A short course in chemistry. Chemistry, short course of lectures

"Saratov State Agrarian University

Chemistry

short course of lectures

for first year students

Direction of training

250100.62 Forestry

Training profile

Forestry

Saratov 2011

Reviewers:

Head of the Department of General and Inorganic Chemistry, Doctor of Chemical Sciences, Professor of the State Educational Institution of Higher Professional Education “SSU named after. Chernyshevsky."

Head of the Department of Chemistry and Fundamentals of Ecology, Candidate of Chemical Sciences, Professor of the Federal State Educational Institution of Higher Professional Education "Saratov State Agrarian University"

Chemistry: a short course of lectures for first-year students of the specialty (direction of training) 250100.62 “Forestry” / Compiled by: , // Federal State Educational Institution of Higher Professional Education “Saratov State Agrarian University”. – Saratov, 2011. – 80 p.

A short course of lectures on the discipline “Chemistry” is compiled in accordance with the work program of the discipline and is intended for students in the training direction 250100.62 “Forestry”. A short course of lectures contains theoretical material on the main issues of general, inorganic and organic chemistry, and issues of identification of chemical substances are considered. 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.

Introduction.

Chemistry is one of the most important natural science disciplines. Chemistry studies substances, their structure, properties and transformations that occur as a result of chemical reactions, as well as the fundamental laws to which these transformations are subject. Modern chemistry is such a vast area of natural science that many of its sections are independent, although interrelated, scientific disciplines.

A short course of lectures on the discipline “Chemistry” is intended for students in the field of study 250100.62 “Forestry”. It reveals the basic laws of general chemistry on which chemical disciplines are based, includes an introduction to inorganic chemistry, introduces the main classes of organic compounds, and involves mastering the theoretical foundations of analytical methods. The course is aimed at developing the key competencies necessary for effectively solving professional problems and organizing professional activities based on a deep understanding of the laws of ecosystem functioning.

Lecture 1

CHEMICAL KINETICS. CHEMICAL EQUILIBRIUM

1.1. WITHchemical reaction rate

Chemical kinetics called the branch of chemistry that studies the rate and mechanisms of chemical reactions .

Under speed of chemical reaction understand the change in the concentration of a substance per unit time. In this case, it does not matter what substance we are talking about - the reactant or the reaction product.

If during the period of time from t1 to t2 the concentration of the substance changed from the value C1 to C2, then the expression for the reaction rate is:

V = ± = ± , mol/l∙s

In this case, the “+” sign in the formula refers to a change in the concentration of the substance formed as a result of the reaction (C2>C1, ΔC>0), and the “–” sign refers to a change in the concentration of the substance entering the reaction (C1>C2, ΔC<0).

1.2. Factors affecting the rate of a chemical reaction

The rate of a chemical reaction depends on the nature of the reactants, their concentration, temperature, pressure - for gases, contact surface area (degree of grinding) - for solids, and the presence of a catalyst.

Influence of the nature of the reacting substances. Different substances have different reactivity. For example, potassium (an alkali metal) reacts violently with water, releasing hydrogen, while gold practically does not react with water.

The reactivity of substances is determined to a large extent by the nature of chemical bonds and the structure of the reagent molecules.

1.3. The influence of the concentration of reactants on the reaction rate.

Law of Mass Action

A necessary condition for chemical interaction is the collision of particles with each other. The more collisions, the faster the reaction occurs. With increasing concentration (the number of particles per unit volume), collisions occur more frequently and, therefore, the reaction rate increases.

The dependence of the reaction rate on the concentration of reactants is characterized law of mass action (K. Guldberg, P. Waage, 1867):

The rate of a chemical reaction is directly proportional to the product of the concentrations of the reactants raised to the power of their stoichiometric coefficients.

For a reaction proceeding according to the equation A A+ V B → With The reaction rate is determined by the expression:

V = k [A] A∙[B] b,

where k is the reaction rate constant, depends on the temperature and nature of the reactants, but does not depend on their concentration.

In relation to specific reactions, the expression of the law of mass action will have the form (Table 1):

Table 1 - Examples of expression of the law of action for various reactions

Reaction equation

Expression of the law of mass action

N2 + 3H2 → 2NH3

V = k ∙ 3

(reaction rate does not depend on solid concentration)

1.4. The effect of temperature on the reaction rate. Van't Hoff's rule.

The temperature dependence of the reaction rate is approximately estimated van't Hoff's rule:

With every 10 degree increase in temperature, the reaction rate increases by 2-4 times.

Where γ – temperature coefficient of reaction rate, equal to 2–4.

The increase in reaction rate with increasing temperature explains activation theory (S. Arrhenius). According to this theory, not all molecules react during a collision, but only active ones - those that have sufficient energy, excess compared to the average energy of molecules at a given temperature - activation energy . So, activation energy Ea (dimension - kJ/mol) is the excess energy that molecules must have in order for their collision to lead to a chemical transformation. In other words, each reaction is characterized by a certain energy barrier. As the temperature rises, the number of active molecules increases, which leads to an increase in the reaction rate.

Arrhenius equation:

where A is the pre-exponential factor and is associated with the frequency of particle collisions and their orientation during collisions.

As follows from the Arrhenius equation, the lower the activation energy and the higher the temperature, the higher the reaction rate.

Figure 1. Energy diagram of a chemical reaction:

A – reagents, B – activated complex

(transition state), C – products.

The reaction proceeds through the stage of formation of an unstable intermediate compound - activated complex. It is for its formation that activation energy is required. This complex is unstable, it exists for a very short time, and as a result of its disintegration reaction products are formed. In the simplest case, an activated complex can be thought of as a configuration of atoms in which old chemical bonds are weakened and new ones are formed.

1.5. Catalysis. Catalysts

Catalysis – the phenomenon of changes in reaction rates under the influence of substances – catalysts. Distinguish positive catalysis (increased reaction rate) and o negative catalysis (reaction slowdown under the influence of substances - inhibitors ). The catalyst itself is not consumed during the reaction, but changes its speed.

There are homogeneous and heterogeneous catalysis. When homogeneous catalysis the catalyst and reactants are in the same state of aggregation. When heterogeneous catalysis – in different states of aggregation.

Examples of catalytic reactions:

Preparation of sulfuric acid by contact method: SO2 + O2 → SO3; (catalyst – V2O5).

Enzymes are protein substances that catalyze biochemical reactions in the cells of living organisms.

Catalyst action explained by a decrease in the activation energy of the reaction. The catalyst interacts with reacting substances to form intermediate compounds, and this requires a lower activation energy and the reaction proceeds quickly.

The reaction A + B = AB without a catalyst proceeds slowly.

In the presence of a catalyst, the reaction occurs in two rapid stages:

AK + B = AB + K.

1.6. Chemical equilibrium, conditions for its displacement. Le Chatelier's principle

Reversible reactions – not flowing completely, they flow simultaneously in two opposite directions.

For example: N2 + 3H2 Û 2NH3

This reaction can proceed in two directions - the formation of ammonia and its decomposition.

The reversible reaction ends with the establishment chemical equilibrium is the state of a system of reacting substances when the rates of forward and reverse reactions are equal:

The state of equilibrium in a reversible system is characterized by an equilibrium constant.

Let's consider the reversible reaction aA + bB Û cC + dD.

The rate of direct reaction proceeding from left to right, according to the law of mass action, has the expression Vpr = k[A]a ∙ [B]b. The rate of the reverse reaction, proceeding from right to left, has the form Vrev = k[C]c ∙ [D]d. If the rates of forward and reverse reactions are equal: k[A]a ∙ [B]b = k[C]c ∙ [D]d. As a result, we obtain the expression for the equilibrium constant:

The equilibrium constant of a reversible reaction is 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.

The equilibrium constant equation shows that the concentrations of all substances participating in the reaction are related to each other. Changing the concentration of any of them will change the concentrations of all the others. As a result, new concentrations will be established, but the relationship between them corresponds to the equilibrium constant.

The principle of shifting equilibrium – Le Chatelier's principle:

If a system in a state of chemical equilibrium is affected in any way (by changing concentration, temperature or pressure), then the equilibrium will shift in the direction where this effect will decrease.

When the concentration of one of the substances increases, the equilibrium will shift towards the consumption of this substance.

As the temperature increases, the equilibrium shifts towards the endothermic reaction, and as the temperature decreases, towards the exothermic reaction.

As pressure increases, the equilibrium will shift towards decreasing volumes.

Questions for self-control

1) The concept of the rate of a chemical reaction. What factors influence the rate of a chemical reaction?

2) The influence of the concentration of reactants on the rate of a chemical reaction. Law of mass action. Task: How will the rate of the reaction 2NO + O2 → 2NO2 change if the NO concentration is doubled?

3) The influence of temperature on the rate of a chemical reaction. Van't Hoff's rule. Task: The temperature coefficient of a certain reaction is 2. How many times will the reaction rate increase when the temperature increases from 10 to 50 ºС?

4) Catalysts and their role in changing the rate of a chemical reaction.

5) Chemical equilibrium. Chemical equilibrium constant. Task: Equilibrium in the system H2 + J2 Û 2HJ was established at concentrations: = 0.025 mol/l, = 0.005 mol/l, = 0.09 mol/l. Calculate the equilibrium constant.

6) Displacement of chemical equilibrium, Le Chatelier's principle. Determine the direction of equilibrium shift in the system:

CO (g.) + O2 (g.) Û 2СО2 (g.) + 566 kJ

a) when the temperature rises; b) with increasing pressure; c) with increasing CO concentration; d) with increasing CO2 concentration?

BIBLIOGRAPHY

Main

1. Glinka, chemistry / .– M.: Integral-press, 2002. – 728 p.

2. Knyazev, D. A., Smarygin, chemistry / , . – M.: Bustard, 2004. – 529 p.

3. Ryazanova, G. E., Samokhina, on general and inorganic chemistry: textbook /,; Federal State Educational Institution of Higher Professional Education "Saratov State Agrarian University". – Saratov, 2007. – 192 p.

Additional

1. Egorov, fundamentals of inorganic chemistry. Short course for students of agricultural universities: textbook /. – Krasnodar: Lan, 2005. – 192 p.

2. Klinsky, G. D., Skopintsev, chemistry for biologists: a textbook for students. agricultural universities / , . – M.: Publishing house MCHA, 2001. – 384 p.

3. Magazines: “Chemistry and Life”, “Agrochemistry”, “Agrochemical Bulletin”, “Ecological Bulletin of Russia”.

Lecture 2

Energy of chemical processes

BIBLIOGRAPHICAL LIST

1. Artemenko, chemistry /. – M.: Higher School, 2005. – 605 p.

2. Handberg, chemistry / . – M.: Bustard, 2002. – 672 p.

3. Glinka, chemistry: textbook /. – M.: Integral-press, 2002. – 728 p.

4. Egorov, fundamentals of inorganic chemistry. Short course for students of agricultural universities: textbook /. – Krasnodar: Lan, 2005. – 192 p.

5. Zolotov, Yu. A., Vershinin, and the methodology of analytical chemistry / , . – M.: Publishing center “Academy”, 2007. – 464 p.

6. Knyazev, D. A., Smarygin, S. N. Inorganic chemistry / , . – M.: Bustard, 2004. – 529 p.

7. Fundamentals of organic chemistry / [etc.]. – M.: Bustard, 2006. – 560 p.

8. Fundamentals of analytical chemistry in 2 books: Textbook for universities/, [etc.] – M.: Higher school, 1999. – 351 p.

9. Ugai, chemistry / . – M.: Higher School, 2002. – 463 p.

Introduction…………………………………………………………………………………..3

Lecture 1. Chemical kinetics. Chemical equilibrium……………………..4

1.1. Rate of chemical reaction…………………………………………………………….4

1.2. Factors influencing the rate of a chemical reaction……………………….4

1.3. The influence of the concentration of reactants on the reaction rate.

Law of mass action……………………………………………………………………………….5

1.4. The effect of temperature on the reaction rate. Van't Hoff's rule.

Activation theory……………………………………………………………………………….6

1.5. Catalysis. Catalysts………………………………………………………..6

1.6. Chemical equilibrium, conditions for its displacement. Le Chatelier's principle……..7

Questions for self-control………………………………………………………..8

References………………………...……………………………………………………8

Lecture 2. Energy of chemical processes ………………………………….....9

……...

………

Bibliography………………………………………………………...10

Content………………………………………………………………………...…11

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 digit 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 valence. 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 will be equal to the sum of the 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 under normal conditions a mole of gas occupies a volume of 22.4 liters, 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 bonds determines such a unique property of each element, which is called valence. In its most general form, valency refers to 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 valence. 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 valence of each element, including in a chemical compound, is a necessary condition for successfully mastering a 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.

Ministry of Education and Science of the Russian Federation

State educational institution of higher professional education

"Ufa State Petroleum Technical University"

USPTU Student Library

A short course of lectures on the discipline “chemistry”

for students of non-chemical specialties

Under general editorship

Professor S.S. Zlotsky and Professor A.K. Mazitova

Approved by the Editorial and Publishing Council of USPTU

as a teaching aid

Authors: O.F. Bulatova, S.B. Denisova, L.N. Zorina, O.I. Mikhailenko, M.A. Molyavko, M.N. Nazarov, L.Z. Rolnik, L.E. Salova, L.G. Sergeeva, O.B. Chalova, A.T. Chanysheva, F.B. Shevlyakov (department of “General and Analytical Chemistry”); Yu.N.Biglova, E.A.Builova, D.R.Galieva, N.M.Shaimardanov (Department of Applied Chemistry and Physics)

Reviewers:

Head of the Department of Chemistry, Ufa State Aviation Technical University, Doctor of Chemical Sciences, Professor V.A. Dokichev

Head of the Department of General Chemistry, Ufa State Academy of Economics and Service, Candidate of Chemical Sciences

Associate Professor I.P. Zhurkina

K78 Short course of lectures on the discipline “Chemistry” / Yu.N. Biglova and others; under

total ed. S.S. Zlotsky and A.K. Mazitova. – Ufa: USNTU, 2010. – 69 p.

ISBN 978-5-7831-0955-3

Lectures on the discipline “Chemistry” are briefly given. The content of lectures corresponds to state educational standards. The modular principle of training is reflected, the content of practical and laboratory classes is given, and a list of basic literature for additional study of the material is provided. Designed for students of non-chemical specialties: AG, AT, AE, BAG, BAT, BAE, BMZ, BMP, BPG, BPS, BTE, VV, GF, DS, MZ, MP, MS, PG, PS, TE, EG, ES, ET, as well as AK, BOS, MH, OS, TS, TN full-time and part-time forms of education.

Technical University, 2010

Preface

The curriculum of students of non-chemical specialties at a technical university, in particular USPTU, includes the discipline “Chemistry”. For the vast majority of specialties in this subject, 12-20 lectures (24-40 hours), 3-5 practical classes (6-10 hours) and 10-15 laboratory classes (20-30 hours) are provided.

The content of the lecture material includes two main sections: structure, general (integral) properties of substances and properties of the most important elements. During practical classes, the key, fundamental issues of the program are discussed in detail in interactive mode, and attention is focused on the sections that are of greatest importance for the entire course. Laboratory work is devoted to the study of a wide range of problems of thermodynamics, kinetics, solutions, electrochemistry and transformations of the most important inorganic compounds. During the experiments, students gain the necessary skills and experience in working with chemicals and reagents. Taken together, classroom lessons, consultations, homework and independent work allow students to successfully master the program material and subsequently use the acquired knowledge in chemistry when studying special disciplines.

Currently, the Chemistry course has a large number of textbooks, study guides, workshops, collections of problems, etc., both in printed form and on electronic media. In 2005-2009, teachers of the departments of general arts and chemical engineering published extensive educational literature for students of non-chemical specialties (see the list of recommended literature).

At the same time, from teaching experience it follows that the lack of a manual containing basic information on the discipline in a concise, accessible form restrains the growth of student performance in the Chemistry course.

In this regard, teams of teachers from the departments of OAH and PCP of USPTU jointly prepared this manual*, the purpose of which is to systematize, simplify and make it easier for first-year students of non-chemical specialties to study and become familiar with the main content of the discipline “Chemistry”. A brief summary of each of the 23 lectures contains a description of basic provisions, terms, formulas and definitions. Questions for self-test and control are given, as well as links to 2 - 4 textbooks, where this section is presented in more detail and detail. At the end of the book there is an extended list of recommended educational literature and lists the main questions for tests and exams.

This manual does not replace existing textbooks and workshops, but, on the contrary, provides a more detailed and detailed introduction and study of sections of the program from the main textbooks. At the same time, the simplicity and accessibility of the textbook, in our opinion, allows students to get acquainted with the topics and content of lectures in advance, better imagine the course outline, and connect individual sections of the program with each other.

The authors, leading teachers of the departments of general education and science and philosophy, summarized and systematized the main parameters, goals and objectives of each lecture in a brief, abstract form. This allows students to minimize wasted time and concentrate on key issues and provisions of the discipline.

We believe that the manual will be useful and interesting to all students studying the discipline “Chemistry” in the first year, without exception, and will also be in demand by young beginning teachers and researchers for preparing for lectures, laboratory and practical classes. We recommend this manual to teachers, teachers of secondary schools, technical schools, colleges, as well as high school students interested in in-depth study of chemistry.

We express our deep gratitude to Associate Professor Builova E.A. and associate professor O.B. Chalova for preparing the manuscript for publication.

Professor Zlotsky S.S., head of the department of general education;

Professor Mazitova A.K., head of the Department of Philology and Philology.




Lecture 1. Quantum mechanical model of the structure of the atom……..……............
Lecture 2. Electronic configurations of atoms. Periodic Law. Periodic system D.I. Mendeleev…………………………………..
Lecture 3. Basic types of chemical bonds. Covalent bond…………...
Lecture 4. Theory of hybridization and geometry of molecules. Polarity and polarizability of covalent bonds and molecules………………………………...
Lecture 5. Intermolecular interactions. Hydrogen bond………….
Lecture 6. Chemical thermodynamics………………………………………………………...
Lecture 7. Chemical kinetics………………………………………………………………..
Lecture 8. Chemical equilibrium……………………………………………..
Lecture 9. Solutions. Methods of expressing the concentration of solutions. Properties of solutions……………………………………………………………...
Lecture 10. Dispersed systems. Surface phenomena…………………..
Lecture 11. Electrolyte solutions. Electrolytic dissociation………
Lecture 12. Dissociation of water. Dissociation of acids and bases. Hydrogen index…………………………………………………………
Lecture 13. Solubility product. Ion exchange reactions……….
Lecture 14. Hydrolysis of salts. Buffer solutions……………………………..
Lecture 15. Redox reactions………………………
Lecture 16. The concept of “Electrode potential”. Electrochemical processes……………………………………………………….………………
Lecture 17. Electrolysis of melts and solutions………………………………..
Lecture 18. General properties of metals………………………………….………
Lecture 19. Corrosion of metals. Methods of corrosion protection…………..........
Lecture 20. Metals of the main subgroup of group II. Hardness of water….........
Lecture 21. Structural metals. Aluminum. Chromium. Iron…………
Lecture 22. Polymers………………………………………………………………………………
Lecture 23. Chemical identification, substance analysis…………………...
Control questions

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 of various elements, and by changing the nature of this compound, one substance can be transformed into another. Democritus created atomic theory in almost its 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. Valence 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 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 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. 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 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 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.