund’s rule), show the distribution of valence electrons at levels and sublevels in the normal and excited states. 13. Name the oxidation states of the element. 14. Write the formulas of the most important compounds corresponding to the oxidation states of the element (oxides, hydroxides, hydrogen compounds). 15. Characterize the properties of the most important compounds. 24 Copyright JSC Central Design Bureau BIBKOM & LLC Kniga-Service Agency 16. Reveal the significance of the element and its compounds for plants, for agriculture, for industry. Examples: Nitrogen N 1) No. 7 9) Non-metal (accepts ē, but can also give 2) +7 (nucleus charge) 3) 7р+ 10) 3rd р – element 4) Ar = 14 11) Electronic formula: 5 ) 7no 1s22s22p3 valence level 6) Total 7 ē 7) II period – two energy levels 12) Electron graphic formula: 8) V group, main subgroup – 5 ē excitation is impossible at the external level 13) Oxidation states: – 3, +1 , +2, +3, +4, +5. 14) Formulas of the most important compounds 3 1 N H3 → ammonia NH4OH → NH4Cl, NH4NO3, (NH4)SO4 ammonium hydroxide ammonium salts Ammonium mineral fertilizers Oxygen compounds: 2 2 1 2 indifferent oxides N 2 O, NO – 3 2 НNO2 → N 2 O3 → nitrous acid 4 2 brown gas N O2 – 5 2 НNO3 → N 2 O5 → nitric acid salts - nitrites – NaNO2 salts - nitrates – KNO3, Ca(NO3)2, NH4NO3 nitrate nitrogen fertilizers (saltpeter) Nitrogen is the most important macroelement. Phosphorus Р Р 1s22s22p63s23р3_ valence level Electron-graphic formula of the nucleus) nt levels Oxidation states: -3, +3, +5 25 Copyright JSC Central Design Bureau BIBKOM & LLC Book-Service Agency The most important phosphorus compounds 3 1 3 2 3 RN 3; P2 H 4 P2 O 3 2 P2 O 5 → H3PO4 → salts - phosphates → H3PO3 → salts-phosphites phosphorous acid Forms of phosphoric acid: HPO3 - meta - H3PO4 - ortho - H4P2O7 - pyro - Na3PO4 Na2HPO4 NaH2PO4 Phosphorus fertilizers Ca (H2PO4)2 – calcium dihydrogen phosphate (superphosphate) CaHPO4∙2H2O – calcium hydrogen phosphate (precipitate) Ca3(PO4)2 ↓ - calcium phosphate (phosphate flour) Phosphorus is the most important macronutrient. d – elements of large periods In large periods (4, 5, 6, 7), 10 d – elements are introduced between the s- and d- elements, in which electrons fill the d – sublevel of the second level from the outside. All d-elements, regardless of the group number, have two electrons in the outer level (or one electron if electron leakage occurs). Therefore d - elements exhibit metallic properties. In d-elements, the valence electrons are the electrons of two levels - the outer level and the d- sublevel of the penultimate level, that is, (n – 1)d and nSnP. d–elements form side subgroups of the periodic table. Most d - metals have several oxidation states, so their compounds exhibit redox properties. Oxides and hydroxides of d-elements have basic, amphoteric or acidic properties depending on the degree of oxidation of the d-element forming them. d– metals are good complexing agents and form stable complex compounds. d – metals manganese, cobalt, copper, zinc are vital essential microelements . Sulfates of these metals are used as microfertilizers, which are used for pre-sowing treatment of seeds and foliar feeding of plants. Example: Iron Fe Fe 1s22s22p63s23р63d64s2 valence levels No. 26 26 ē 6th d-element Electron graphic formulas Oxidation states of Fe: +2, +3, +6. 26 Copyright JSC Central Design Bureau BIBKOM & LLC Kniga-Service Agency The most important compounds 2 Fe O → Fe(OH)2 main basic oxide properties 3 H3FeO3 → ferrite salts KFeO2 Fe2 O3 → Fe(OH)3 weak weakly amphoteric amphoteric properties of Н2О НFeO2 6 Fe O3 → acidic properties of H2FeO4 → iron acid K2FeO4 salt-ferrates (oxidizing agents) Iron is an essential trace element. 3.4. Periodicity of changes in the properties of elements and their compounds The properties of chemical elements change periodically. Periodically changing: atomic radii (R); ionization energy (ionization potential I); electron affinity (E) electronegativity (χ) χ=I+E or χ= I E . 2 The radii of atoms (R) in a period decrease with increasing atomic number (horizontal periodicity). Reason: The charge of the atomic nucleus increases, but the number of energy levels in a period is constant. The smaller the radius, the more difficult it is for the atom to give up electrons (decreasing metallic properties). In long periods the following is observed: 1. d – compression (decreasing the radius in the row of d – elements); 2. f – compression (decreasing the radius in the row of f – elements); a) in period VI – lanthanide compression; b) in the VII period – actinoid compression. In groups, as the serial number increases, the radii of the elements increase (the main pattern). Reason: Increase in the number of energy levels. Consequence: Increased ability to donate electrons → increased metallic properties. Table 3.2 - Normal change in radius in the IA group Element R, nm Li 0.155 Na 0.189 K 0.236 Rb 0.248 The metallic properties of the elements increase. 27 Cs 0.268 Fr 0.280 Copyright JSC Central Design Bureau BIBKOM & LLC Kniga-Service Agency Table 3.3 - Violation of the pattern of changes in R due to f - compression in the IB group Element R, nm Cu 0.128 Ag 0.144 Au 0.144 In secondary subgroups, metallic properties elements decrease (the influence of lanthanide compression). Ionization energy is the energy required to remove one electron from an atom. Ionization potential I is the smallest voltage (in electron volts) that must be applied to an atom to remove one electron from it. Ionization potential I characterizes the metallic and reducing properties of elements. E0 = E+ + ē atom ion General pattern: In periods with increasing atomic number, the ionization potential increases. The value of I reaches its maximum towards the end of the period. General pattern: The ionization potentials in groups decrease with increasing atomic number. Violation of the general pattern: (influence of the degree of filling of the sublevel). Sublevels that are completely filled or half filled with electrons are the most stable (I increases). Table 3.4 - Influence of the degree of sublevel filling on the ionization potential Element Valence electrons I, eV N 2s22p3 14.53 O 2s22p4 13.61 Table 3.5 - Influence of lanthanide compression on the ionization potential Element R, nm I, eV Cu 0.128 7.73 Ag 0.144 7 .57 Au 0.144 9.22 Electron affinity (E, eV) is the energy that is released when one electron is added to an atom. Electron affinity characterizes the non-metallic, oxidizing properties of elements. General pattern: In a period with increasing atomic number, electron affinity increases. In a group, as the atomic number increases, the electron affinity decreases. 28 Copyright JSC Central Design Bureau BIBKOM & LLC Kniga-Service Agency Table 3.6 - Change in the value of electron affinity in group VII - A Element E, eV F 3.62 Cl 3.82 Br 3.54 J 3.24 Electronegativity (χ) is the ability of atoms in a molecule to attract electrons to themselves. A measure of electronegativity is considered to be energy equal to the sum of ionization energies (I) and electron affinity (E) χ = I+E. It was conventionally assumed that Li 1. Conventionally, the boundary between metals and non-metals is 2 (or 1.7). nonmetals metals 2< 2 χ > 2 Figure 3.1. Conventional boundary of metals and non-metals Table 3.7 - Electronegativity of atoms of some elements Element χ Na 1.01 Mg 1.23 Al 1.47 metals Ge 2.02 Cl 2.83 F 4.10 non-metals The chemical properties of elements change periodically. During the period, the properties of the elements change from metallic to non-metallic and to inert (noble) gas. The oxidation states of elements change periodically. The oxidation state is the charge of an atom in a compound conventionally assumed to be ionic. The highest oxidation state is its greatest positive value. It is equal to the group number. Exceptions: O, F, Fe, Cu, Ag, Au. The lowest oxidation state is its lowest value. All other oxidation states are intermediate or intermediate. The oxidation state has a (+) or (–) sign. Constant oxidation states: 1) alkali metals of I A-group: Li, Na, K, Rb, Cs, Fr: +1 2) metals of II A-group: Be, Mg, Ca, Sr, Ba, Ra: +2 3 ) Al: +3 4) H: +1 (exception: NaH hydrides) 5) O: – 2 (exception: OF2) 6) The oxidation state of the atom is zero (the atom is electrically neutral) 7) The oxidation state of the molecule is zero (the molecule is electrically neutral) neutral). The valency of elements changes periodically. Valence is measured by the number of chemical bonds by which a given atom is connected to other atoms. Valence has no sign. The values of oxidation state and valence may not coincide. 29 Copyright JSC Central Design Bureau BIBKOM & LLC Kniga-Service Agency Table 3.8 – Oxidation state and valence of nitrogen in NH3 and NH4Cl Substance NH3 NH4Cl Oxidation state of nitrogen –3 –3 Valency of nitrogen 3 4 The properties of chemical elements affect the properties of their corresponding complex connections. Scheme 3.2. Characteristics of the properties of elements Questions for self-control 1. Give the formulation of the periodic law of D.I. Mendeleev. Explain how you understand it. 2. Compare the atomic structure and properties of any macroelement and microelement. 3. Explain why noble gases are inert and alkali metals are active. 4. Compare the radii of atoms and ionization potentials. 1) sodium and potassium; 2) oxygen and fluorine. 5. Give examples of the dependence of the properties of complex compounds on the properties of the elements that form them. REFERENCES Main 1. Glinka, N.L. General chemistry / N.L. Glinka – M.: KNORUS, 2009. – 752 p. 2. Knyazev, D.A. Inorganic chemistry/D.A. Knyazev, S.N. Smarygin – M.: Bustard, 2004. – 592 p. 3. Ryazanova, G.E. General and inorganic chemistry. Tables and diagrams / G.E. Ryazanov. – Saratov: Federal State Educational Institution of Higher Professional Education “Saratov State Agrarian University”, 2006 – 284 p. 30 Copyright JSC Central Design Bureau BIBKOM & LLC Book-Service Agency Additional 1. Inorganic chemistry (biogenic and abiogenic elements): Textbook / ed. V.V. Egorova. – St. Petersburg: Lan Publishing House, 2009. – 320 p. 2. Guzey, L.S. General chemistry / L.S. Guzey, V.N. Kuznetsov, A.S. Guzey. – M.: Publishing house. Moscow State University, 1999. – 333 p. 3. Agafoshin, N.P. Periodic law and periodic system D.I. Mendeleev. / N.P. Agafoshin - M.: Education, 1973. - 208 p. 4. Dmitriev, S.N. Chemical identification and study of the properties of superheavy metals. Evolution of the periodic system D.I. Mendeleev / S.N. Dmitriev – Abstracts of reports of the XVIII Mendeleev Congress on General and Applied Chemistry: In 5 volumes; v.1. – M.: Granitsa, 2007. – P. 47. 31 Copyright JSC “CDB “BIBKOM” & LLC “Agency Kniga-Service” Lecture 4 MANIFESTATION OF THE PERIODIC LAW IN THE ACID-BASE PROPERTIES OF INORGANIC COMPOUNDS 4.1. Genetic connection of the main classes of inorganic compounds In nature, there are a huge number of inorganic compounds that belong to a small number main classes of inorganic compounds. Oxides and hydroxides differ in the degree to which they exhibit basic or acidic properties. Hydroxides vary in electrolyte strength. Each substance, in addition, also has its own individual characteristics. It is necessary to clearly imagine this dialectic of the chemical properties of elements in order to be able to imagine the dialectic of the chemical properties of complex compounds. The properties of the main classes of inorganic compounds help to understand the periodic law and allow us to rely on it as reliable foundation knowledge. In all phenomena and processes of nature, there are contradictory, mutually exclusive, opposing tendencies. These opposing trends are manifested in the properties of metal atoms that can donate electrons, and in the properties of non-metal atoms that can gain electrons. As the composition of substances becomes more complex, these opposing trends continue to appear. Typical metals and transition elements in low oxidation states form basic oxides, while typical nonmetals and transition elements in high oxidation states form acidic oxides with opposite properties. With further complication of the composition of substances, hydroxides are formed, with bases corresponding to basic oxides, and acids corresponding to acidic oxides. Bases and acids are also opposite in properties. Thus, a change in the composition of substances leads to a change in their properties. The contrast in the properties of metals and non-metals determines the contrast in the properties of both basic and acidic oxides and bases and acids. This is proof that the properties of complex substances are in dialectical dependence on the properties of the atoms of the elements, and, therefore, they are interconnected and interdependent. The development of the chemical form of the movement of matter is manifested in the complication of the composition and properties of substances. Indeed, the periodic law contains the idea that the properties of complex compounds depend on the properties of the elements from which they are formed. Therefore, one should expect, for example, that metals of different activity will correspond to complex compounds that differ in properties from each other, and metals and non-metals that are fundamentally different from each other will correspond to complex compounds that are sharply different from each other. Complex compounds are formed from simple substances. The simplest of complex compounds are oxides. Oxides are complex substances consisting of two elements, one of which is oxygen. When the oxide molecule is complicated by adding water (directly or indirectly), hydroxides are obtained. The existence of the main classes of inorganic compounds reflects the development of inorganic matter. 32 Copyright JSC Central Design Bureau BIBKOM & LLC Book-Service Agency Scheme 4.1. Genetic relationship of the main classes of inorganic compounds and levels of development of a chemical substance Table 4.1 - Levels of development of a chemical substance 4 3 2 1 Salts Hydroxides Oxides Atoms of elements Chemical phenomena are characterized by internal inconsistency. It is the reason for the diversity of chemical properties and the impossibility of explaining them within the framework of the best schemes without any exceptions that go beyond this framework. 4.2. Chemical properties of oxides, bases, acids and salts Basic oxides Basic oxides are complex substances consisting of a metal and oxygen, to which bases correspond as hydroxides. Table 4.2 – Patterns of development of metals Chemical properties of basic oxides Basic oxides react: 1) with acidic oxides: Na2O + CO2 = Na2CO3; 2) with acids: CaO + H2SO4 = CaSO4 + H2O; 3) with water (oxides of only the most active metals - alkali and alkaline earth) react: CaO + H2O = Ca(OH)2 CuO + H2O 33 does not work Copyright JSC Central Design Bureau BIBKOM & LLC Kniga-Service Agency Acidic oxides Acidic Oxides are complex substances consisting of a non-metal and oxygen, to which acids correspond as hydroxides. Table 4.3 – Patterns of development of non-metals Chemical properties of acid oxides Acidic oxides react: 1) with basic oxides: SO3 + Na2O = Na2SO4; 2) with bases: CO2 + 2NaOH = Na2CO3 + H2O; 3) with water: SO3 + H2 O = H2SO4. Bases Bases are substances consisting of a metal and hydroxyl groups OH–. Bases are substances that dissociate to form hydroxide ions OH–. Very active alkali metals (group 1, main subgroup: Na, K, etc.) alkaline earth metals (group 2, main subgroup: Ca, Sr, Ba) form strong bases (alkalis). They are soluble in water, dissociate completely (NaOH, KOH , Ca(OH)2, Ba(OH)2). Low-active metals (all except alkali and alkaline earth): Al, Zn, Cu, Fe, Pb form weak bases. They are poorly soluble in water and partially dissociate: (Cu(OH)2↓, Fe(OH)3↓). More active metals with a lower degree of oxidation form stronger bases 2 3 Fe (OH)2 > Fe (OH)3 Preparation of strong bases 1) oxide with water: CaO + H2O = Ca(OH)2; 2) metal with water: 2Na + H2O = 2NaOH + H2; 3) electrolysis of salt solution: NaCl, KCl. 34 Copyright JSC Central Design Bureau BIBKOM & LLC Kniga-Service Agency Preparation of weak bases Salt of a low-active metal with an alkali: FeSO4 + 2NaOH = Fe(OH)2↓ + Na2SO4. Chemical properties of bases Bases react: 1) with acids (neutralization reaction): 2NaOH + H2SO4 = Na2SO4 + 2H2O; 2) with acid oxides: 2NaON + CO2 = Na2CO3 + H2O; 3) with salts (if a precipitate forms): 2NaOH + CuSO4 = Cu(OH)2 ↓ + Na2SO4. Rule: Soluble salts of low-reactive metals react with strong bases. Professional competence: The use of microfertilizers CuSO4, FeSO4, MnSO4, etc. on alkaline soils is ineffective, since they become insoluble. Amphoteric hydroxides Amphoteric hydroxides have dual properties - both bases and acids at the same time. They react with both acids and bases. Table 4.4 – Metals that form amphoteric hydroxides (internal inconsistency of properties) Group Elements Oxidation state I Au +3 II Be, Zn +2 III Al +3 IV Ge, Sn, Pb +2; +4 VI Cr +3 Amphoteric properties of zinc hydroxide Zn(OH)2: amphoteric. Zn(OH)2 ↓+ H2SO4 = ZnSO4 + 2H2O base acid In the melt: amphote. Zn(OH)2 + 2NaOH = H2ZnO2 + 2NaOH = Na2ZnO2 + 2H2O acid base sodium zincate In solution: Zn(OH)2 + 2NaOH = Na2 or Na Metals with the same oxidation state (+2) form amphoteric hydroxides with similar properties: Zn (OH)2, Be(OH)2, Ge(OH)2, Zn(OH)2, Pb(OH)2 Amphoteric properties of aluminum hydroxide: amphote. 2Al(OH)3 + 3H2SO4 = Al2(SO4)3 + 6H2O base acid 35 Copyright JSC Central Design Bureau BIBKOM & LLC Kniga-Service Agency In the melt: amphote. Al(OH)3 + NaOH → H3AlO3 + NaOH → NaAlO2 + 2H2O acid base H2O sodium metaaluminate metaaluminum acid HAlO2 In solution: Al(OH)3 + 3NaOH = Na3 or Na or Na Metals with the same oxidation state (+3) form amphoteric hydroxides with similar properties: Al(OH)3, Cr(OH)3, Au(OH)3. Acids Acids are complex compounds that dissociate to form hydrogen cations H+ (protons). Table 4.5 - Acids and salts Name of acid Hydrochloric (hydrochloric) Nitric Acid formula Acid residue Name of salt Examples of salts HCl Cl - Chlorides NaCl HNO3 Nitrates Sulfates hydrosulfates Phosphates hydrophosphates dihydrogen phosphates sulfites hydrosulfites sulfides hydrosulfides carbonates hydrocarbonates silicates hydrosilicates Nitrites KNO3 Na2SO4 NaHSO4 Ca3( PO4)2 CaHPO4 Ca(H2PO4)2 Na2SO3 NaHSO3 Na2S NaHS Na2CO3 NaHCO3 Na2SiO3 NaHSiO3 NaNO2 Cyanides KCN Sulfur H2SO4 Phosphorus H3PO4 Sulfur H2SO3 Hydrogen sulfide H2S Coal H2CO3 Silicon H2SiO3 Nitrous Hydrogen cyanide HNO2 NO3– SO42– HSO4 – PO43– HPO42– H2PO4– SO32– HSO3 – S2– HS– CO32– HCO3– SiO32– HSiO3– NO2– HCN CN– Chemical properties of acids Acids interact with: 1) bases 2HCl + Ca(OH)2 = CaCl2 + 2H2O; 2) basic oxides H2SO4 + CaO = CaSO4 + H2O; 3) salts (if a precipitate or gas is formed): 36 Copyright JSC Central Design Bureau BIBKOM & LLC Kniga-Service Agency a) H2SO4 + BaCl2 = BaSO4 + 2HCl, b) 2HCl + CaCO3 = CaCl2 + CO2 + H2O; 4) Interaction of acids with metals Ordinary (typical) acids H Cl, H 2 SO4. (detailed). The oxidizing agent is H+. Acids react with metals to release hydrogen. A series of voltages of metals: Li, K, Na, Mg, Al, Mn, Zn, Fe, Ni, Sn, Pb, H Cu, Hg, Ag, Pt, Au Metals located in a series of voltages up to hydrogen displace hydrogen from ordinary acids . In dilute H2SO4, the oxidizing agent is hydrogen cations H+ (they accept electrons and change the oxidation state). 0 2 H 2 SO4 Z n abd. 0 2 2е Zn Zn 0 2е 2H 0 Zn SO4 H2 ok H2 1 1 does not go. H 2 SO 4 Cu r dil. 6 5 5 Oxidizing acids: H2 S O4, H N O3, H N O3 conc. conc. div. 6 5 The oxidizing agent is an acid-forming element: S, N. Oxidizing acids react with metals without releasing hydrogen. Oxidizing acids can even react with Cu, Hg, Ag (without releasing hydrogen). Concentrated H2SO4 reacts with metals without releasing hydrogen (the oxidizing agent is the SO 42 ion, i.e. S). 0 2е Cu sol 6 S 2 Cu 2е okl 4 S 1 1 Concentrated sulfuric acid H2SO4 passivates Al, Fe, Cr: 3H2SO4 + 2Fe → Fe2O3 + 3SO2 + 3H2O conc. film 37 Copyright JSC Central Design Bureau BIBKOM & LLC Book-Service Agency Properties nitric acid Nitric acid has oxidizing properties. 5 In nitric acid, the oxidizing agent is NO3 ions (N accepts electrons, changes the oxidation state). When nitric acid reacts with metals, hydrogen is not released. Concentrated HNO3 usually turns into brown gas NO2. 5 4 HNO3 NO 2 brown end HNO3 passivates Al, Fe, Cr. conc. HNO 3 Al Al 2 O 3 NO 2 conc. H 2O HNO3 film dissolves Cu, Hg, Ag: conc. HNO 3 Cu conc. Cu NO 3 2 NO 2 H 2O Table 4.6 - Interaction of nitric acid with metals 5 HNO3 dil. low active Cu, Ag, Pb metals of medium activity Mg, Zn very active Na, K, Ca 2 3 NO N2 ;N2 O HNO 3 Cu Cu(NO 3) 2 NO H 2O HNO 3 Zn Zn(NO 3) 2 N 2 H 2O dil. div. HNO 3 Na dil. NaNO 3 NH 3 H 2 O or HNO 3 Na dil. 3 NH 3, NH 4 NO 3 0 NaNO 3 NH 4 NO 3 H 2 O The products of interaction of dilute acid with metals depend on the activity of the metal and the degree of dilution of nitric acid. The more active the metal, the more reduced the product of its transformation is. Usually several products are formed at once. The kinetics of the process determines the advantage of the formation of one of them. 38 Copyright JSC Central Design Bureau BIBKOM & LLC Kniga-Service Agency Salts Salts are electrolytes that dissociate into metal cations and anions of an acid residue. Medium salts consist of a metal and an acid residue. 1 Example: Na 3 PO 43 – sodium phosphate. Acid salts (hydrosalts) consist of a metal cation and a complex anion containing hydrogen and an acid residue. 1 Example: Na 2 HPO 42 – sodium hydrogen phosphate. Basic salts (hydroxo salts) consist of a complex cation containing a metal and a hydroxyl group, and an acid residue anion. 2 Example: AlO H Cl 41 – aluminum hydroxychloride. Chemical properties of salts Salts react: 1) with bases (if a precipitate is formed): FeCl3 + 3NaOH = Fe(OH)3 + 3NaCl; 2) with acids (if a precipitate forms or gas is released): AgNO3 + HCl = AgCl + HNO3, Na2CO3 + 2HCl = 2NaCl + CO2 + H2O; 3) with salts (if a precipitate forms): Na2SO4 + BaCl2 = BaSO4 + 2NaCl 4) with metals (more active than the metal forming the salt): CuSO4 + Fe = Cu + FeSO4. 4.3. Periodicity of changes in the acid-base properties of chemical substances The acid-base properties of complex compounds depend on the properties of the elements from which they were formed: 1. The properties of elements in periods change from metallic to non-metallic. In accordance with this, the properties of complex compounds (oxides and hydroxides) periodically change from basic to acidic. 2. Active metals (alkali and alkaline earth) form strong bases (alkalis); 3. Less active metals form weak bases; 4. More active metals with a lower oxidation state correspond to stronger bases; 5. More active nonmetals with a higher oxidation state correspond to stronger acids; 6. Active nonmetals correspond to strong acids. 39 Copyright JSC Central Design Bureau BIBKOM & LLC Book-Service Agency 4.4. Questions of professional competence 1. Superphosphate Ca(H2PO4)2 is not advisable to use on alkaline soils, since it turns into insoluble Ca3(PO4)2: Ca(H2PO4)2 + 2Ca(OH)2 = Ca3(PO4)2↓+ 4H2O soluble alkaline medium insoluble 2. Phosphorite flour Ca3(PO4)2 is used on acidic soils, where it gradually turns into soluble calcium dihydrogen phosphate: Ca3 (PO4)2 + H3PO4 → Ca(H2PO4)2 insoluble acidic medium soluble 3. Microfertilizers (CuSO4, MnSO4 etc.) is not advisable to use on alkaline soils, since they turn into insoluble compounds: CuSO4 + 2NaOH = Cu(OH)2↓ + Na2SO4 soluble alkaline insoluble medium 4. When liming acidic soils, limestone CaCO3 or dolomite flour CaCO3 MgCO3 is added to the soil . In acidic soil, processes can occur due to which the acidity of the soil decreases: +H+ Ca2+ + CO2 + H2O a) CaCO3 +H+ Ca(HCO3)2 → Ca2+ + HCO 3 b) Limestone CaCO3 interacts with the H+ cations of the soil absorption complex ( PPK): N PPK N CaCO 3 PPK Ca 2 CO 2 H 2O 5. When gypsuming alkaline saline soils, gypsum CaSO4 2H2O is added to the soil. In the soil of soda salinization, the following processes occur: a) Na 2 CO3 CaSO4 CaCO3 Na 2SO 4 improves the soil - it is easily washed out by rain, irrigation b) PPC Na Na CaSO 4 PPC Ca 2 Na 2 SO 4 is easily washed away Questions for self-control 1. List the main classes inorganic compounds in order of their complexity. 2. Using a genetic diagram, explain the chemical properties of each class of inorganic compounds. 3. Explain the algorithm for compiling and checking formulas of complex chemical substances. Give examples. 4. Complete the equations of possible reactions: 1) CaO + H2O → 6) NaOH + Zn(OH)2 → 2) CuO + H2O → 7) Na2SO4 + KCl → 3) Cu + HCl → 8) Na2SO4 + BaCl2 → 4) Cu + HNO3 (conc.) → 9) CaCO3 + HCl → 5) NaOH + Ca(OH)2 → 10) CaCO3 + H2SiO3 → 5. Draw genetic diagrams of the main classes of inorganic compounds. 40 Copyright JSC Central Design Bureau BIBKOM & LLC Kniga-Service Agency 6. Write the formulas of the following compounds: sodium sulfate, aluminum chloride, calcium phosphate, calcium dihydrogen phosphate, copper nitrate, aluminum hydroxychloride, aluminum hydrosulfate, sodium hydrosulfide. 7. Give an example of the dependence of the acid-base properties of a chemical compound on the properties of the element that forms it. 8. Explain why it is inappropriate to use superphosphate and microfertilizers (manganese sulfate, copper sulfate) on alkaline soils. 9. What processes can occur in the soil during chemical reclamation of a) acidic soils b) alkaline soils 10. Which microfertilizer FeSO4 or ZnSO4 will be more easily absorbed on alkaline soils? REFERENCES Main 1. Ryazanova, G.E. General and inorganic chemistry. Tables and diagrams / G.E. Ryazanov. – Saratov: Federal State Educational Institution of Higher Professional Education “Saratov State Agrarian University”, 2006 – 284 p. 2. Ryazanova, G.E. Inorganic and analytical chemistry / G.E. Ryazanova - Saratov: Federal State Educational Institution of Higher Professional Education "Saratov State Agrarian University", 2006. - 172 p. Additional 1. Klinsky, G.L. Inorganic chemistry / G.L. Klinsky, V.D. Skopintsev. – M.: Publishing house. MSHA, 2001. – 384 p. 2. Egorov, V.V. Environmental chemistry. /V.V. Egorov. – St. Petersburg: Lan Publishing House, 2009. – 192 p. 3. Vlasov, V.M. Mistakes leading to explosion. / V.M. Vlasov // Chemistry and life. – 2006. – No. 7. – p. 60. 4. http://www.online-knigi.com/biologiya/agrohimiya-uchebnik-skachat.html 41 Copyright JSC Central Design Bureau BIBKOM & LLC Kniga-Service Agency Table 4.7 - Properties of elements of the III period and corresponding im oxides and hydroxides Element R of the R atom of the ion J, eV Properties Na 0.189 0.80 5.14 active metal Oxide Na2O basic oxide Hydroxide NaOH strong base Mg 0.160 0.740 7.65 less active metal MgO weak basic oxide Mg(OH)2 weak base 5.90 0.074 Kd (dissociation constant) Al 0.143 0.570 5.99 even less active amphoteric metal Al2O3 amphoteric oxide Al(OH)3 weak base with amphoteric (dual - both basic and acidic) 4 10-13 (acidic) 1 .38 10-9 (basic) 42 Si 0.134 0.390 8.15 slightly active non-metal SiO2 weakly acidic oxide H2SiO3 weak acid, insoluble in water P 0.130 0.350 10.49 more active non-metal P2O5 moderately acidic oxide H3PO4 medium-strength acid 2.2 10-10 7.52 10-3 S 0.104 0.290 10.36 active nonmetal SO3 strongly acidic oxide H2SO4 strong acid Cl 0.999 0.260 13.01 very active nonmetal Cl2O7 strongly acidic oxide HClO4 very strong acid Copyright JSC Central Design Bureau BIBKOM & LLC "Agency Book-Service" Lecture 5 CHEMICAL BONDING 5.1. Modern ideas about chemical bonding Under ordinary conditions, very few substances consist of individual atoms. These are inert gases (for example, helium He). The remaining substances consist of more complex particles (molecules, molecular ions, radicals). A molecule is a collection of atoms that has a number of characteristic properties. The properties of a molecule depend on the strength of the chemical bonds and the geometry of the molecule. The geometry of a molecule is the spatial structure of the molecule, which is determined by bond angles and bond lengths. Molecules are formed from atoms between which a chemical bond occurs. There are different definitions of the concept of “chemical bond”. A chemical bond is the phenomenon of joining atoms into molecules (L.S. Guzei). A chemical bond is the interaction of atoms, leading to the formation of molecules and crystals of simple and complex substances and ensuring their stability (D.A. Knyazev). A chemical bond is a set of interactions between electrons and nuclei, leading to the joining of atoms into molecules (Yu.A. Ershov). Chemical bond theory is a branch of chemistry in which the properties of chemical bonds are used to describe the properties of chemical compounds. Currently there is no unified theory of chemical bonding. There are several theories of chemical bonding (valence bond method, molecular orbital method, crystal field theory). 5.2. Basic principles of the valence bond method (VBC) W. Heitler and F. London (1927) A chemical bond is carried out by two electrons with opposite spins. The chemical bond is two-electron, two-center, localized. A chemical bond carried out by electron pairs is called covalent. Currently, the method of valence bonds is important for a qualitative understanding of the nature of chemical bonds 5.2.1. Mechanisms of formation of a covalent bond Exchange mechanism A chemical bond is formed due to the overlap of two one-electron orbitals (Fig. 1). Figure 5.1. Scheme of overlap of atomic electron clouds according to the exchange mechanism. When a chemical bond is formed, the energy of the system decreases, reaching a minimum. 43 Copyright JSC Central Design Bureau BIBKOM & LLC Kniga-Service Agency Donor-acceptor mechanism of bond formation A chemical bond is formed when the two-electron orbital of one atom overlaps with the free orbital of another atom. Figure 5.2. Scheme of overlap of atomic electron orbitals according to the donor-acceptor mechanism. Chemical properties of ammonia The chemical properties of ammonia are determined by its ability to interact with water and acids according to the donor-acceptor mechanism, forming ammonium compounds. In this case, the donor is ammonia nitrogen, which has a non-bonding electron pair, and the acceptor is a hydrogen cation, which has a free orbital. Ammonium compounds – NH4NO3 (ammonium nitrate), (NH4)2SO4 (ammonium sulfate), NH4Cl (ammonium chloride), NH4H2PO4 and (NH4)2HPO4 (ammophos) – are valuable mineral fertilizers. Their production in industry is based on the reactions: HCl → NH4Cl: NH3 + HNO3 → NH4NO3 2 NH3 + H2SO4 → (NH4)2SO4 NH3 + H3PO4 → NH4H2PO4 2 NH3 + H3PO4 → (NH4)2HPO4: NH3 + 5.2.2. Covalent bond When a chemical bond is formed, the energy of the system decreases and reaches a minimum. The formation of a chemical bond is an exothermic process. A covalent bond is characterized by the following properties: Bond length is the distance between the nuclei in the molecule. The shorter the bond length, the stronger the chemical bond. Bond energy is the energy released during the formation of a chemical bond (kJ/mol). The higher the binding energy, the stronger the bond. Bond saturation is the ability of atoms to form a certain limited number of covalent bonds. 44 Copyright JSC Central Design Bureau BIBKOM & LLC Book-Service Agency Scheme 5.1. Properties of covalent bonds 5.2.3. Types of chemical bonds - a bond is formed when electron orbitals overlap on a line connecting the centers (nuclei) of atoms. Figure 5.3. Scheme of formation-bonds The orbitals overlap deeply, ensuring maximum energy gain. If a single (single) bond occurs between two atoms, then it is a bond. -bond is formed by overlapping p-orbitals located perpendicular to the main bond line. Figure 5.4. Scheme of -bond formation Two areas of overlap arise, on both sides of the straight line connecting the nuclei of atoms. The π bond, which is less strong than the - bond, breaks first. A π bond is formed only in the presence of an - bond. These connections differ significantly. Around the connection, free axial rotation is possible. Rotation around the π bond is impossible, since it has two overlapping regions. 45 Copyright JSC Central Design Bureau BIBKOM & LLC Kniga-Service Agency Multiplicity of chemical bonds A covalent bond carried out by one electron pair is called single. The multiplicity of a chemical bond is the number of electron pairs forming bonds (double or triple bond). The properties of chemical substances, their activity, depend on the number of electron pairs that form chemical bonds in the molecules. Thus, unsaturated hydrocarbons with a double or triple bond (alkenes or alkynes) are oxidized by potassium permanganate, discoloring bromine water, and saturated hydrocarbons with a single bond with an unbranched chain do not enter into such reactions under normal conditions. 5.2.4. Types of chemical bonds Covalent nonpolar bond A covalent nonpolar bond is formed between identical atoms of the same element with the same electronegativity. Non-polar molecules of simple substances: H2, O2, N2, Cl2 H + H H:H electron pair is located symmetrically Covalent polar bond A covalent polar bond is formed between atoms of different elements with different electronegativity. HCl, H2O, NH3, H2S The more polar a molecule is, the more reactive it is usually. Polarization is the displacement of the connecting electron cloud. A measure of polarization is the effective atomic charges (q). Many symmetrically constructed molecules of complex substances are non-polar, although the bonds between them are polar: CH4, CCI4, SO3, BeCl2, CO2, C6H6 O=C=O 8.99∙10-30 C∙m ← → 8.99∙10-30 C ∙m The electric dipole moment between charges is equal to the product of the dipole charge and the distance = ql (C∙m) Electric dipole moments – vector quantities, therefore the dipole moments of the bonds in the molecule are summed up as vectors. 5.2.5. Degree of ionicity of a bond 1. The calculated theoretical charge of an atom in a molecule is equal to the oxidation state 1 1 HF 1 2 H2 O 46 3 1 N H3 Copyright JSC Central Design Bureau BIBKOM & LLC Kniga-Service Agency 2. The effective charge of an atom in a molecule characterizes the displacement electron density and is determined experimentally. 0.42 0.42 0.18 0.18 0.05 0.05 H F H Cl H I The degree of ionicity i (%) is equal to the ratio of the effective qeff and the calculated qcalc charges of the atom. qeff i 100%. q calc Table 5.1 – Degree of ionicity i(%) of some chemical compounds Substance HF LiF NaCl NaF RbCl i (%) 42 84 90 97 99 The degree of ionicity is equal to the difference in electronegativity Δχ of the atoms forming a given bond. The value Δχ = 1.7 is conventionally taken to correspond to the degree of ionicity of the bonds, equal to 50%. It is conventionally assumed that bonds with Δχ >1.7 and i >50% are ionic, and bonds with Δχ<1,7 и i <50% - ковалентными полярными. Степень ионности в неполярных молекулах (H2, O2, Cl2) равна нулю. Внесистемной единицей дипольного момента является Дебай (Д). 1Д = 3,33∙10-30 Кл∙м Если дипольный момент молекулы равен нулю, то молекула является неполярной. Если дипольный момент молекулы отличен от нуля, то молекула является полярной. Таблица 5.2 – Зависимость типа химической связи от величины μ (Д) Дипольный момент (Д) Тип химической связи 0 До 4Д 4-11Д неполярная полярная ионная Ионная связь Ионная связь является предельным случаем ковалентной полярной связи. Ионная связь – это химическая связь между ионами, осуществляемая электростатическим притяжением Ионная связь образуется между элементами с резко противоположными свойствами (активными металлами и неметаллами) 47 Ионная связь – это предельно поляризованная связь Copyright ОАО «ЦКБ «БИБКОМ» & ООО «Aгентство Kнига-Cервис» Свойства ионной связи ненаправленность ненасыщаемость соединения с ионной связью легко диссоциируют образуются твердые тела с ионным кристаллическим строением Под действием полярных молекул растворителя происходит электролитическая диссоциация электролитов. Чем более полярна химическая связь, тем легче идет диссоциация. Диссоциация легче всего идет по ионной связи. Таблица 5.3 – Условное деление вещества на полярные и ионные Степень ионности i XA - XB > 50 % < 50% > 1,7 < 1,7 (Д) Тип вещества > 4-11 > 0-4 ionic polar 5.2.6. Hybridization of atomic orbitals Hybridization is the mixing of different atomic orbitals, leading to their alignment in shape and energy. Figure 5.5. Scheme of hybridization of valence orbitals The number of hybrid orbitals formed is equal to the number of initial atomic orbitals participating in hybridization. The simplest cases of hybridization occur when the s- and p-orbitals are mixed. 1. With sp - hybridization, two orbitals are mixed - s (ball shape) and p (dumbbell shape). In this case, two hybrid orbitals are formed, having the shape of an asymmetric dumbbell. They are located on the same line at an angle of 180°. This determines the geometry of molecules with sp-hybridization of the central atom. This molecule has a linear structure. 48 Copyright JSC Central Design Bureau BIBKOM & LLC Kniga-Service Agency An example is the BeF2 molecule, in which there is sp - hybridization of the beryllium atom in the excited state. Figure 5.6. Diagram of a linear molecule BeF2 2 2. During sp - hybridization, three orbitals are mixed - one s and two p - orbitals. In this case, three hybrid orbitals are formed. They are located in the same plane and oriented at an angle of 120° to each other. This type of hybridization corresponds to the geometry of a flat triangle. An example is the boron fluoride molecule BF3. Figure 5.7. Diagram of a triangular BF3 molecule 3. During sp3 hybridization, four hybrid orbitals are formed. They are oriented at an angle of 109°28" to each other, elongated to the vertices of the tetrahedron. An example is the CH4 methane molecule, which has the shape of a tetrahedron. Figure 5.8. Diagram of a tetrahedral methane molecule 4. In the molecules of ammonia NH3 and water H2O, sp3 also occurs - hybridization of nitrogen atoms and oxygen. The bond angle in these molecules is close to the tetrahedral one, but not equal to it. In the NH3 molecule the bond angle is 107.3°. The distortion of the tetrahedron angle occurs due to the influence of one non-bonding orbital. In the water molecule the bond angle is 104.5° due to the influence of two non-bonding orbitals orbitals Ammonia NH3 (trigonal pyramid shape) Water H2O Angular shape of the molecule Figure 5.9. Diagram of the influence of non-bonding orbitals on the geometry of molecules 49 Copyright JSC Central Design Bureau BIBKOM & LLC Kniga-Service Agency 5.2.7 Metallic bond Metallic bond is communication between a large number of metal cations located in the nodes of the metal crystal lattice, carried out using freely moving electrons (electron gas). The structural features of metals lead to such properties as electrical conductivity, thermal conductivity, ductility, and metallic luster. The most common crystal lattices of metals are cubic and face-centered. The concept of binding energy does not apply to metallic bonds due to their multicenter nature. 5.2.8. Hydrogen bond A hydrogen bond is a bond between a hydrogen bonded to an atom of a strongly electronegative element and an atom of the same element of another (or the same) molecule. Hydrogen bonding is due to: 1. Electrostatic attraction between a proton and a polar group. 2. Donor-acceptor interaction. Intermolecular hydrogen bond: 1. H2O association; (H2O)2; (H2O)5; (H2O)n Hydrogen bonds Scheme 5.2. The influence of hydrogen bonds on the properties of chemical substances 5.3. The concept of the molecular orbital method (1928-1932) R. Mulliken American physical chemist Mulliken Robert Sanderson. For his study of chemical bonds and electronic structures of molecules using the molecular orbital method, he was awarded the Nobel Prize in Chemistry (1966). 50 Copyright JSC Central Design Bureau BIBKOM & LLC Kniga-Service Agency Basic principles of the molecular orbital method A molecule is a set consisting of nuclei and electrons, in which each electron moves in the field of all other electrons and all nuclei. A chemical bond can be not only two-electron, but also three-electron, multielectron and multicenter. The method of molecular orbitals is the method of atomic orbitals extended to a molecule. A molecular orbital is a linear combination of atomic orbitals. The molecular orbital wave function is obtained by adding or subtracting atomic wave functions. A bonding orbital is an orbital whose wave function is obtained by adding the wave functions of atomic orbitals. It has less energy than an atomic orbital. An antibonding orbital is an orbital whose wave function is obtained by subtracting the wave functions of atomic orbitals. It has more energy than an atomic orbital. Figure 5.10. Scheme of the formation of molecular orbitals in the hydrogen molecule H2 The method of molecular orbitals is more complex than the method of valence bonds. However, its use makes it possible to explain some properties of substances that cannot be explained using valence bonds (for example, the paramagnetic properties of oxygen). Intermolecular interactions The properties of a substance depend not only on the properties of individual molecules, but also on their associations. In the gaseous state, the interaction forces between molecules are weak, so the properties of individual molecules are most important. In liquid and solid states, the interaction forces between molecules are of great importance. Universal intermolecular interaction forces are called van der Waals forces. 51 Copyright JSC Central Design Bureau BIBKOM & LLC Kniga-Service Agency They have 4 components: orientation forces (between polar molecules); inductive forces (between polar and non-polar molecules); dispersion forces; energy of repulsion of electron shells. The practical significance of intermolecular interactions is great. They influence the physical and chemical properties of substances. To determine them, approaches based on the laws of classical physics and theoretical quantum chemical methods are used. Questions for self-control 1. What is a chemical bond? 2. Give examples of substances with non-polar, polar and ionic bonds. 3. Explain the manifestation of the donor-acceptor mechanism of bond formation using the example of the chemical properties of ammonia. 4. Explain the geometry of molecules with SP3 - hybridization of the central atom (CH4, NH3, H2O). 5. Compare the properties of covalent and ionic bonds. 6. Substances with which type of chemical bond dissociate most easily? 7. How many π bonds are there in a sulfuric acid molecule? 8. How does the polarity of a chemical bond change in the series HCl –––––– NaCl? REFERENCES Main 1. Glinka, N.L. General chemistry / N.L. Glinka – M.: KNORUS, 2009. – 752 p. 2. Knyazev, D.A. Inorganic chemistry/D.A. Knyazev, S.N. Smarygin – M.: Bustard, 2004. – 592 p. 3. Ryazanova, G.E. General and inorganic chemistry. Tables and diagrams / G.E. Ryazanov. – Saratov: Federal State Educational Institution of Higher Professional Education “Saratov State Agrarian University”, 2006 – 284 p. Additional 1. Klinsky, G.D. Inorganic chemistry /G.D. Klinsky, V.D. Skopintsev. – M: Publishing house MCHA, 2001. – 384 p. 2. Gelfman, M.I. Inorganic chemistry / M.I. Gelfman, V.P. Yustratov. – St. Petersburg: Lan Publishing House, 2009. – 528 p. 3. http://www.xumuk.ru/encyclopedia/2/2994.html 4. ru.wikipedia.org 52 Copyright JSC Central Design Bureau BIBKOM & LLC Kniga-Service Agency Lecture 6 MODERN THEORY OF SOLUTIONS 6.1. Classification of dispersed systems A dispersed system is a system in which one substance is evenly distributed in the form of particles inside another substance. The dispersed phase (DP) is a finely divided substance. A dispersion medium (DS) is a substance in which the dispersed phase is distributed. Table 6.1 - Classification of dispersed systems by degree of dispersion Coarsely dispersed systems Colloidal systems True solutions The following definition of a solution is currently accepted: Solutions are homogeneous, thermodynamically stable systems of variable composition, consisting of a dissolved substance, a solvent and the products of their interaction. The solvent predominates quantitatively. The solute is present in less quantity. The dissolution of a substance is accompanied by thermal effects: Table 6.2 - Thermal effects of the dissolution process Endothermic process (heat absorption) - breaking of bonds ΔH1 > 0 Exothermic process (heat release) - formation of new bonds (solvation, hydration ΔH2< 0 общий тепловой эффект: ΔН = ΔН1 + ΔН2 53 Copyright ОАО «ЦКБ «БИБКОМ» & ООО «Aгентство Kнига-Cервис» Правило: Подобное растворяется в подобном. Полярные вещества – в полярных растворителях HCl – в H2O полярн. полярн. Неполярные вещества – в неполярных растворителях I2 – в бензоле неполярн. неполярн. 6.2. Способы выражения состава растворов Концентрация раствора – это величина, характеризующая содержание растворенного вещества в определенном массовом или объемном количестве раствора или растворителя. Таблица 6.3 – Способы выражения состава растворов 6.3. Концентрация почвенного раствора и осмос Осмос – это односторонняя диффузия молекул растворителя через полупроницаемые мембраны. Движущими силами осмоса являются: переход в состояние с более низким уровнем энергии; выравнивание концентраций по обе стороны мембраны; увеличение энтропии (неупорядоченности системы). Величина осмотического давления черноземов при орошении составляет 0,5-5,0 МПа, а солонцов в засушливых районах 5-17 МПа. 54 Copyright ОАО «ЦКБ «БИБКОМ» & ООО «Aгентство Kнига-Cервис» Таблица 6.4 – Осмотическое давление растительной клетки 6.4. Растворы электролитов 6.4.1. Водные растворы Важнейшим растворителем является вода. Водные растворы имеют огромное значение для живых организмов, в них проходят биохимические процессы. Распространение в природе и применение в промышленности и сельском хозяйстве. Таблица 6.5 – Биологическая роль воды Свойства воды 1. Молекула воды имеет угловую конфигурацию, так как имеет место SP3 – гибридизация электронных орбиталей атома кислорода 2. Электрический дипольный момент μ = 6,17∙10-30 Кл∙м (очень большой) 3. Диэлектрическая проницаемость = 81 (высокая) 4. Молекула воды полярна. 5. Степень ионности связи = 33% 6. Вода - уникальный растворитель. 55 Copyright ОАО «ЦКБ «БИБКОМ» & ООО «Aгентство Kнига-Cервис» 7. Молекулы воды образуют водородные связи; вступают в диполь – дипольные и ион-дипольные взаимодействия. В liquid state water molecules are associated: (H2O)2; (H2O)3; (H2O)4; (H2O)x. In the crystal lattice of ice, each oxygen atom is connected to the 4 oxygen atoms of 4 water molecules by four hydrogen bonds. There are cavities between water molecules. Figure 6.1. Information impact on a water crystal 6.4.2. Theory of electrolytic dissociation The concepts of “electrolytes” and “non-electrolytes” were already known in the 19th century. Electrolytes are substances whose solutions (or melts) do not conduct electric current. Nonelectrolytes are substances whose solutions (or melts) do not conduct electric current. For the general (colligative) properties of non-electrolyte solutions, the Raoult–van't Hoff law was discovered: a decrease in the saturated vapor pressure above the solution, an increase in the boiling point, a decrease in the freezing point is proportional to the molar concentration of the solution. It was found that electrolyte solutions do not obey the Raoult–van’t Hoff law. For electrolytes, all colligative properties are of greater importance. For example, for KCl ≈ 2 times, for BaCl2 ≈ 3 times. The problem has arisen: why is this so? In 1887, S. Arrhenius (Sweden) proposed the theory of electrolytic dissociation. Electrolytic dissociation is the decomposition of molecules into ions under the influence of a solvent (for aqueous solutions - under the influence of water). Ions are particles that have a charge. Cations have a charge (+). Anions have a charge (–). Electrolytes are substances that conduct current with their ions. 56 Copyright JSC Central Design Bureau BIBKOM & LLC Kniga-Service Agency The colligative properties of electrolyte solutions are greater because during dissociation the number of particles of the dissolved substance increases. HCl = H+ + Cl– 1 particle 2 particles Arrhenius was a proponent of the physical theory of solutions, which stated that the solute does not interact with the solvent. DI. Mendeleev created chemical theory solutions (1887). He proved that a solute and a solvent interact to form solvates (for aqueous solutions - hydrates). They are berthollides. I.A. Kablukov combined the physical theory of Arrhenius and the chemical theory of Mendeleev. The theory of electrolytic dissociation has risen to a new, higher level. The main cause of dissociation is solvation (hydration). Some ions are highly hydrated. Thus, the hydration energy of the hydrogen cation (proton) is high: H+: H+ + H2 O: = H3O+ acceptor donor hydronium cation There are no free H+ cations in an aqueous solution; they are converted into strong hydronium cations H3O+. 6.4.3. Quantitative characteristics of electrolytic dissociation There are two quantitative characteristics of electrolytic dissociation - the degree of dissociation and the dissociation constant. 1. The degree of electrolytic dissociation () is the ratio of the number of molecules disintegrated into ions to the total number of molecules of the dissolved substance. The degree of dissociation is measured in fractions of a unit or as a percentage: O 1 or O 100%. (%) = Diss 100%. Commun The degree of dissociation depends on: the nature of the dissolved substance (substances with ionic and highly polar types of chemical bonds dissociate best); on the nature of the solvent (on the value of its dielectric constant); on the concentration of the solution (when the solution is diluted, the degree of dissociation increases); on temperature; from the presence of ions of the same name (the degree of dissociation decreases). 57 Copyright JSC Central Design Bureau BIBKOM & LLC Kniga-Service Agency Table 6.6 – Dependence of electrolyte strength on the degree of dissociation and dissociation constant (Kd) Degree of dissociation () 30% 3% Electrolyte strength Kd 30%< 3% >10-2 Strong 10-2 – 10–4 Medium< 10–4 слабый 2. Константа диссоциации Кд находится на основании закона действия масс для слабых электролитов. КА ↔ К+ + А. Кд = . Кд не зависит от концентрации. 6.4.4. Свойства сильных электролитов Сильные электролиты в растворе диссоциированы полностью. Движение ионов стеснено притяжением друг к другу ионов с противоположными зарядами. Каждый ион окружен «ионной атмосферой», состоящей из противоположно заряженных ионов, которая тормозит его движение. Поэтому все свойства растворов электролита, зависящие от концентрации, проявляются не в соответствии с полной диссоциацией электролита, а в меньшей степени. Для оценки состояния ионов в растворе применяют величину, называемую активностью (а). Активность иона – это его условная, эффективная концентрация, в соответствии с которой он действует при химических реакциях. Активность иона (а) равна его концентрации (с), умноженной на коэффициент активности f: a = f∙C. Коэффициент активности иона зависит от его заряда и от ионной силы раствора I. Ионная сила раствора равна полусумме произведения концентраций всех находящихся в растворе ионов на квадрат их заряда: I 1 C1 Z 12 2 C 2 Z 22 ... C n Z n2 . Если вместо значений концентраций пользоваться значениями активности, то закон действия масс можно применить и к сильным электролитам. При этом можно получить значения констант диссоциации сильных кислот. Это дает возможность сравнивать свойства не только слабых, но и сильных электролитов (Kд>10–2 – strong; Kd<10–4 – слабые; Kд = 10–2 – 10–4 – средние). 58 Copyright ОАО «ЦКБ «БИБКОМ» & ООО «Aгентство Kнига-Cервис» Таблица 6.7 – Константы диссоциации некоторых электролитов при 25оС в водных растворах Слабые электролиты Сильные электролиты Электролит Кд Электролит Угольная кислота Н2СО3 К1 = 4,5∙10-7 Сероводородная кислота К1 = 6∙10-8 К2 = 4,7∙10 К2 = 1∙10 -11 -14 К1 = 8∙10-3 Ортофосфорная кислота К2 = 6∙10 -8 К3 = 1∙10-12 Кд Азотная кислота НNO3 Кд = 43,6 Хлороводородная кислота HCl 1∙107 Бромоводородная кислота HBr 1∙109 Йодоводородная кислота HJ 1∙1011 Упрощенно можно применить к электролитам более простую классификацию, разделив их на две группы – сильные и слабые. При этом все электролиты, не являющиеся сильными, относятся к слабым. 6.4.5. Типы электролитов Сильные электролиты диссоциируют практически полностью, необратимо. Слабые электролиты диссоциируют частично, обратимо. Таблица 6.8 – Сильные и слабые электролиты Кислоты HCl HBr Hl HNO3 H2SO4 HClO4 Основания Соли Гидроксиды Практически активных – все щелочных и (растворимые) щелочноземельных NaCl металлов KNO3 CuSO4 NaOH FeCl3 KOH Al2(SO4)3 Ca(OH)2 ZnSO4 Ba(OH)2 59 Кислоты Основания Вода H2CO3 H2SiO3 H2SO3 H2S HCN H3PO4 Органические кислоты Гидроксиды малоактивных металлов и гидроксид аммония Cu(OH)2↓ Fe(OH)3↓ Zn(OH)2 NH4OH(р-р) H2O Copyright ОАО «ЦКБ «БИБКОМ» & ООО «Aгентство Kнига-Cервис» 6.4.6. Диссоциация электролитов Таблица 6.9 - Диссоциация кислот, оснований и солей Диссоциация кислот Кислотность определяется присутствием ионов водорода Н+ Диссоциация оснований Щелочность среды определяется присутствием гидроксид-ионов ОН– Диссоциация солей 60 Copyright ОАО «ЦКБ «БИБКОМ» & ООО «Aгентство Kнига-Cервис» 6.4.7. Реакции в растворах электролитов Реакции в растворах электролитов идут в сторону образования прочных, устойчивых соединений, являющихся осадком, газом или слабым электролитом. При составлении молекулярно-ионных уравнений реакций сильные электролиты записывают в виде отдельных ионов, а слабые – в виде молекул. Пример: 1. Молекулярное уравнение реакции CuCl2 + 2NaOH = Cu(OH)2↓ + 2NaCl сильн. сильн. слаб. сильн. 2. Полное молекулярно-ионное уравнение реакции Cu2+ + 2Cl– + 2Na+ + 2OH– = Cu(OH)2↓ + 2Na+ + 2Cl– 3. Сокращенное молекулярно-ионное уравнение Cu2+ + 2OH– = Cu(OH)2↓ 6.4.8. Гидролиз солей Гидролиз – это обменное взаимодействие соли с водой. Гидролизу подвергаются соли, при взаимодействии которых с водой образуется хотя бы одно прочное, устойчивое соединение, являющееся слабым электролитом (или сложный, малодиссоциирующий ион). Гидролизу подвергаются соли, содержащие катион слабого основания или анион слабой кислоты или то и другое вместе (AlCl3, Na2CO3, Al2S3). Гидролизу не подвергаются соли, содержащие катионы сильного основания и анионы сильной кислоты (NaCl, K2SO4, KNO3). Периодический закон Д.И. Менделеева может быть применен к свойствам растворов. Периодический закон помогает понять взаимосвязь между свойствами атомов элементов и свойствами систем, содержащих ионы этих элементов. Так, например, изменение теплоты гидратации катионов металлов является функцией потенциалов ионизации атомов этих металлов. Чем больше заряд катиона металла и меньше радиус, тем более сильнее электрическое поле он создает, сильнее подвергается гидролизу. Поэтому гидролиз по катиону Al3+ идет, а гидролиз по катионам Na+ и K+ не идет. Алгоритм составления уравнений гидролиза 1. Составить уравнение диссоциации соли. Выявить ион, образующий слабое основание или слабую кислоту: AlCl3 = Al3+ + 3Cl- катион слабого основания 2. Составить сокращенное молекулярно-ионное уравнение по иону слабого электролита. Определить реакцию среды: Al3+ + HOH ↔ AlOH2+ + H+ кислая среда 61 Copyright ОАО «ЦКБ «БИБКОМ» & ООО «Aгентство Kнига-Cервис» 3. Составить молекулярное уравнение гидролиза: AlCl3 + HOH ↔ AlOHCl2 + HСl Чем слабее основание или кислота, образующие соль, тем сильнее идет гидролиз. Степень гидролиза (h) - это отношение концентрации соли, подвергшейся гидролизу (Сгидр.) к общей концентрации соли в растворе (Ссоли): h= С гидр С соли 100% . Константа гидролиза (Kг) выводится на основании константы равновесия (Kр) CN– + HOH ↔ HCN + OH– Kp = ; Кг = . Управление процессом гидролиза Усиливают гидролиз: разбавление; нагревание; удаление продуктов гидролиза (осадок, газ). Подавляет гидролиз добавление одного из продуктов гидролиза. Значение гидролиза для почвы Систематическое применение в качестве удобрения солей аммония увеличивает кислотность почвы: 1) NH4Cl = NH4+ + Cl–; слаб. 2) NH4+ + HOH ↔ NH4OH + H+; кислая среда 3) NH4Сl + HOH ↔ NH4OH + HCl. 6.4.9. Ионное произведение воды. Водородный показатель рН Диссоциация воды: H2O ↔ H+ + OH – слабый электролит Константа диссоциации: Kр = на основании закона действия масс [ H 2 O] Ионное произведение воды при 25оС: = 10–14 Для воды и разбавленных растворов электролитов произведение концентрации ионов водорода и гидроксид-ионов есть величина постоянная. 62 Copyright ОАО «ЦКБ «БИБКОМ» & ООО «Aгентство Kнига-Cервис» Из этого следует: 1. Нейтральная среда а) = [ОH–] 2. Кислая среда >10–7(10–6, 10–5, 10–3) b) ∙ = 10–14 c) 2 = 10–14 10 14 d) [H] 2 e) = 10–7 – in a neutral environment 3. Alkaline environment< 10–7(10–8, 10–9, 10–10) Таблица 6.10 – Реакция среды Среда , моль/л Нейтральная 10-7 Кислая >10-7 -6 (10 , 10-5, 10-3) Alkaline > 10–7 -8 (10 , 10-9, 10-10) The degree of acidity or alkalinity of a solution is characterized by the concentration of hydrogen ions. Hydrogen index The acidity or alkalinity of a solution can be characterized by the pH value. pH is the negative decimal logarithm of the concentration of hydrogen ions. pH = - log - molar concentration of hydrogen cations, or more precisely pH = - log a H a H - active concentration of hydrogen cations. = 10-pH Hydroxyl index The hydroxyl index pH is the negative decimal logarithm of the concentration of hydroxide ions. pOH = - log = 10–pOH Formulas for solving problems 1) ∙ = 10–14; 4) pH + pOH = 14; 2) pH = – log; 5) = 10–pH; 3) pOH = – log[ОH–]; 6) = 10–pH. + 63 Copyright JSC Central Design Bureau BIBKOM & LLC Book-Service Agency Figure 6.1. Scheme of changes in pH when the environment changes Table 6.11 - Value of acidity of the environment Soil Chernozem Podzolic Solonchak pH Plant Optim. pH Solution pH Potatoes 4.5-6.3 Blood 7.4 3.5-6 Cabbage 7.8-7.4 Gastric juice 1.7 8-9 Beetroot 7.0-7.5 Tap water 7.5 7 6.4.10. The importance of solutions Solutions are widespread in nature, used in technology, and are of great importance to living organisms. Table 6.12 - Electrolyte solutions in nature Natural waters of rivers, lakes, seas, oceans Soil solutions Biological fluids and tissues of living organisms Electrolyte solutions are used in technology to obtain valuable chemical compounds (acids, bases, salts), mineral fertilizers. In biology, the role of solutions is great because all cellular processes take place in aqueous solutions. Plants absorb macro- and micronutrients from aqueous soil solutions. The liquid allows you to create multifunctional controls and Information Systems . The prototypes of such a system are the living cell and the human brain. Questions for self-control 1. What is electrolytic dissociation? 2. Give examples of strong and weak electrolytes. 3. Why do not all salts undergo hydrolysis? 4. Calculate the pH and pOH of a 0.0001 mol/L KOH solution. 5. How to suppress the hydrolysis of FeCl3? 6. How to prepare 3 kg of 6% NaCl solution? 7. Why is iron sulfate FeSO4∙7H2O sometimes used for chemical reclamation of slightly alkaline soils? 8. Why is the use of microfertilizers CuSO4, FeSO4 ineffective on alkaline soils? 9. How can soil pH change with the systematic use of ammonium nitrate NH4NO3? 10. Why do plants wither when salt concentrations in the soil are too high? 64 Copyright JSC Central Design Bureau BIBKOM & LLC Book-Service Agency REFERENCES Main 1. Glinka, N.L. General chemistry / N.L. Glinka – M.: KNORUS, 2009. – 752 p. 2. Knyazev, D.A. Inorganic chemistry/D.A. Knyazev, S.N. Smarygin – M.: Bustard, 2004. – 592 p. 3. Ryazanova, G.E. General and inorganic chemistry. Tables and diagrams / G.E. Ryazanov. – Saratov: Federal State Educational Institution of Higher Professional Education “Saratov State Agrarian University”, 2006 – 284 p. Additional 1. Egorov, V.V. Environmental chemistry. /V.V. Egorov. – St. Petersburg: Lan Publishing House, 2009. – 192 p. 2. Ryazanova, G.E. Inorganic and analytical chemistry / G.E. Ryazanova - Saratov: Federal State Educational Institution of Higher Professional Education "Saratov State Agrarian University", 2006. - 172 p. 3. http://www.ecology-portal.ru/publ/1-1-0-124 65 Copyright JSC Central Design Bureau BIBKOM & LLC Kniga-Service Agency Lecture 7 REDOX REACTIONS 7.1. Modern theory of redox reactions (ORR) 7.1.1. The importance of redox processes Redox processes are among the most important processes in the world around us. The transformations of many elements and compounds in nature are associated with a chain of sequential redox reactions. Respiration, self-renewal of protein bodies, photosynthesis, absorption of food by animals, absorption of carbon dioxide and some macro- and microelements by plants, etc. – everything is based on redox reactions. Of great importance are the redox properties of the soil, in which a cycle of substance transformations constantly occurs, reminiscent of the metabolism in a living organism. The circulation in nature of the most important “elements of life” - nitrogen, sulfur, carbon - is associated with redox processes. Both in soil and in plants, these elements undergo redox transformations. If scientifically based recommendations for the use of nitrogen fertilizers are violated (fertilizer rates, application doses, timing, failure to take into account the dynamics of soil composition, temperature conditions, irrigation regulation), conditions are created for disruption of the ecological balance, poisoning of agricultural products with nitrates and nitrites due to the impossibility of normal reduction of nitrates to ammonia. Table 7.1 - The importance of ORR for plants Source of energy resources for living organisms Assimilation of certain nutrients by plants Formation of humus in the soil, increasing its fertility ORR Processes of transformation of nutrients in the soil (iron, manganese, nitrogen, sulfur) The cycle of “elements of life” in nature ( nitrogen, carbohydrate, sulfur) Redox reactions are of great technological importance. They are used to obtain metals in industry, to produce acids, phosphorus, halogens, hydrogen, etc. With the help of redox reactions, energy is obtained that is used in automobile, rocket, and aircraft technology. Many environmental problems of our time are associated with redox reactions, for example, atmospheric pollution with nitrogen and sulfur oxides, industrial waste and the possibility of converting them into substances that have beneficial applications. The study of redox processes creates a basis for understanding a number of complex issues in inorganic chemistry, organic, physical, biological, agronomic chemistry, soil science and special disciplines. Table 7.2 - Intrasubject and interdisciplinary connections theory of OVR 7.1.2. Basic provisions of the ORR theory Redox reactions are reactions in which the oxidation states of elements change due to the transfer of electrons from one atom to another. Reduction is the process of gaining electrons by an atom, ion or molecule of a substance. Oxidation is the process of losing electrons by an atom, ion or molecule of a substance. A reducing agent is a particle (atom, ion, molecule) that donates electrons. An oxidizing agent is a particle (atom, ion, molecule) that gains electrons. During reduction, the oxidation state of the element decreases. During oxidation, the oxidation state of the element increases. The redox process is the unity of two opposing processes. A reducing agent (oxidized) -e– – +e B oxidizing agent (reduced) The most important reducing agents 2 Metals, H2; N B r; H I ; NaBr; K I ; H2S2; Na2S2; FeSO4; CO. The most important oxidizing agents are Halogens (F2, Cl2, Br2, I2), oxygen O2; HNO3 (conc. and diluted), H2SO4 (conc.), KMnO4, K2Cr2O7. The redox properties of substances change periodically. 67 Copyright OJSC Central Design Bureau BIBKOM & LLC Kniga-Service Agency In the period with an increase in the serial number, the properties of elements change from metallic to non-metallic, therefore, from reducing to oxidizing. In the group from top to bottom, the metallic properties of the elements increase, and, consequently, the reducing properties (the ability to donate electrons). Electron balance rule The number of electrons given up by the reducing agent must be equal to the number of electrons accepted by the oxidizing agent. 7.1.3. Calculation of the oxidation state of an element Oxidation state The following concept of oxidation state is currently accepted: The oxidation state is the conditional charge of an element’s atom, calculated on the assumption that the molecule consists of ions. The following definitions of the oxidation state can be proposed: The oxidation state is the conditional charge of an atom of an element in a compound formally accepted as ionic. The oxidation number is the nominal charge of an atom in a hypothetical molecule consisting of ions. Using the concept of oxidation state, we want to evaluate how the charge of an element changes when its atoms give up or gain electrons. Quantitatively, the oxidation state is determined by the number of electrons given or accepted by an atom. Basic rules: 1. The oxidation state of an element in the atomic state is zero 2. The molecule is electrically neutral 3. The oxidation state of an element in molecules of simple substances is zero. A constant oxidation state in compounds is: H+ (with the exception of H 2 2 2 in Na + H hydrides); O (exclusion of O in O F 2); metals of group IA (Na+, K+); metals of group IIA (Ca+2, Ba+2), aluminum Al+3 (III A-group). Oxidation state: 2 metals can only be positive: Na 2 O, C aO; 2 6 of a nonmetal can be either positive or negative: H 2 S, S O 3; more electronegative nonmetal in a compound of two nonmetals – 6 2 negative: S O 3 ; is equal to zero if the number of electrons in the atom is equal to the charge of the nucleus; positive if the number of electrons in the atom is less than the charge of the nucleus; negative if the number of electrons is greater than the charge of the nucleus; The magnitude and sign of the oxidation state depend on the structure of the atom. 68 Copyright JSC Central Design Bureau BIBKOM & LLC Kniga-Service Agency Algorithm for calculating the unknown oxidation state in a molecule 1. Designate the known oxidation states of elements; 2. Create an equation for the algebraic sum of the oxidation states of all elements, which is equal to zero (taking into account the number of atoms of each element); 3. Find the unknown oxidation state by solving the equation. Example: Find the oxidation number of Mn in KMnO4 Solution: KMnO4 +1 + x + 4(-2) = 0 x = +8 –1 = +7 Answer: The oxidation number of manganese in KMnO4 is + 7. To find the oxidation number in a complex ion , you need to solve an algebraic equation. Example: Find the oxidation state of nitrogen in N O3 X 2 Solution: x + 3(-2) = -1 (N O3) x = -1 + 6 = +5 Answer: The oxidation state of nitrogen in NO 3 is +5. Electronic equations Electronic equations reflect the processes of oxidation of a reducing agent due to the donation of electrons to it and reduction of the oxidizing agent due to the addition of electrons. To compile an electronic equation, the oxidation state of an element in its initial and final states must be compared. Table 7.3 – Basic concepts for composing electronic equations 0 Examples: 1. S 2e 2 0 S 1. Na e 3 2 4 2. Fe e Fe 2. S 2e 6 3. S 8e 2 2 S 3. S 8e 69 1 Na 6 S 6 S Copyright JSC Central Design Bureau BIBKOM & LLC Book-Service Agency 7.1.4. Algorithm for characterizing the redox properties of complex compounds 1. Calculate the oxidation state of the element that determines the redox properties of the substance molecule; 2. Assess the electronic structure of the atom in a given oxidation state, the presence of valence electrons; 3. Draw conclusions about the possibility of manifestation of reducing or oxidizing properties. Table 7.4 – Analysis of the redox properties of complex compounds Electronic structure of an atomic particle Maximum positive oxidation state 5 H N O3 No valence electrons N+5)0 ē 1s22s02p0 NH 3 3 External electronic level completed N 3)8 ē 1s22s22p8 3 K N O2 Presence of valence electrons N+3)2 ē 1s22s22p0 Properties HNO3 - only oxidizing negative NH3 - only reducing intermediate KNO2 and reducing and oxidizing 7. 1.5. Types of redox reactions Intermolecular oxidation-reduction reactions Intramolecular oxidation-reduction reactions The reducing agent and the oxidizing agent are in disparity The oxidizing agent and the reducing agent are constituents of the same molecule Disproportionation reactions (self-oxidation-self-reduction) The oxidizing agent and the reducing agent are the same element Reactions comporportionation (averaging) Reductant and oxidizing agent - atoms of an element with different degrees of oxidation are part of different molecules 70 Copyright OJSC Central Design Bureau BIBKOM & LLC Kniga-Service Agency 7.2. Algorithms for composing equations for redox reactions 7.2.1. Algorithm for composing the ORR equation using the electronic balance method 1. Draw up a reaction scheme (write the formulas of the starting substances and reaction products; 2. Indicate the oxidation states of elements that changed during the reaction; 3. Draw up electronic equations reflecting the processes of oxidation of the reducing agent and reduction of the oxidizing agent; 4 Based on the electronic balance rule, find the coefficients for the reducing agent and the oxidizing agent; 5. Place the found coefficients on the left and right sides of the reaction equation in front of the corresponding substances; 6. Equalize for the metal that has not changed the oxidation state; 7. Equalize for the non-metal of the acid residue; 8. Equalize for hydrogen; 9. Check for oxygen. Examples: 7.2.2. Drawing up the ORR equation using the electron-ion method (half-reaction method) The half-reaction method is used to find the coefficients of redox reactions occurring in solutions. In this method, the charges of oxidizing and reducing ions are indicated , and not the oxidation states of the elements. Rule of equality of charges The sums of charges on the left and right sides of the electron-ion equation of the redox process must be equal. Algorithm for composing the ORR equation using the electron-ionic method (half-reaction method) 1. Draw up a molecular scheme of the reaction: KI + KMnO4 + H2SO4 → I2 + MnSO4 + K2SO4 + H2O 2. Determine the oxidizing agent, reducing agent and products of their transformation: KI + KMnO4 + H2SO4 → I2 + MnSO4 + K2SO4 + H2O medium 71 Copyright JSC Central Design Bureau BIBKOM & LLC Kniga-Service Agency 3. Draw up electron-ion equations (half-reactions) for oxidation and reduction processes. Check the sum of the charges on the left and right sides of the equation. If the sum of the charges of the original particles more than the amount charges of the reaction products, then add the corresponding number of electrons to the left side of the equation (if less, subtract). Example: 4. Find the coefficients for the reducing agent, the oxidizing agent and the products of their transformation: 5. Put the coefficients in the molecular equation diagram: 6. Equalize for the metal that has not changed the charge (potassium). 7. Check for acid residues. 8. Check for hydrogen. 9. Check for oxygen. Table 7.5 – Scheme for the absorption of nitrates by plants Electronic equations Electron-ionic equations 72 Copyright JSC Central Design Bureau BIBKOM & LLC Kniga-Service Agency Factors influencing redox reactions 1. Nature of the reducing agent and oxidizing agent 2. Temperature. 3. Wednesday. 4. Concentration of oxidizing agent and reducing agent. 7.3. Redox properties of compounds of biogenic elements The most important biogenic elements - nitrogen, iron, manganese, sulfur, etc. exhibit redox properties during reactions in the living cell of a plant organism and in the soil. Of great importance in the life of plants are reactions associated with redox transformations of the “metals of life” - iron, manganese, copper, molybdenum, cobalt. Thus, iron is part of the enzymes involved in the synthesis of chlorophyll, which plays an important role in respiration and energy metabolism of cells; manganese is involved in the reduction of nitrates, molybdenum in the process of nitrogen fixation, etc. The soil contains redox systems Fe3+/Fe2+, manganese compounds in oxidation states +2, +3, +4. If there is a lack of manganese in the soil, manganese fertilizers are used - MnSO4 and KMnO4. Properties of potassium permanganate KMnO4 7 K Mn O4 – strong oxidizing agent 7 Mn)0 – oxidizing agent, no valence electrons (only accepts electrons) 0 7 Mn)3d 5 4s 2 7å Mn)3d 0 4s 0 7 The role of the medium for the oxidizing properties of KMnO4 (Mn ) The course of the redox reaction depends on the environment. You can control the redox reaction by measuring the pH of the environment. The concentration (activity) of hydrogen cations H+ is important for such reactions that occur with their participation. Hydrogen ions take part in reactions with an oxygen-containing oxidizing agent (KMnO4, K2Cr2O7). They take oxygen away from the oxidizing agent, facilitating the reduction process. Figure 7.1. Properties of potassium permanganate KMnO4 in various environments An acidic environment, in which the H+ concentration is high, is the most favorable 7 for reactions involving KMnO4. Under these conditions, Mn is reduced most deeply, and the largest number of electrons is involved in the reduction process. 73 Copyright JSC Central Design Bureau BIBKOM & LLC Kniga-Service Agency Acidic environment (favorable, a lot of H+) Neutral environment (less favorable, less H+) Alkaline environment (unfavorable) 7.4. Redox potentials. Direction of ORR The redox properties of a substance are characterized by the value of the standard redox potential E0. The standard electrode redox potential (E0) is the potential of a redox system measured relative to a standard hydrogen electrode, provided that the ratio of the activities of the oxidized and reduced forms is equal to unity. The higher the E0 value, the stronger the oxidizing agent the substance is. A process occurs spontaneously with the participation of a stronger oxidizing agent, that is, at 0 EMF = Eok Evost > 0 Changing conditions affects the direction of the process. 74 Copyright JSC Central Design Bureau BIBKOM & LLC Kniga-Service Agency Example: HClO4 + Br2 + H2O ↔ HClO3 + HBrO3 stronger oxidizing agent The reaction occurs with the participation of a stronger oxidizing agent Nernst equation RT a ok a H ln E=E + nF a restore 0 or m 0.059 a ok a H E=E + . ln n a restore 0 An increase in the concentration of the oxidizer increases E. An increase in the concentration of H+, a change in the pH of the medium can change the direction of the reaction (if Eok and Evost have close values). Questions for self-control 1. What is reduction, oxidation, reducing agent, oxidizing agent? 2. Calculate the oxidation state of Mn in KMnO4. 3. Compare the algorithms for finding coefficients for the exchange reaction and for the redox reaction. 4. Explain the difference in redox properties for NH3 and HNO3. 5. Find the coefficients for the reaction KMnO4 + FeSO4 + H2SO4 → MnSO4 + Fe2(SO4)3 + K2SO4 + H2O 6. Compose electronic (or electron-ion) equations corresponding to the uptake of sulfur by plants: SO 42 → SO 23 → S 2 REFERENCES Main 1. Glinka, N.L. General chemistry / N.L. Glinka – M.: KNORUS, 2009. – 752 p. 2. Knyazev, D.A. Inorganic chemistry/D.A. Knyazev, S.N. Smarygin – M.: Bustard, 2004. – 592 p. 3. Ryazanova, G.E. General and inorganic chemistry. Tables and diagrams / G.E. Ryazanov. – Saratov: Federal State Educational Institution of Higher Professional Education “Saratov State Agrarian University”, 2006 – 284 p. 4. Ryazanova, G.E. Inorganic and analytical chemistry / G.E. Ryazanova - Saratov: Federal State Educational Institution of Higher Professional Education "Saratov State Agrarian University", 2006. - 172 p. Additional 1. Inorganic chemistry (biogenic and abiogenic elements): Textbook / edited by prof. V.V. Egorova. – St. Petersburg: Lan Publishing House, 2009. – 320 p. 2. Klinsky, G.D. Inorganic chemistry /G.D. Klinsky, V.D. Skopintsev. – M: Publishing house MCHA, 2001. – 384 p. 75 Copyright JSC Central Design Bureau BIBKOM & LLC Kniga-Service Agency Lecture 8 COMPLEX CONNECTIONS 8.1. Brief history of complex compounds (CS) The history of complex compounds can be divided into four periods. I period (from ancient times to the beginning of the 18th century) Application of complex compounds: precious stones (emerald, garnet, turquoise); natural dyes (indigo, purple, saffron); cleaning the surface of metals for soldering; tanning leather; receiving ink. II period (beginning of the 18th century until 1893) Synthesis of complex compounds; An attempt to explain the structure and properties of CS. III period (1893-1940) Creation, justification and victory of the coordination theory of Alfred Werner (Switzerland). IV period (from 1940 to the present) Development of coordination theory using the achievements of chemistry, physics and mathematics. In Russia, the chemistry of complex compounds developed in an original way: the work was ahead of European achievements, but was little known in the scientific world. Also M.V. Lomonosov studied the dissolution of salts in saturated solutions, and P.B. Bagration (the commander's nephew) discovered the reaction for the industrial production of gold. Yu.V. Lermontova (the poet’s second cousin) carried out the separation of platinum group metals. DI. Mendeleev, L.A. Chugaev, N.S. Kurnakov introduced the chemistry of complex compounds modern forms . I.I. Chernyaev, A.A. Grinberg, V.V. Lebedinsky, K.B. Yatsimirsky carried out further development of the theory of complex compounds. M.V. Lomonosov P.B. Bagration 76 Yu.V. Lermontov Copyright JSC Central Design Bureau BIBKOM & LLC Book-Service Agency D.I. Mendeleev L.A. Chugaev N.S. Kurnakov I.I. Chernyaev V.V. Lebedinsky A.A. Greenberg Problematic situation - by the middle of the 19th century, information about unusual chemical reactions and compounds had been accumulated. Example 1: Why does AgCl, insoluble in water, easily dissolve in ammonia NH3? AgCl + 2 NH3 white precipitate Cl solution The reaction was discovered by I. Glauber in 1648. Example 2: Why does CuSO4 react by combining with ammonia NH3? CuSO4 + 4NH3 SO4 blue solution blue-violet solution The reaction was discovered by Andrei Libavi in 1597. Butlerov A.M. believed that “Facts that are inexplicable by existing theories are the most precious for science; their development should primarily be expected to develop in the near future.” 8.2. Werner's coordination theory and modern ideas The word “complex” (Latin) is complex; combination. In 1893, Swiss scientist Alfred Werner proposed coordination theory to explain the formation of complex compounds. This theory is the main one in the chemistry of complex compounds at the present time. Over time, ideas about the forces acting between particles forming complex compounds change and become more precise. Questions explaining the properties of complex compounds are very complex. 77 Copyright JSC Central Design Bureau BIBKOM & LLC Kniga-Service Agency Werner introduced the concept of main and secondary valency, but did not explain their reason and difference. Currently, various theories are used to explain the structure and properties of complex compounds: the method of molecular orbitals; crystal field theory; ligand field theory; valence bond method (VBC). The method of valence bonds provides qualitative answers to questions within the framework of visual representations. From the point of view of MBC: the main valence is explained by the formation of ionic (or covalent) bonds; secondary valence - a covalent bond formed by a donor-acceptor mechanism. There are various definitions of complex compounds. Let us present one of them, based on the valence bond method. Complex compounds (CCs) are complex compounds of higher order, characterized by the presence of covalent bonds that arise through a donor-acceptor mechanism. 8.2.1. Composition of molecules of complex compounds 1. The complexing agent occupies a central place in the molecule of a complex compound. The most common complexing agents are metal cations. Strong complexing agents: 3 2 0 2 2 cations and d-metal atoms of side subgroups (Fe, Fe, Fe, Cu, Ag, Zn) 3 2 2 p-metal cations (Al, Sn, Pb). Poor complexing agents are cations of alkali and alkaline earth metals. However, in living organisms they participate in the formation of CS. 2. The complexing agent coordinates LIGANDS around itself (ions of opposite charge or neutral molecules) Anions: F-(fluoro-), Cl– (chloro-), NO2– (nitro-), CN– (cyano-), SO32– (sulphito- ), OH– (hydroxo-) Neutral molecules: NH3 (ammin-), H2O (aqua-), etc. 3. The coordination number of a complexing agent (CCN) is the number of ligands associated with the complexing agent. The coordination number of the complexing agent is usually twice the oxidation number of the complexing agent. 78 Copyright JSC Central Design Bureau BIBKOM & LLC Kniga-Service Agency Table 8.1 – Dependence of the coordination number on the degree of oxidation of the complexing agent Oxidation degree of the complexing agent KCC +1 (Ag) 2 2 2 4 (3.6) +2 (Cu, Fe ) 3 3 6 (4) +3 (Fe, Al) 4. The complexing agent, together with the ligands, forms the internal coordination sphere of the complex compound, or a complex ion. When writing a formula, it is enclosed in square brackets. The charge (oxidation state) of the inner sphere is equal to the algebraic sum of the oxidation states of the complexing agent and ligands. 5. The internal sphere of the CS is connected with its external sphere. The composition of the outer sphere, depending on the charge of the complex ion, can include both anions of acidic residues and metal cations (for example, alkali). The charge on the outer sphere is equal in magnitude and opposite in sign to the charge on the inner sphere. Table 8.2 - Examples of compiling formulas of complex compounds Example 1. Compose the formula of tetraammine copper (II) chloride 2 [Cu (NH 3) 04 ] complexing agent ligands Example 2: Compose the formula of sodium tetrahydroxycinate 2 2 Cl 2 Na 2 [ Zn (OH) 4 ] 2 external complex ligands collaborator external environment CCC CCC internal coordination sphere internal coordination sphere 8.2.2. Nomenclature of complex compounds When composing the name of a complex compound, you should: 1. First name the anion. 2. In the name of the internal coordination sphere, list from right to left all its components. 2.1. Name the number of ligands using Greek numerals (1 – mono, 2 – di, 3 – three, 4 – tetra, 5 – penta, 6 – hexa) 2.2. Name the ligands (first anions, then molecules) F–(fluoro-), Cl–(chloro-), Br–(bromo-), CN–(cyano-), H2O(aqua-), NH3(ammin-) . 2.3. Name the complexing agent. Indicate its oxidation state using Roman numerals in parentheses. If the inner coordination sphere is a (+) cation, use Russian name complexing agent in the genitive case. 79 Copyright JSC Central Design Bureau BIBKOM & LLC Kniga-Service Agency Example: +Cl– diammine silver chloride (I) If the internal coordination sphere is an anion (–), use Latin name complexing agent with the ending “at”. K3 –3 potassium hexacyanoferrate (III) 8.3. Chemical bonding in complex compounds Chemical bonding in complex compounds is a complex scientific problem. To explain the formation, structure and properties of complex compounds, several theories are used - ligand field theory, crystal field theory, and the valence bond method. Crystal field theory is applied to ionic complexes. It takes into account electrostatic forces and does not take into account covalent forces. The ligand field theory assumes that the formation of complex compounds occurs both due to electrostatic forces and due to covalent interactions when the orbitals of the central ion and ligands overlap. The valence bond method is based on the assumption that chemical bonding in complex compounds occurs through lone electron pairs of ligands and free orbitals of the complexing agent. To clarify the qualitative side of issues, all three theories are sometimes used. The theory of chemical bonds in complex compounds requires further development and improvement. The valence bond method (MVM) is the most visual. According to MBC, the chemical bond between the complexing agent and the ligands occurs via a donor-acceptor mechanism. Example: How does the 3+ coordination sphere form? 1) Formation of complexing ion Al3+: 2). The oxygen atom in a water molecule has non-bonding electron pairs and can play the role of a donor. 80 Copyright JSC Central Design Bureau BIBKOM & LLC Kniga-Service Agency Electron pairs are oriented in space at an angle of 90° (four in one plane, two more perpendicular to it). The resulting coordination sphere has the shape of an octahedron Figure 8.1. Scheme of ion formation 3+ CN = 2 sp CN = 3 sp2 Dumbbell Triangle Square –3 – 2- sp2d CN = 4 sp3 CN = 5 sp3d CN = 6 sp3d2 Tetrahedron Trigonal bipyramid Octahedron +2 2+ sd3 Figure 8.2. Geometry of coordination spheres of complexes 8.4. Electrolytic dissociation of complex compounds Complex compounds that are electrolytes dissociate in two stages. First stage. Primary dissociation is the dissociation of the CS into internal and external spheres. It follows the type of strong electrolytes, since the inner and outer spheres of the CS are connected by an ionic bond. K3 = 3K+ + 3– Second stage. Secondary dissociation is the dissociation of the inner sphere of the CS (complex ion). It is of the type of weak electrolytes, insignificant. 3- ↔ Fe3+ + 6CN– Stability of complex compounds. The stability of complex compounds is characterized by the instability constant of the complex ion Kn. The instability constant Kn of a complex compound is derived based on the law of mass action. Kn is the ratio of the product of the concentration of ions formed during the dissociation of a complex ion, to a degree equal to the coefficients in the dissociation equation of a complex ion, to the concentration of the complex ion. 81 Copyright OJSC Central Design Bureau BIBKOM & LLC Kniga-Service Agency Example: Create expressions for the instability constant for tetraammine copper (II) sulfate SO4. Primary dissociation: SO4 = 2+ + SO42– Secondary dissociation: 2+ ↔ Cu2+ + 4NH3 4 Kn = The smaller the Kn value, the more stable the complex ion is. Table 8.3 – Instability constants of some complex ions Complex ion + 2+ 2+ 34– 2– – Кн 7.2∙10–8 3.5∙10–10 2.1∙10–13 1.0∙10–44 1 .0∙10–37 2.4∙10–6 1.0∙10–21 Extremely unstable complex compounds are double salts: (K2SO4∙Al2(SO4)3) – potassium alum 2KCl∙CuCl2∙2H2O – double chloride (NH4 )2SO4∙Fe2(SO4)3∙24H2O – Mohr’s salt or NH4 Fe(SO4)2∙12H2O In the form of complex compounds, double salts exist only in the solid state. In solution they completely dissociate. Example: Dissociation of Mohr's salt NH4Fe(SO4)2 = NH4+ + Fe3+ + 2SO42. Exchange reactions involving complex compounds Preparation Prussian blue (Prussian blue) F. Diesbach (1704) 3 1 4 Fe Cl 3 3K 4 1 4 Fe 4 3 3 4 12KCl . The meaning and application of complex compounds Complex compounds perform important functions in living organisms and are used in industry, agriculture, and to solve environmental problems. 82 Copyright JSC Central Design Bureau BIBKOM & LLC Kniga-Service Agency Table 8.4 - The meaning of complex compounds Complex compounds For the production of dense and durable metal coatings For the separation of rare earth elements Perform specific functions in living organisms (respiration, photosynthesis, enzymatic catalysis) In analytical chemistry - for the determination of many elements In agriculture - to combat calcareous chlorosis of plants, to dissolve iron carbonates and hydroxides, to eliminate zinc deficiency To solve environmental problems, to obtain easily neutralized complexing agents Hemoglobin Carries out the transfer of oxygen from the lungs to the tissues. It has a unique significance for the life of the human body. Chlorophyll An intracomplex compound of magnesium ion with complex organic ligands. It has a unique significance for plants (for photosynthesis). Figure 21. Scheme of manifestation of the donor properties of nitrogen in metal-containing biomolecules Life metals (iron, zinc, molybdenum, copper, manganese) are complexing agents in the molecules of many enzymes that catalyze redox processes: peroxidase, catalase (iron); carboxypeptidase (zinc); xanthine oxidase (molybdenum and iron). Currently, a new direction of chemistry has emerged - bioinorganic chemistry, which explores the essence of such important manifestations of life as metabolism, heredity, immunity, thinking, memory. 83 Copyright JSC Central Design Bureau BIBKOM & LLC Kniga-Service Agency The molecules of many enzymes are built according to the type of intra-complex compounds, in which the role of complexing agents is played by “life metals” - iron, copper, manganese, molybdenum, cobalt. The specificity of enzyme activity depends on which metal is the complexing agent in protein molecules. At the same time, the role of complex compounds in the mineral nutrition of plants is being studied. It has been established that the absorption of microelements and the fixation of atmospheric nitrogen occurs with the participation of complexing agents. In 1964 M.E. Volpin and V.B. Schur made a discovery at the Institute of Organoelement Compounds. They turned molecular nitrogen into a complex compound under normal conditions, without heating. It was found that in the presence of compounds of iron, molybdenum and vanadium, nitrogen is activated and behaves as a ligand, forming complex compounds that are decomposed by water to ammonia. Questions for self-control 1. The concept of a complex compound. 2. Write a formula for a complex compound if the complexing agent is 3+ Al and the ligand is OH ions. 3. Complete the reaction equation: CuSO4 + NH3 → 4. Write the reaction equation: FeCl3 + K4 → 5. Determine the coordination number and oxidation state of the complexing agent in SO4 compounds. 6. Write an expression for the instability constant of the complex compound K3. REFERENCES Main 1. Glinka, N.L. General chemistry / N.L. Glinka – M.: KNORUS, 2009. – 752 p. 2. Knyazev, D.A. Inorganic chemistry/D.A. Knyazev, S.N. Smarygin – M.: Bustard, 2004. – 592 p. 3. Ryazanova, G.E. General and inorganic chemistry. Tables and diagrams / G.E. Ryazanov. – Saratov: Federal State Educational Institution of Higher Professional Education “Saratov State Agrarian University”, 2006 – 284 p. Additional 1. Inorganic chemistry (biogenic and abiogenic elements): Textbook / edited by prof. V.V. Egorova. – St. Petersburg: Lan Publishing House, 2009. – 320 p. 2. Ugay, Ya.A. General and inorganic chemistry /Ya.A. Ugai. – M.: Higher School, 2004. – 528 p. 3. Lensky, A.S. Introduction to bioinorganic and biophysical chemistry / A.S. Lensky. – M.: Higher School, 1989. – 256 p. 84 Copyright JSC Central Design Bureau BIBKOM & LLC Kniga-Service Agency Lecture 9 CHEMICAL KINETICS. CHEMICAL EQUILIBRIUM 9.1. The concept of chemical kinetics Chemical kinetics is the study of the rates and mechanisms of chemical reactions. Chemical kinetics studies the course of a reaction over time, studies the factors influencing the rate of a chemical reaction, and provides information about the reaction mechanism. Most chemical reactions go through a number of intermediate stages. The intermediate stages are called elementary stages. The reaction mechanism is the complete sequence of all its elementary stages. The rate of a reaction is determined by the rate of the slowest (limiting) stage of the reaction. Reaction mechanism for the synthesis of hydrogen iodide Stage 1 Stage 2 Stage 3 Stage 4 (fast) (slow) Total reaction equation: H2 + I2 2HI Reaction of sodium thiosulfate with sulfuric acid Total reaction equation Na2S2O3 + H2SO4 = Na2SO4 + SO2 + S + H2O Stage 1 Na2S2O3 + H2SO4 = H2S2O3 + Na2SO4 (fast) Stage 2 H2S2O3 = H2SO3 + S ↓ (slow) Stage 3 H2SO3 = SO2 + H2O (fast) 9.2. Rate of a chemical reaction The rate of a chemical reaction is the number of active collisions between molecules leading to the formation of a reaction product per unit volume or per unit surface per unit time. The rate of a chemical reaction is measured by the change in the concentration of one of the reacting substances per unit of time. The concepts of average and instantaneous reaction rates are used. The average reaction rate is determined by the change in concentration over a given period of time: 85 Copyright JSC Central Design Bureau BIBKOM & LLC Kniga-Service Agency ΔС Δt Vav The true (instantaneous) value of the reaction rate is calculated as the limit to which the average speed tends at Δt → 0 , i.e. as a derivative of concentration over time: dС dt Vist. The sign + or – depends on whether the change in the concentration of which substance - the original or the product - is used for calculations. If the concentration of the reacting substance is used, then the reaction rate will have a minus sign (–), since its concentration decreases during the reaction. Rate of chemical reaction Concentration C, mol/l Figure 9.1. Dependence of changes in concentration on reaction time 9.3. Factors influencing the rate of a chemical reaction Scheme 9.1. The influence of kinetic factors on the rate of a chemical reaction 86 Copyright JSC Central Design Bureau BIBKOM & LLC Kniga-Service Agency Table 9.1 – The importance of the nature of the reacting substances for the possibility and activity of the reaction Au + HCl = does not occur Zn + 2HCl = ZnCl2 + H2 Reactivity is the ability of molecules of a substance to participate in reactions at a certain rate. 9.3.1. The influence of concentration on the reaction rate Law of mass action (K. Guldberg and P. Waage (1864-1867 Norway): The rate of the elementary stage of a chemical reaction (V) is proportional to the product of the molar concentrations of the reacting substances (C) raised to the power of their stoichiometric coefficients. aA + bB = cC + dD; V = KC àÀ C b , where K is the reaction rate constant. The rate constant depends on the nature of the reactants and on temperature, but does not depend on their concentration. The rate constant is numerically equal to the reaction rate if the concentration of each of the reacting substances is equal to 1 mol/dm3. For heterogeneous reactions, the concentration of the solid substance is not included in the kinetic equation, since the surface area is practically constant, and substances react only on the surface. The region of manifestation of the law of mass action. The law of mass action is valid for a reaction occurring in one stage or for each individual stage of a multi-stage reaction.Applying the law of mass action to calculate the rate of a reaction without taking into account the multi-stage mechanism will lead to erroneous results. Examples 1. Homogeneous reaction 4HCl(g) + O2(g) 2H2O(g) + 2Cl2(g) 4 For the forward reaction 1 = K1C HCl CO 2 For the reverse reaction 2 = K2C 2H 2O C Cl2 2 2. Heterogeneous reaction Fe3O4(s) + H2(g) 3FeO(s) + H2O(g) For the forward reaction V1 = K1 CH 2(r) For the reverse reaction V2 = K2 C H2O(g) 87 Copyright JSC Central Design Bureau BIBKOM & Book-Service Agency LLC 9.3.2. Effect of temperature on the reaction rate 1. Van't Hoff's rule For every 10 degrees the temperature increases, the reaction rate increases by 2-4 times: Vt 2 = Vt 1 t1 t 2 10 Vt 2 ; Vt1 Δt 10 γ, where is the temperature coefficient, taking values from 2 to 4. Van't Hoff's rule is approximate. It is applicable for an approximate assessment of the effect of temperature changes on changes in reaction rates in a small temperature range. 2. To more accurately determine the effect of temperature on the rate of a chemical reaction in a wide temperature range, the Arrhenius equation is used: K = Ae Ea /RT, where K is the rate constant; A – constant; e – base of natural logarithm; Ea – activation energy; R – gas constant; T – absolute temperature. The Swedish scientist Svante Arrhenius created the theory of electrolytic dissociation in 1887, and in 1889 he proposed a new measure of the reactivity of compounds - activation energy. The equation he introduced into chemical kinetics has been used for more than 100 years to study the rates and mechanisms of reactions. The main provisions of the activation theory of S. Arrhenius: The reaction rate depends on the number of active molecules, and not on their total number. Only active molecules react. Activation energy Ea is the energy that must be imparted to molecules in order to transfer them to the active state (for 1 mole of a substance). The lower the activation energy, the greater the reaction rate. In order for the starting materials to form reaction products, it is necessary to overcome the energy barrier Ea. To do this, the starting substances need to absorb the energy necessary to break existing chemical bonds. New chemical bonds do not form immediately. First, an intermediate transition state (activated complex) with the maximum energy of the system is formed. Then it breaks down and molecules of the reaction product are formed with new chemical bonds, and energy is released. A reaction is exothermic if the amount of energy released is greater than the activation energy. In the presence of a catalyst, the activation energy decreases and the reaction rate increases. 88 Copyright JSC Central Design Bureau BIBKOM & LLC Book-Service Agency Figure 9.2. Energy diagram of the reaction 9.3.3. The influence of catalysts on the rate of a chemical reaction Catalysts are substances that affect the rate of a reaction, but are not consumed during its course. Positive catalysts increase the reaction rate because they reduce the activation energy. Negative catalysts (inhibitors) reduce the rate of a chemical reaction. Catalysis is a change in the rate of chemical reactions under the influence of catalysts. Catalytic reactions are reactions that occur under the influence of catalysts. Catalytic reactions play an important role in chemical production and biochemistry. 90% of all productions chemical industry catalysts are used (production of ammonia, nitric acid, sulfuric acid, rubber, etc.). Catalysts are used in fuel combustion and wastewater treatment. High-tech ammonia plants use less energy and produce less waste than traditional plants. Catalysis can be homogeneous or heterogeneous. In homogeneous catalysis, the reagents and the catalyst are in the same state of aggregation. In heterogeneous catalysis, a solid catalyst is used, and the reagents are part of a solution or gas mixture. In the presence of a catalyst, the reaction proceeds faster because the activated complex formed with its participation has a lower activation energy than in the absence of a catalyst. Currently, new catalysts are being developed using nanocompositions, semiconductor materials, and complex compounds. Much attention is paid to nanocomposite catalysts containing highly dispersed metals or metal oxides stabilized in a zeolite matrix having an ordered three-dimensional structure of pores (channels) of molecular size (0.3 - 1.2 nm). They have a wide range of applications due to their catalytic selectivity, acid resistance and thermal stability. 89 Copyright JSC Central Design Bureau BIBKOM & LLC Kniga-Service Agency Catalysis takes place in biological systems. Most chemical reactions occurring in living organisms are catalytic reactions. Biological catalysts - ENZYMES - participate in protein metabolism processes, catalyze the conversion of starch into sugar, and accelerate redox reactions. “Comparison with biocatalysis shows that, despite great successes, we are still far from exploiting all the possibilities opened up in this regard in catalysis. This teaches modesty” (S.Z. Roginsky). It is necessary to store crop products at low temperatures, as this slows down unwanted enzymatic processes. Many substances used in seed dressing disable enzymes and kill unwanted microorganisms. Chemical kinetics explains the mechanisms of reactions in a living cell. Genetic engineering methods are methods for controlling the rates of biochemical reactions. 9.4. The concept of oscillatory reactions Oscillatory reactions are reactions during which the concentration of intermediate compounds and the rate of the reaction undergo fluctuations, which can be periodic. Nobel laureate Ilya Prigogine called the discovery of oscillatory reactions a scientific feat of the 20th century. The feat was that in the twentieth century it was possible to prove that oscillatory reactions exist, since before that they had been rejected. Information about oscillatory reactions is in the archives of the 17th–19th centuries. Book “Theory of Oscillations” by A.A. Andronova, A.A. Witta and S.E. Khaikina was published in 1937. YES. Frank-Kamenetsky published the book “Diffusion and Heat Transfer in Chemical Kinetics” in 1947. These works were innovative, but they were not accepted and were consigned to oblivion. In 1951 B.P. Belousov discovered and described a visual oscillatory reaction, proposed ideas about the key moments of its mechanism, but it was rejected. After 10 years, the study of oscillatory reactions was continued by a group of scientists led by A.M. Jabotinsky. It has been proven experimentally and mathematically that oscillatory reactions are possible. The Belousov-Zhabotinsky reaction gained worldwide fame. Many works are devoted to the experimental and model study of the Belousov-Zhabotinsky reaction, since it makes it possible to observe the features of complex self-organization processes in a simple chemical system and allows for various types of control. The results of the study of the Belousov-Zhabotinsky reaction gave a powerful impetus to the development of such new sections modern science, as synergetics, thermodynamics of nonequilibrium processes. Synergetics studies the processes of self-organization in open nonequilibrium systems. Synergetics teaches you to see the world in a new way. Synergetics demonstrates how chaos can act as a mechanism of evolution, how order can arise from chaos. It allows you to look at the world as a single and very complex system , developing according to the laws of nonlinear dynamics. It is necessary to study a variety of processes - reversible and irreversible, equilibrium and nonequilibrium, deterministic and non-deterministic, which take place in phenomena of varying levels of complexity. 90 Copyright JSC Central Design Bureau BIBKOM & LLC Book-Service Agency 9.5. Chemical balance. Possibilities for controlling chemical processes 9.5.1. Concept of chemical equilibrium Chemical reactions are divided into two groups: 1. Reversible reactions are reactions that proceed in two mutually opposite directions: N2 + 3H2 2NH3 2. Irreversible reactions proceed in one direction: Zn + 2HCl → ZnCl2 + H2 Chemical equilibrium - This is a state of a system of reacting substances in which the rates of forward and reverse reactions are equal. For a reversible reaction at a given temperature, the ratio of the product of the equilibrium concentrations of the reaction products to the product of the equilibrium concentrations of the starting substances raised to the power of their stoichiometric coefficients is a constant value called the equilibrium constant (Kp). Application of the law of mass action for a reversible reaction allows us to derive the equilibrium constant Kp. The equilibrium constant shows how many times the rate of the forward reaction is greater than the rate of the reverse reaction at a given temperature and at constant concentrations of all participating substances equal to unity. The concentrations of substances established at chemical equilibrium are called equilibrium. They are indicated by the formulas of the reacting substances enclosed in square brackets. For the reaction aA + bB V1 = K1[A]a [B]b; cC + dD V2 = K1[C]c [D]d In equilibrium, V1 = V2. Then K1[A]a [B]b = K1[C]c [D]d Kr = K1 K2 Kp = [C]C [D]d . [A]a [B]b The equilibrium constant Kp depends on the nature of the reactants and on temperature, but does not depend on concentration. Kr is a dimensionless quantity. 9.5.2. Le Chatelier's Principle In 1901, French chemist Henri Louis Le Chatelier developed and patented a method for synthesizing ammonia from nitrogen and hydrogen. Studying the conditions under which this reaction occurs, he discovered and formulated the principle by which the direction of the shift in chemical equilibrium when changing the concentration of reactants, temperature and pressure (for gases) is determined. reactions). Table 9.2 – Le Chatelier’s principle Example: 2CO2 + 568 kJ/mol 2CO + O2 ∆Н = – 568 kJ/mol Conditions for shifting the equilibrium towards the product: 1. Increase in O2 concentration 2. Increase in pressure 3. Temperature reduction 9.5.3. Elements of thermochemistry Le Chatelier's principle is a consequence of the second law of thermodynamics. The shift in chemical equilibrium with temperature changes is associated with the energy of chemical reactions. These questions are studied in the course of physical chemistry in the section “Thermodynamics of chemical processes”. The study of the influence of temperature on the shift of chemical equilibrium should be carried out using the concepts of thermochemistry - a branch of chemical thermodynamics. Thermochemistry is a branch of chemical thermodynamics that studies the thermal effects of chemical processes. Exothermic reactions are reactions that release heat. Endothermic reactions are reactions that involve the absorption of heat. The thermochemical equation of a reaction is an equation that specifies the thermal effect of a reaction. The thermal effect of a reaction is the amount of heat released or absorbed per 1 mole of a substance. Two forms of writing thermochemical equations are used: 1. Heat characterizes changes in the environment during the reaction process (thermochemical effect of reaction Q). In this case, for an exothermic reaction, the thermal effect Q is indicated on the right side of the equation with a plus sign (+), since heat is released into the environment. 92 Copyright JSC Central Design Bureau BIBKOM & LLC Kniga-Service Agency For an endothermic reaction, the thermal effect Q is indicated with a minus sign (–), since heat is absorbed from the environment: 2H2 + O2 = 2H2O + 571 kJ/mol (exo ); N2 + O2 = 2NO – 180.5 kJ/mol (endo) 2. The thermal effect of a reaction reflects changes in the system of reactants itself (thermodynamic thermal effect of a reaction). In this case, the concept of enthalpy change ΔH is used. Enthalpy H is the energy reserve of a substance at constant pressure and temperature. For an exothermic reaction ΔH< 0, так как система теряет теплоту и запас еѐ энергии уменьшается (знак –). Для эндотермической реакции ΔН >0, since the system absorbs heat and its energy reserve increases (+ sign): 2H2 + O2 = 2H2O; ΔHo = – 571 kJ/mol (exo) N2 + O2 = 2NO; ΔHo = 180.5 kJ/mol (endo) The thermal effect of the reaction is usually given for standard conditions: normal pressure 101.3 kPa (1 atm) and temperature 298 K (ΔHo). The thermochemical effect of the reaction is equal in magnitude and opposite in sign to the thermodynamic thermal effect: Q = – ΔН. 9.6. Possibilities for controlling chemical processes The existence of reversible reactions is a manifestation of the law of unity and struggle of opposites operating in nature. Understanding this law in its specific manifestation is of great practical importance for optimal control of chemical processes in favorable conditions, to realize production efficiency and quality. Table 9.3 - Kinetic and thermodynamic factors influencing the chemical process Chemical reaction Nature of the reacting substances Thermodynamic factors: shift of equilibrium, possibility of reaction Kinetic factors influencing the rate of reaction (nature of the substance, concentration, temperature, pressure) Possibility of contradictory action 93 External environment of the reaction Copyright JSC "CDB "BIBKOM" & LLC "Agency Kniga-Service" Control of chemical reactions is possible only on the basis of understanding the laws to which they obey. It should be based on knowledge of chemical kinetics and thermodynamics - both the classical thermodynamics of equilibrium processes of isolated systems, and the modern thermodynamics of open nonequilibrium systems in which self-organization phenomena are possible. Questions for self-control 1. Give examples of reactions occurring at high and low speed. 2. Define activation energy. 3. How will the reaction rate change when heated from 20° to 50° if the Van't Hoff coefficient is 2? 4. Calculate the van’t Hoff coefficient if, when the temperature changes from 20° to 50°, the reaction rate increases 27 times. 5. The concentration of which substances is not included in the expression for the rate and equilibrium constant of the reaction: FeO3 + 3CO 2Fe + 3CO2 6. What factors will help shift the equilibrium of the reaction towards the formation of the reaction product: 2H2 + O2 = 2H2O + 571 kJ/mol. ∆Н = – 571 kJ/mol REFERENCES Main 1. Knyazev, D.A. Inorganic chemistry/D.A. Knyazev, S.N. Smarygin – M.: Bustard, 2004. – 592 p. 2. Glinka, N.L. General chemistry / N.L. Glinka – M.: KNORUS, 2009. – 752 p. 3. Ryazanova, G.E. General and inorganic chemistry. Tables and diagrams / G.E. Ryazanov. – Saratov: Federal State Educational Institution of Higher Professional Education “Saratov State Agrarian University”, 2006 – 284 p. Additional 1. Barkovsky, E.V. General chemistry: Course of lectures /E.V. Barkovsky, S.V. Tkachev. – Minsk: BSMU, 2009. – 132 p. 2. Trubetskov, D.I. Introduction to the theory of self-organization open systems . /DI. Trubetskov, E.S. Mchedlova, L.V. Krasichkov. – M.: Publishing house of physical and mathematical literature, 2002. – 200 p. 3. Ruzavin, G.I. Problems of simple and complex in the evolution of sciences. //G.I. Ruzavin. – Questions of philosophy. – 2008. No. 3. – p. 102. 94 Copyright JSC Central Design Bureau BIBKOM & LLC Book-Service Agency BIBLIOGRAPHICAL LIST 1. Agafoshin, N.P. Periodic law and periodic system D.I. Mendeleev. / N.P. Agafoshin - M.: Education, 1973. - 208 p. 2. Barkovsky, E.V. General chemistry: Course of lectures /E.V. Barkovsky, S.V. Tkachev. – Minsk: BSMU, 2009. – 132 p. 3. Vlasov, V.M. Errors leading to an explosion / V.M. Vlasov // Chemistry and life. – 2006. – No. 7. p. 60. 4. Gelfman, M.I. Inorganic chemistry / M.I. Gelfman, V.P. Yustratov. – St. Petersburg: Publishing house “Lan”, 2009. – 528 p. 5. Glinka, N.L. General chemistry / N.L. Glinka – M.: KNORUS, 2009. – 752 p. 6. Guzey, L.S. General chemistry / L.S. Guzey, V.N. Kuznetsov, A.S. Guzey. – M.: Moscow State University Publishing House, 1999. – 333 p. 7. Dmitriev, S.N. Chemical identification and study of the properties of superheavy metals. Evolution of the periodic system D.I. Mendeleev / S.N. Dmitriev – Abstracts of reports of the XVIII Mendeleev Congress on General and Applied Chemistry: In 5 volumes; v.1. – M.: Granitsa, 2007. – P. 47. 8. Egorov, V.V. Environmental chemistry. /V.V. Egorov. – St. Petersburg: Lan Publishing House, 2009. – 192 p. 9. Klinsky, G.D. Inorganic chemistry /G.D. Klinsky, V.D. Skopintsev. – M: Publishing house MCHA, 2001. – 384 p. 10. Knyazev, D.A. Inorganic chemistry/D.A. Knyazev, S.N. Smarygin – M.: Bustard, 2004. – 592 p. 11. Lensky, A.S. Introduction to bioinorganic and biophysical chemistry / A.S. Lensky. – M.: Higher School, 1989. – 256 p. 12. Mineev, V.G. In defense of nitrates and phosphates / V.G. Mineev // Chemistry and life. – 2008. No. 5. – P. 20. 13. Naydysh, V.M. Concepts of modern natural science. /V.M. Naydysh. – M.: Alfa-M; INFRA-M, 2004, – 622 p. 14. Inorganic chemistry (biogenic and abiogenic elements): Textbook / ed. V.V. Egorova. – St. Petersburg: Lan Publishing House, 2009. – 320 p. 15. Ruzavin, G.I. Problems of simple and complex in the evolution of sciences. //G.I. Ruzavin. – Questions of philosophy. – 2008. No. 3. – p. 102. 16. Ryazanova, G.E. Inorganic and analytical chemistry / G.E. Ryazanova - Saratov: Federal State Educational Institution of Higher Professional Education "Saratov State Agrarian University", 2006. - 172 p. 17. Ryazanova, G.E. General and inorganic chemistry. Tables and diagrams / G.E. Ryazanov. – Saratov: Federal State Educational Institution of Higher Professional Education “Saratov State Agrarian University”, 2006 – 284 p. 18. Sirotkin, O.S. Chemistry in its place / O.S. Sirotkin // Chemistry and life. – 2003. - No. 5. – P. 26. 19. Trubetskov, D.I. Introduction to the theory of self-organization of open systems. / D.I. Trubetskov, E.S. Mchedlova, L.V. Krasichkov. – M.: Publishing house of physical and mathematical literature, 2002. – 200 p. 20. Ugai, Ya.A. General and inorganic chemistry /Ya.A. Ugai. – M.: Higher School, 2004. – 528 p. 95 Copyright OJSC "CDB "BIBKOM" & LLC "Agency Kniga-Service" CONTENTS Introduction................................... ........................................................ ........................................... 3 Lecture 1. Basic concepts and basic laws of chemistry................................... 4 1.1. Goals of studying the discipline............................................................. ............................... 4 1.2. The importance of chemistry for agriculture......................................................... ........ 4 1.3. Chemistry subject........................................................ ................................................... 5 1.4 . Dialectics of basic concepts and laws of chemistry.................................................. 5 1.5. Methods for studying chemistry........................................................ .................................. 8 Questions for self-control.................. ........................................................ .................. 11 References.................................... ........................................................ ............... 12 Lecture 2. Modern teaching about the structure of the atom.................................. ................... 13 2.1. Dialectics of ideas about the structure of the atom.................................................... 13 2.2. Modern quantum mechanical model of the structure of the atom.................................. 14 2.3. Electronic formulas of elements................................................... ............... 17 Questions for self-control .................................... ........................................................ 20 References.................................................................... ................................................... 20 Lecture 3. Periodic law and periodic system D.I. Mendeleev........................................................ ................................................... 21 3.1 . Periodic law and periodic system of elements D.I. Mendeleev – the basis of modern chemistry.................................................... ... 21 3.2. Structure of the periodic table......................................................... .............. 23 3.3. Algorithm for general characteristics of an element atom.................................................. 24 3.4. Frequency of changes in the properties of elements and their compounds.................. 27 Questions for self-control.................................... ........................................................ ..... 30 References.................................................... ........................................................ .. 30 Lecture 4. Manifestation of the periodic law in the acid-base properties of inorganic compounds................................................. ........................ 32 4.1. Genetic relationship of the main classes of inorganic compounds....... 32 4.2. Chemical properties of oxides, bases, acids and salts.................................... 33 4.3. Frequency of changes in the acid-base properties of chemicals.................................................... ........................................................ 39 4.4. Questions of professional competence................................................................... .. 40 Questions for self-control.................................................... .................................... 40 References......... ........................................................ .................................. 41 Lecture 5. Chemical bond.......... ........................................................ ...................... 43 5.1. Modern ideas about chemical bonding.................................................... 43 5.2. Fundamentals of the valence bond method (MVM) W. Heitler and F. London (1927) .................................. ........................................... 43 5.2.1. Mechanisms of covalent bond formation.................................................... 43 5.2.2. Covalent bond................................................... ................................ 44 5.2.3. Types of chemical bonds................................................................... ........ 45 5.2.4. Types of chemical bonds................................................................... ........................... 46 5.2.5. Degree of bond ionicity................................................................... ....................... 46 5.2.6. Hybridization of atomic orbitals.................................................... ...... 48 5.2.7. Metal connection........................................................ ............................... 50 5.2.8. Hydrogen bond................................................... ................................... 50 5.3. Concept of the molecular orbital method................................................................. 50 96 Copyright OJSC Central Design Bureau BIBKOM & LLC Book-Service Agency Questions for self-control............................. ........................................................ .. 52 References................................................... ........................................................ ....... 52 Lecture 6. Modern theory of solutions.................................... ............................... 53 6.1. Classification of disperse systems................................................................... ................... 53 6.2. Methods of expressing the composition of solutions.................................................. .............. 54 6.3. Soil solution concentration and osmosis.................................................... ...... 54 6.4. Electrolyte solutions................................................... ..................................... 55 6.4.1. Aqueous solutions........................................................ ........................................... 55 6.4.2. Theory of electrolytic dissociation................................................................. .. 56 6.4.3. Quantitative characteristics of electrolytic dissociation................................................................. ........................................................ .......... 57 6.4.4. Properties of strong electrolytes................................................................... ............... 58 6.4.5. Types of electrolytes................................................... .................................... 59 6.4.6. Dissociation of electrolytes................................................... ....................... 60 6.4.7. Reactions in electrolyte solutions.................................................. ............. 61 6.4.8. Hydrolysis of salts................................................... ........................................... 61 6.4.9. Ionic product of water. Hydrogen pH................................... 62 6.4.10. The meaning of solutions................................................... .................................. 64 Questions for self-control............... ........................................................ ........................... 64 References.................................... ........................................................ ........................ 65 Lecture 7. Redox reactions.................. ........................... 66 7.1. Modern theory of redox reactions (ORR) ...... 66 7.1.1. The importance of redox processes........................ 66 7.1.2. Basic provisions of the theory of ODD.................................................. ............. 67 7.1.3. Calculation of the oxidation state of an element.................................................... 68 7.1.4. Algorithm for characterizing the redox properties of complex compounds. ........................................................ ........................... 70 7.1.5. Types of redox reactions.................................................... 70 7.2. Algorithms for composing equations of redox reactions.................................................... ........................................................ ........................... 71 7.2.1. Algorithm for compiling the OVR equation using the electronic balance method.................................................... ........................................................ 71 7.2.2. Drawing up the ORR equation using the electron-ion method (half-reaction method) .................................................... ........................................................ .... 71 7.3. Redox properties of compounds of biogenic elements.................................................... ........................................................ ....................... 73 7.4. Redox potentials. Direction of ODD......................... 74 Questions for self-control.................................... ........................................................ ...... 75 References................................................... ........................................................ ......... 75 Lecture 8. Complex compounds.................................... ....................................... 76 8.1. A Brief History of Complex Compounds (CCs) .................................................... 76 8.2. Werner's coordination theory and modern ideas................................ 77 8.2.1. Composition of molecules of complex compounds.................................................... 78 8.2.2. Nomenclature of complex compounds................................................................... 79 8.3. Chemical bonding in complex compounds.................................................... .. 80 8.4. Electrolytic dissociation of complex compounds.................................... 81 Questions for self-control...... ........................................................ ....................... 84 References.................................... ........................................................ ........................ 84 97 Copyright JSC Central Design Bureau BIBKOM & LLC Kniga-Service Agency Lecture 9. Chemical kinetics. Chemical equilibrium................................................... 85 9.1. The concept of chemical kinetics.................................................... ........................... 85 9.2. Rate of chemical reaction................................................................... ............................ 85 9.3. Factors influencing the rate of a chemical reaction.................................................... 86 9.3.1. The influence of concentration on the rate of reaction.................................................... 87 9.3. 2. The influence of temperature on the reaction rate................................................... 88 9.3.3. The influence of catalysts on the rate of chemical reactions.................................... 89 9.4. The concept of oscillatory reactions................................................................. .................... 90 9.5. Chemical balance. Possibilities for controlling chemical processes................................................................. ........................................................ ........................... 91 9.5.1. The concept of chemical equilibrium.................................................... ............ 91 9.5.2. Le Chatelier's principle......................................................... ................................... 91 9.5.3. Elements of thermochemistry................................................... ........................... 92 9.6. Possibilities for controlling chemical processes.................................................... 93 Questions for self-control.... ........................................................ .................................. 94 References.................. ........................................................ .................................... 94 Bibliography........... ........................................................ ......................... 95 Contents................... ........................................................ ................................................ 96 98
Chapter 1.
General chemical and environmental patterns.
Where does chemistry begin?
Is this a difficult question? Everyone will answer it differently.
In secondary school, students study chemistry over a number of years. Many people do quite well on their final exam in chemistry. However…
Conversations with applicants and then first-year students indicate that residual knowledge in chemistry after secondary school is insignificant. Some get confused in various definitions and chemical formulas, while others cannot even reproduce the basic concepts and laws of chemistry, not to mention the concepts and laws of ecology.
Their chemistry never started.
Chemistry, apparently, begins with a deep mastery of its fundamentals, and above all, the basic concepts and laws.
1.1. Basic chemical concepts.
In D.I. Mendeleev’s table there are numbers next to the element symbol. One number indicates the atomic number of the element, and the second atomic mass. The serial number has its own physical meaning. We will talk about it later, here we will focus on atomic mass and highlight in what units it is measured.
It should be noted right away that the atomic mass of an element given in the table is a relative value. The unit of relative atomic mass is taken to be 1/12 of the mass of a carbon atom, an isotope with a mass number of 12, and is called the atomic mass unit /amu/. Therefore, 1 amu equal to 1/12 of the mass of the carbon isotope 12 C. And it is equal to 1.667 * 10 –27 kg. /The absolute mass of a carbon atom is 1.99 * 10 –26 kg./
Atomic mass, given in the table, is the mass of the atom expressed in atomic mass units. The quantity is dimensionless. Specifically for each element, atomic mass shows how many times the mass of a given atom is greater or less than 1/12 of the mass of a carbon atom.
The same can be said about molecular weight.
Molecular mass is the mass of a molecule expressed in atomic mass units. The magnitude is also relative. The molecular mass of a particular substance is equal to the sum of the masses of the atoms of all the elements that make up the molecule.
An important concept in chemistry is the concept of “mole”. Mole– such an amount of substance that contains 6.02 * 10 23 structural units /atoms, molecules, ions, electrons, etc./. Mole of atoms, mole of molecules, mole of ions, etc.
The mass of one mole of a given substance is called its molar / or molar / mass. It is measured in g/mol or kg/mol and is designated by the letter “M”. For example, the molar mass of sulfuric acid M H 2 SO4 = 98 g/mol.
The next concept is “Equivalent”. Equivalent/E/ is the weight amount of a substance that interacts with one mole of hydrogen atoms or replaces such an amount in chemical reactions. Therefore, the equivalent of hydrogen E H is equal to one. /E N =1/. The oxygen equivalent E O is equal to eight /E O =8/.
A distinction is made between the chemical equivalent of an element and the chemical equivalent of a complex substance.
The equivalent of an element is a variable quantity. It depends on the atomic mass /A/ and valence /B/ that the element has in a particular compound. E=A/B. For example, let's determine the equivalent of sulfur in the oxides SO 2 and SO 3. In SO 2 E S =32/4=8, and in SO 3 E S =32/6=5.33.
The molar mass of an equivalent, expressed in grams, is called equivalent mass. Therefore, the equivalent mass of hydrogen ME H = 1 g/mol, the equivalent mass of oxygen ME O = 8 g/mol.
The chemical equivalent of a complex substance /acid, hydroxide, salt, oxide/ is the amount of the corresponding substance that interacts with one mole of hydrogen atoms, i.e. with one equivalent of hydrogen or replaces that amount of hydrogen or any other substance in chemical reactions.
Acid equivalent/E K/ is equal to the quotient of the molecular weight of the acid divided by the number of hydrogen atoms participating in the reaction. For the acid H 2 SO 4, when both hydrogen atoms react H 2 SO 4 +2NaOH=Na 2 SO+2H 2 O the equivalent will be equal to EN 2 SO4 = M H 2 SO 4 /n H =98/2=49
Hydroxide equivalent /E hydr. / is defined as the quotient of the molecular weight of the hydroxide divided by the number of hydroxo groups that react. For example, the equivalent of NaOH will be equal to: E NaOH = M NaOH / n OH = 40/1 = 40.
Salt equivalent/E salt/ can be calculated by dividing its molecular weight by the product of the number of metal atoms that react and their valency. Thus, the equivalent of the salt Al 2 (SO 4) 3 will be equal to E Al 2 (SO 4) 3 = M Al 2 (SO 4) 3 /6 = 342/2.3 = 342/6 = 57.
Oxide equivalent/E ok / can be defined as the sum of the equivalents of the corresponding element and oxygen. For example, the equivalent of CO 2 would be equal to the sum equivalents of carbon and oxygen: E CO 2 =E C +E O =3+8=7.
For gaseous substances it is convenient to use equivalent volumes /E V /. Since when normal conditions A mole of gas occupies a volume of 22.4 liters, then based on this value, it is easy to determine the equivalent volume of any gas. Let's consider hydrogen. The molar mass of hydrogen 2g occupies a volume of 22.4 liters, then its equivalent mass of 1g occupies a volume of 11.2 liters / or 11200 ml /. Therefore E V N =11.2l. The equivalent volume of chlorine is 11.2 l /E VCl = 11.2 l/. The equivalent volume of CO is 3.56 /E VC O =3.56 l/.
The chemical equivalent of an element or complex substance is used in stoichiometric calculations of exchange reactions, and in the corresponding calculations of redox reactions, oxidative and reduction equivalents are used.
Oxidative equivalent is defined as the quotient of the molecular weight of the oxidizing agent divided by the number of electrons it accepts in a given redox reaction.
The reducing equivalent is equal to the molecular weight of the reducing agent divided by the number of electrons it gives up in a given reaction.
Let's write the redox reaction and determine the equivalent of the oxidizing agent and reducing agent:
5N 2 aS+2KMnO 4 +8H 2 SO 4 =S+2MnSO 4 +K 2 SO 4 +5Na 2 SO 4 +8H 2 O
The oxidizing agent in this reaction is potassium permanganate. The equivalent of the oxidizing agent will be equal to the mass of KMnO 4 divided by the number of electrons accepted by the oxidizing agent in the reaction (ne=5). E KMnO 4 =M KMnO 4 /ne=158/5=31.5. The molar mass of the equivalent of the oxidizing agent KMnO 4 in an acidic medium is 31.5 g/mol.
The equivalent of the reducing agent Na 2 S will be: E Na 4 S = M Na 4 S / ne = 78/2 = 39. The molar mass of Na 2 S equivalent is 39 g/mol.
In electrochemical processes, in particular during the electrolysis of substances, an electrochemical equivalent is used. The electrochemical equivalent is determined as the quotient of the chemical equivalent of the substance released at the electrode divided by the Faraday number /F/. The electrochemical equivalent will be discussed in more detail in the corresponding paragraph of the course.
Valence. When atoms interact, a chemical bond is formed between them. Each atom can only form a certain number of bonds. The number of connections determines this unique property each element, which is called valency. In the most general view Valency is the ability of an atom to form a chemical bond. One chemical bond that a hydrogen atom can form is taken as a unit of valency. In this regard, hydrogen is a monovalent element, and oxygen is a divalent element, because No more than two hydrogens can form a bond with an oxygen atom.
The ability to determine the valency of each element, including in a chemical compound, is a necessary condition successful completion of the chemistry course.
Valence is also related to such a concept of chemistry as oxidation state. The oxidation substate is the charge that an element has in an ionic compound or would have in a covalent compound if the shared electron pair were completely shifted to a more electronegative element. The oxidation state has not only a numerical expression, but also a corresponding charge sign (+) or (–). Valence does not have these signs. For example, in H 2 SO 4 the oxidation state is: hydrogen +1, oxygen –2, sulfur +6, and the valency, accordingly, will be 1, 2, 6.
Valency and oxidation state in numerical values do not always coincide in value. For example, in a molecule of ethyl alcohol CH 3 –CH 2 –OH the valence of carbon is 6, hydrogen is 1, oxygen is 2, and the oxidation state, for example, of the first carbon is –3, the second is –1: –3 CH 3 – –1 CH 2 –OH.
1.2. Basic environmental concepts.
Recently, the concept of “ecology” has deeply entered our consciousness. This concept, introduced back in 1869 by E. Haeckel, comes from the Greek oikos- house, place, dwelling, logos– the teaching / is disturbing humanity more and more.
In biology textbooks ecology defined as the science of the relationship between living organisms and their environment. An almost consonant definition of ecology is given by B. Nebel in his book “Science of the Environment” - Ecology is the science of various aspects of the interaction of organisms with each other and with the environment. A broader interpretation can be found in other sources. For example, Ecology – 1/. The science that studies the relationship between organisms and their systemic assemblies and the environment; 2/. A set of scientific disciplines that study the relationship of systemic biological structures /from macromolecules to the biosphere/ among themselves and with the environment; 3/. A discipline that studies the general laws of functioning of ecosystems at various hierarchical levels; 4/. A comprehensive science that studies the habitat of living organisms; 5/. Study of the position of man as a species in the biosphere of the planet, his connections with ecological systems and the impact on them; 6/. The science of environmental survival. / N.A. Agidzhanyan, V.I. Torshik. Human ecology./. However, the term “ecology” refers not only to ecology as a science, but to the state of the environment itself and its impact on humans, flora and fauna.
The manual is intended for schoolchildren, applicants and teachers. The manual outlines the modern fundamentals of chemistry in a brief but informative and clear manner. These are the basics that every high school graduate must understand and absolutely must know for anyone who sees himself as a chemistry, medical, or biologist student of the 21st century.
Atomic-molecular theory.
The atomic-molecular theory of the structure of matter arose as a result of scientists’ attempts to solve two main issues. 1) What do substances consist of? 2) Why are substances different and why can some substances transform into others? The main provisions of this theory can be formulated as follows:
1. All substances are made up of molecules. A molecule is the smallest particle of a substance that has its chemical properties.
2. Molecules are made up of atoms. An atom is the smallest particle of an element in chemical compounds. Different elements correspond to different atoms.
3. During chemical reactions, molecules of some substances are transformed into molecules of other substances. Atoms do not change during chemical reactions.
Let us briefly consider the history of the creation and development of atomic-molecular theory.
Atoms were invented in Greece in the 5th century. BC e. The philosopher Leucippus wondered whether every piece of matter, no matter how small, could be divided into even smaller pieces. Leucippus believed that as a result of such division one could obtain such a small particle that further division would become impossible. Leucippus' student and philosopher Democritus called these tiny particles "atoms." He believed that the atoms of each element have a special size and shape and that this explains the differences in the properties of the elements. The substances that we see and feel are compounds of atoms various elements, and by changing the nature of this compound, one substance can be converted into another. Democritus created the atomic theory almost modern form. However, this theory is only the fruit of philosophical reflection, not confirmed by experimental observations.
TABLE OF CONTENTS
Preface 3
PART 1. THEORETICAL CHEMISTRY 5
CHAPTER 1. Basic concepts and laws of chemistry 5
§ 1.1. Chemistry subject 5
§1.2. Atomic-molecular theory 7
§ 1.3. Law of conservation of mass and energy 10
§ 1.4. Periodic Law 12
§ 1.5. Basic Chemistry Concepts 14
§ 1.6. Stoichiometric ratios in chemistry 18
§ 1.7. Gas laws 19
CHAPTER 2. Atomic structure 22
§ 2.1. Development of ideas about the complex structure of the atom 22
§ 2.2. Quantum numbers of electrons 25
§ 2.3. Distribution of electrons in atoms 28
§ 2.4. Radioactive transformations 33
§ 2.5. Periodicity of properties of atoms of elements 37
CHAPTER 3. Chemical bonding and molecular structure 41
§ 3.1. The nature of the chemical bond 41
§ 3.2. Covalent bond 44
§ 3.3. Ionic bond 48
§ 3.4. Metal connection 50
§ 3.5. Intermolecular chemical bonds 51
§ 3.6. Valency and oxidation state 55
§ 3.7. Spatial structure of molecules 58
CHAPTER 4. States of matter 63
§ 4.1. Characteristic properties gases, liquids and solids 63
§ 4.2. Phase diagrams of substances 66
§ 4.3. Gases 68
§ 4.4. Liquids 70
§ 4.5. Crystalline substances 73
§ 4.6. Various forms of existence of substances 80
CHAPTER 5. Energy effects of chemical reactions 81
§ 5.1. Release and absorption of energy in chemical reactions 81
§ 5.2. Exothermic and endothermic reactions. Thermochemical law of Hess 87
CHAPTER 6. Kinetics of chemical reactions 93
§ 6.1. Basic concepts and postulates of chemical kinetics 93
§ 6.2. Effect of temperature on reaction rate 97
§ 6.3. Catalysis 99
CHAPTER 7. Chemical equilibrium 103
§ 7.1. Determination of equilibrium state 103
§ 7.2. Chemical equilibrium constant 105
§ 7.3. Shift in chemical equilibrium. Le Chatelier's Principle 108
§ 7.4. On optimal conditions for obtaining substances on an industrial scale 111
CHAPTER 8. Solutions 114
§ 8.1. Dissolution as a physicochemical process 114
§ 8.2. Factors affecting the solubility of substances 117
§ 8.3. Ways to express the concentration of solutions 121
CHAPTER 9. Electrolytic dissociation and ionic reactions in solutions 122
§ 9.1. Electrolytes and electrolytic dissociation 122
§ 9.2. Degree of dissociation. Strong and weak electrolytes. Dissociation constant 123
§ 9.3. Ionic reaction equations 126
§ 9.4. Hydrolysis of salts 128
CHAPTER 10. Basic types of chemical reactions 129
§ 10.1. Symbolism and classification characteristics of reactions 129
§ 10.2. Classification by the number and composition of reagents and reaction products 131
§ 10.3. Classification of reactions according to phase characteristics 136
§ 10.4. Classification of reactions according to the type of particles transferred 137
§ 10.5. Reversible and irreversible chemical reactions 138
CHAPTER 11. Redox processes 140
§ 11.1. Redox reactions 140
§ 11.2. Selection of stoichiometric coefficients in OVR 144
§ 11.3. Standard potentials OVR 148
§ 11.4. Electrolysis of solutions and melts of electrolytes 152
PART II. INORGANIC CHEMISTRY 154
CHAPTER 12. general characteristics inorganic compounds, their classification and nomenclature 154
§ 12.1. Oxides 155
§ 12.2. Bases (metal hydroxides) 158
§ 12.3. Acids 160
§ 12.4. Salts 165
CHAPTER 13. Hydrogen 168
§ 13.1. Atomic structure and position in the periodic table D.I. Mendeleeva 168
§ 13.2. Chemical properties of hydrogen 171
§ 13.3. Production of hydrogen and its use 173
§ 13.4. Hydrogen oxides 174
CHAPTER 14. Halogens 178
§ 14.1. Physical properties halogens 178
§ 14.2. Chemical properties and production of halogens 180
§ 14.3. Hydrogen halides, hydrohalic acids and their salts 185
§ 14.4. Oxygen-containing halogen compounds 187
CHAPTER 15. Chalcogens 190
§ 15.1. General characteristics 190
§ 15.2. Simple substances 191
§ 15.3. Sulfur compounds 196
CHAPTER 16. Nitrogen subgroup 204
§ 16.1. General characteristics 204
§ 16.2. Properties of simple substances 205
§ 16.3. Ammonia. Phosphine. Phosphorus halides 207
§ 16.4. Nitrogen oxides. Nitric and nitrous acids 210
§ 16.5. Phosphorus oxides and acids 214
CHAPTER 17. Carbon subgroup 218
§ 17.1. General characteristics 218
§ 17.2. Carbon 219
§ 17.3. Carbon oxides 223
§ 17.4. Carbonic acid and its salts 226
§ 17.5. Silicon 228
§ 17.6. Silicon compounds with oxidation state +4 230
§ 17.7. Silicon compounds with oxidation state -4 233
CHAPTER 18. Properties of s-metals and their compounds 234
§ 18.1. General characteristics 234
§ 18.2. Chemical properties of metals 236
§ 18.3. Compounds of s-metals 239
CHAPTER 19. Aluminum and boron 240
§ 19.1. General characteristics 240
§ 19.2. Properties and preparation of simple substances 242
§ 19.3. Boron and aluminum compounds 247
CHAPTER 20. Main transition metals 249
§ 20.1. General characteristics 249
§ 20.2. Chromium and its compounds 251
§ 20.3. Manganese and its compounds 253
§ 20.4. Iron triad 255
§ 20.5. Iron and steel production 258
§ 20.6. Copper and its compounds 261
§ 20.7. Zinc and its compounds 263
§ 20.8. Silver and its compounds 264
CHAPTER 21. Noble gases 265
§ 21.1. General characteristics 265
§ 21.2. Chemical compounds of noble gases 267
§ 21.3. Application of noble gases 269
PART III. ORGANIC CHEMISTRY 271
CHAPTER 22. Basic concepts and patterns in organic chemistry 271
§ 22.1. Organic Chemistry Subject 271
§ 22.2. Classification of organic compounds 272
§ 22.3. Nomenclature of organic compounds 274
§ 22.4. Isomerism of organic compounds 278
§ 22.5. Electronic effects and reactivity of organic compounds 279
§ 22.6. General characteristics 281
CHAPTER 23. Saturated hydrocarbons 283
§ 23.1. Alkanes 283
§ 23.2. Cycloalkanes 286
CHAPTER 24. Alkenes and alkadienes 289
§ 24.1. Alkenes 289
§ 24.2. Diene hydrocarbons 293
CHAPTER 25. Alkynes 295
§ 25.1. General characteristics 295
§ 25.2. Preparation and chemical properties 296
CHAPTER 26. Arenas 300
§ 26.1. General characteristics 300
§ 26.2. Preparation and chemical properties 303
§ 26.3. Orientants (deputies) of the first and second kind 308
CHAPTER 27. Alcohol and phenols 310
§ 27.1. General characteristics 310
§ 27.2. Monohydric alcohols 311
§ 27.3. Polyhydric alcohols 315
§ 27.4. Phenols 316
CHAPTER 28. Aldehydes and ketones 321
§ 28.1. General characteristics 321
§ 28.2. Ways to get 323
§ 28.3. Chemical properties 324
CHAPTER 29. Carboxylic acids 327
§ 29.1. Classification, nomenclature and isomerism 327
§ 29.2. Monobasic saturated carboxylic acids 334
§ 29.3. Monobasic unsaturated carboxylic acids 339
§ 29.4. Aromatic carboxylic acids 342
§ 29.5. Dibasic carboxylic acids 343
CHAPTER 30. Functional derivatives of carboxylic acids 345
§ 30.1. Classification of functional derivatives 345
§ 30.2. Carboxylic acid anhydrides 346
§ 30.3. Carboxylic acid halides 348
§ 30.4. Amides of carboxylic acids 350
§ 30.5. Esters 352
§ 30.6. Fats 353
CHAPTER 31. Carbohydrates (sugars) 357
§ 31.1. Monosaccharides 357
§ 31.2. Individual representatives of monosaccharides 363
§ 31.3. Oligosaccharides 366
§ 31.4. Polysaccharides 368
CHAPTER 32. Amines 371
§ 32.1. Saturated aliphatic amines 371
§ 32.2. Aniline 375
CHAPTER 33. Amino acids. Peptides. Proteins 377
§ 33.1. Amino acids 377
§ 33.2. Peptides 381
§ 33.3. Proteins 383
CHAPTER 34. Nitrogen-containing heterocyclic compounds 387
§ 34.1. Six-membered heterocycles 387
§ 34.2. Compounds with a five-membered ring 390
CHAPTER 35. Nucleic acids 393
§ 35.1. Nucleotides and nucleosides 393
§ 35.2. Structure of nucleic acids 395
§ 35.3. Biological role of nucleic acids 398
CHAPTER 36. Synthetic high-molecular compounds (polymers) 400
§ 36.1. General characteristics 400
§ 36.2. Plastics 402
§ 36.3. Fiber 404
§ 36.4. Rubbers 405
Recommended reading 410.
