Atomic-molecular science

The concept of matter and motion

Modern chemistry is one of the natural sciences, the subject of study of which is matter and is a system of individual chemical disciplines - inorganic, analytical, physical, organic, colloidal, etc.

The entire diverse world around us, the entire set of objects and phenomena are united by a common concept - matter, for which two forms of existence are known - substance and field.

Matter is a material formation consisting of particles that have their own mass or rest mass. Modern science Various types of material systems and the corresponding structural levels of matter are known. These include both elementary particles (electrons, protons, neutrons, etc.) and macroscopic bodies of various sizes (geological systems, planets, stars, star clusters, Galaxy, galaxy systems, etc.) modern knowledge of the structure of matter extends from 10 -14 cm before 10 28 cm(approximately 13 billion light years).

Unlike matter, a field is a material medium in which particles interact. For example, in an electronic field, interaction occurs between charged particles, and in a nuclear field, interaction occurs between protons and neutrons.

The universal forms of existence of matter are space and time, which do not exist outside of matter, just as there cannot be material objects that do not have spatiotemporal properties.

The fundamental and integral property of matter is movement - the way of its existence. The forms of motion of matter are very diverse, they are mutually related and can move from one to another. For example, the mechanical form of motion of matter can transform into an electrical form, an electrical form into a thermal form, etc. The measure of the movement of matter, its quantitative characteristic, is energy.

Definition of Chemistry

Various forms of motion of matter are studied by various sciences - physics, chemistry, biology, etc. Chemistry studies the chemical form of motion of matter, which is understood as a qualitative change in substances, the transformation of some substances into others. In this case, chemical bonds between the atoms that make up the substance are broken, re-emerged or redistributed. As a result chemical processes new substances with new physicochemical properties arise.

Thus, chemistry is a science that studies the processes of transformation of substances, accompanied by changes in composition, structure and properties, as well as mutual transitions between these processes and other forms of movement of matter.

The objects of study in chemistry are chemical elements and their connections. By studying the properties of substances and their transformations, chemistry reveals the laws of nature, cognizes matter and its movement. The study of chemistry as one of the most important fundamental natural sciences is necessary for the formation of a scientific worldview.

Atomic-molecular science

Atomic-molecular science was developed and first applied in chemistry by the great Russian scientist M.V. Lomonosov. The main provisions of his teaching are set out in the work “Elements of Mathematical Chemistry”. The essence of the teachings of M.V. Lomonosov comes down to the following.

1. All substances consist of “corpuscles” (as M.V. Lomonosov called molecules). 2. Molecules are made up of elements (atoms). 3. Particles - molecules and atoms are in continuous motion. 4. Molecules simple substances consist of identical atoms, molecules of complex substances - from different atoms.

This doctrine was later developed in the works of D. Dalton and J. Berzelius. The atomic-molecular theory in chemistry was finally established in the middle of the 19th century. At the International Congress of Chemists in Karlsruhe in 1860, definitions of the concepts of a chemical element, atom and molecule were adopted.

An atom is the smallest particle of a chemical element that has its chemical properties and is indivisible in chemical reactions.

A molecule is the smallest particle of a substance that has its chemical properties. Chemical properties molecules are determined by its composition and chemical structure.

All substances are divided into simple and complex.

A simple substance is made up of atoms of the same element.

A complex substance is made up of atoms of different elements. For example, copper oxide (II) formed by atoms of the elements copper and oxygen.

Just 100 years ago, the atom was viewed as an indivisible entity. However, in accordance with modern concepts, the atom has a complex structure and consists of three subatomic particles: protons, neutrons and electrons. Protons have a positive charge; neutrons have no charge and electrons have a negative charge. The charges on the proton and electron are the same in magnitude. Protons and neutrons together occupy a very small volume of an atom, called the nucleus. Most of the rest of the volume of an atom is the space in which electrons move. Because atoms have no net electrical charge, each atom contains an equal number of electrons and protons. The charge of the nucleus is determined by the number of protons.

A chemical element is a type of atoms characterized by the same nuclear charge and, accordingly, characterized by a certain set of properties. Atoms of the same element that differ in the number of neutrons, and therefore in mass, are called isotopes. Symbol 12 6 C or simply 12 C means a carbon atom with six protons and six neutrons. The number of protons in the nucleus of an atom is called the atomic number. Superscript (12) is called the mass number and indicates the total number of protons and neutrons in the nucleus of an atom.

The concept of “chemical element” cannot be identified with the concept of “simple substance”. A simple substance is characterized by a certain density, solubility, melting and boiling points, etc. These properties relate to a set of atoms and they are different for different simple substances.

A chemical element is characterized by a certain nuclear charge, isotopic composition, etc. The properties of an element relate to its individual atoms.

Complex substances are not made up of simple substances, but from elements. For example, water does not consist of the simple substances hydrogen and oxygen, but of the elements hydrogen and oxygen.

Many chemical elements form several simple substances that differ in structure and properties. This phenomenon is called allotropy, and the resulting substances are called allotropic modifications or modifications. Thus, the element oxygen forms two allotropic modifications: oxygen O 2 and ozone O 3; carbon element - three: diamond, graphite and carbine.

The chemical form of the movement of matter is studied and known by measuring the physical properties and physical quantities inherent in each substance. A physical quantity is, for example, the mass of a substance, its density, melting point. In chemistry, the concepts of relative atomic and molecular mass of a substance are widely used.

Relative atomic mass. The masses of atoms are extremely small. Thus, the mass of a hydrogen atom is 1.674×10 -27 kg, oxygen - 2.667×10 -26 kg. In chemistry, they traditionally use relative rather than absolute mass values. The unit of relative mass is the atomic mass unit (abbreviated a.e.m.), which represents 1/12 carbon atom mass - 12 , i.e. carbon isotope 6 C - 1.66×10 -27 kg. Since most elements have atoms with different masses, the relative atomic mass of a chemical element is a dimensionless quantity equal to the ratio of the average mass of an atom of the natural isotopic composition of the element to 1/12 mass of a carbon atom.

The relative atomic mass of an element is denoted by A r. For example,

Where 1.993·10 -26 kg– mass of a carbon atom.

Relative molecular weight. Relative molecular masses, like atomic masses, are expressed in atomic mass units. The relative molecular mass of a substance is a dimensionless quantity equal to the ratio of the average mass of a molecule of the natural isotopic composition of the substance to 1/12 carbon atom mass 12 6 C.

Relative molecular weight is denoted by M r. It is numerically equal to the sum of the relative atomic masses of all atoms that make up the molecule of a substance, and is calculated using the formula of the substance. For example, M r (H 2 O) will be composed of 2 A r (N)» 2; A r (O) = 1 × 16 = 16; M r (H 2 O) = 2 + 16 = 18.

Mol. In the international system of units (SI) The mole is taken as a unit of quantity of a substance. A mole is the amount of a substance containing so many structural or formulaic (FE) units (molecules, atoms, ions, electrons or others), how many atoms are there in 0.012 kg carbon isotope 12 6 C.

Knowing the mass of one carbon atom 12 C (1.993×10 -26 kg), calculate the number of atoms N A V 0.012 kg carbon.

Number of particles in 1 mole any substance is the same. It is equal 6.02×10 23 and is called Avogadro's constant (denoted N A, dimension 1/mol or mole -1). Obviously, in 2 mol carbon will be contained 2 × 6.02 × 10 23 atoms, in 3 mol - 3 × 6.02 × 10 23 atoms.

Molar mass. It is usually designated M. Molar mass is a value equal to the ratio of the mass of a substance to the amount of substance. It has a dimension kg/mol or g/mol. For example, M = m/n or M = m/n, Where m- mass in grams; n(nude) or n- amount of substance in moles, M- molar mass in g/mol- a constant value for each given substance. So, if the mass of a water molecule is equal to 2.99×10 -26 kg, then the molar mass M(H2O) = 2.99×10 -26 kg × 6.02×10 23 mol -1 = 0.018 kg/mol or 18 g/mol. In general, the molar mass of a substance, expressed in g/mol, is numerically equal to the relative atomic or relative molecular mass of this substance.

For example, relative atomic and molecular masses C, O2, H2S respectively equal 12, 32, 34, and their molar masses are respectively 12, 32, 34 g/mol.

Question 31. Modern preschool education for children with disabilities

Question 8. Non-ionizing electromagnetic fields and radiation. Laser radiation. Ionizing radiation.

Harmful and dangerous factors in the working environment in premises where modern computer equipment, telecommunication networks and various electronic devices are used.

Basic concepts of chemistry, laws of stoichiometry

Chemical atomism (atomic-molecular theory) is historically the first fundamental theoretical concept that forms the basis of modern chemical science. The formation of this theory took more than a hundred years and is associated with the activities of such outstanding chemists as M.V. Lomonosov, A.L. Lavoisier, J. Dalton, A. Avogadro, S. Cannizzaro.

Modern atomic-molecular theory can be presented in the form of a number of provisions:

1. Chemical substances have a discrete (discontinuous) structure. Particles of matter are in constant chaotic thermal motion.

2. The basic structural unit of a chemical substance is the atom.

3. Atoms in a chemical substance are bonded to each other to form molecular particles or atomic aggregates (supramolecular structures).

4. Complex substances (or chemical compounds) consist of atoms of different elements. Simple substances consist of atoms of one element and should be considered as homonuclear chemical compounds.

When formulating the main provisions atomic-molecular theory we had to introduce several concepts that need to be discussed in more detail, since they are fundamental in modern chemistry. These are the concepts of “atom” and “molecule,” more precisely, atomic and molecular particles.

Atomic particles include the atom itself, atomic ions, atomic radicals, and atomic radical ions.

An atom is the smallest electrically neutral particle of a chemical element, which is the carrier of its chemical properties, and consists of a positively charged nucleus and an electron shell.

Atomic ion is an atomic particle that has an electrostatic charge, but does not have unpaired electrons, for example, Cl - is a chloride anion, Na + is a sodium cation.

Atomic radical- an electrically neutral atomic particle containing unpaired electrons. For example, the hydrogen atom is actually an atomic radical - H × .

An atomic particle that has an electrostatic charge and unpaired electrons is called atomic radical ion. An example of such a particle is the Mn 2+ cation, which contains five unpaired electrons at the d-sublevel (3d 5).

One of the most important physical characteristics of an atom is its mass. Since the absolute value of the mass of an atom is negligible (the mass of a hydrogen atom is 1.67 × 10 -27 kg), chemistry uses a relative mass scale, in which 1/12 of the mass of a carbon atom of isotope-12 is chosen as a unit. Relative atomic mass is the ratio of the mass of an atom to 1/12 the mass of a carbon atom of the 12 C isotope.

It should be noted that in the periodic system D.I. Mendeleev presents the average isotopic atomic masses of elements, which are mostly represented by several isotopes that contribute to the atomic mass of an element in proportion to their content in nature. Thus, the element chlorine is represented by two isotopes - 35 Cl (75 mol.%) and 37 Cl (25 mol.%). The average isotopic mass of the element chlorine is 35.453 amu. (atomic mass units) (35×0.75 + 37×0.25).

Similar to atomic particles, molecular particles include molecules themselves, molecular ions, molecular radicals, and radical ions.

A molecular particle is the smallest stable collection of interconnected atomic particles, which is the bearer of the chemical properties of a substance. The molecule is devoid of electrostatic charge and has no unpaired electrons.

molecular ion is a molecular particle that has an electrostatic charge, but does not have unpaired electrons, for example, NO 3 - is a nitrate anion, NH 4 + is an ammonium cation.

molecular radical is an electrically neutral molecular particle containing unpaired electrons. Most radicals are reaction particles with a short lifetime (on the order of 10 -3 -10 -5 s), although fairly stable radicals are currently known. So methyl radical × CH 3 is a typical low-stable particle. However, if the hydrogen atoms in it are replaced by phenyl radicals, then a stable molecular radical triphenylmethyl is formed

Molecules with an odd number of electrons, such as NO or NO 2 , can also be considered highly stable free radicals.

A molecular particle that has an electrostatic charge and unpaired electrons is called molecular radical ion. An example of such a particle is the oxygen radical cation – ×O 2 + .

An important characteristic of a molecule is its relative molecular weight. Relative molecular mass (M r) is the ratio of the average isotopic mass of a molecule, calculated taking into account the natural content of isotopes, to 1/12 of the mass of a carbon atom of the 12 C isotope.

Thus, we have found out that the smallest structural unit of any chemical substance is an atom, or rather an atomic particle. In turn, in any substance, excluding inert gases, atoms are connected to each other by chemical bonds. In this case, the formation of two types of substances is possible:

· molecular compounds in which it is possible to isolate the smallest carriers of chemical properties that have a stable structure;

· compounds of a supramolecular structure, which are atomic aggregates in which atomic particles are linked by covalent, ionic or metallic bonds.

Accordingly, substances having a supramolecular structure are atomic, ionic or metallic crystals. In turn, molecular substances form molecular or molecular-ionic crystals. Substances that are under normal conditions in a gaseous or liquid state of aggregation also have a molecular structure.

In fact, when working with a specific chemical substance, we are not dealing with individual atoms or molecules, but with a collection of very large number particles, the levels of organization of which can be represented by the following diagram:

For a quantitative description of large arrays of particles, which are macrobodies, a special concept of “amount of matter” was introduced, as a strictly defined number of its structural elements. The unit of quantity of a substance is the mole. A mole is an amount of substance(n) , containing as many structural or formula units as there are atoms contained in 12 g of carbon isotope 12 C. Currently, this number is quite accurately measured and is 6.022 × 10 23 (Avogadro's number, N A). Atoms, molecules, ions, chemical bonds and other objects of the microworld can act as structural units. The concept of “formula unit” is used for substances with a supramolecular structure and is defined as the simplest relationship between its constituent elements (gross formula). In this case, the formula unit takes on the role of a molecule. For example, 1 mole of calcium chloride contains 6.022 × 10 23 formula units - CaCl 2.

One of important characteristics a substance is its molar mass (M, kg/mol, g/mol). Molar mass is the mass of one mole of a substance. The relative molecular mass and molar mass of a substance are numerically the same, but have different dimensions, for example, for water M r = 18 (relative atomic and molecular masses are dimensionless values), M = 18 g/mol. The amount of substance and molar mass are related by a simple relationship:

The basic stoichiometric laws that were formulated at the turn of the 17th and 18th centuries played a major role in the formation of chemical atomism.

1. Law of Conservation of Mass (M.V. Lomonosov, 1748).

The sum of the masses of the reaction products is equal to the sum of the masses of the substances that interacted. In mathematical form, this law is expressed by the following equation:

An addition to this law is the law of conservation of mass of an element (A. Lavoisier, 1789). According to this law During a chemical reaction, the mass of each element remains constant.

Laws M.V. Lomonosova and A. Lavoisier found a simple explanation within the framework of atomic theory. Indeed, during any reaction, the atoms of chemical elements remain unchanged and in constant quantities, which entails both the constancy of the mass of each element individually and the system of substances as a whole.

The laws under consideration are of decisive importance for chemistry, since they allow one to model chemical reactions using equations and perform quantitative calculations based on them. It should be noted, however, that the law of conservation of mass is not absolutely accurate. As follows from the theory of relativity (A. Einstein, 1905), any process that occurs with the release of energy is accompanied by a decrease in the mass of the system in accordance with the equation:

where DE is the released energy, Dm is the change in the mass of the system, c is the speed of light in vacuum (3.0×10 8 m/s). As a result, the equation of the law of conservation of mass should be written in the following form:

Thus, exothermic reactions are accompanied by a decrease in mass, and endothermic reactions are accompanied by an increase in mass. In this case, the law of conservation of mass can be formulated as follows: in an isolated system the sum of masses and reduced energies is a constant quantity. However, for chemical reactions whose thermal effects are measured in hundreds of kJ/mol, the mass defect is 10 -8 -10 -9 g and cannot be detected experimentally.

2. Law of Constancy of Composition (J. Proust, 1799-1804).

An individual chemical substance of molecular structure has a constant qualitative and quantitative composition, independent of the method of its preparation.. Compounds that obey the law of constant composition are called colorblind. Daltonides are all currently known organic compounds (about 30 million) and some (about 100 thousand) inorganic substances. Substances having a non-molecular structure ( Bertolides), do not obey this law and may have a variable composition, depending on the method of obtaining the sample. These include the majority (about 500 thousand) of inorganic substances. These are mainly binary compounds of d-elements (oxides, sulfides, nitrides, carbides, etc.). An example of a compound of variable composition is titanium(III) oxide, the composition of which varies from TiO 1.46 to TiO 1.56. The reason for the variable composition and irrationality of the Bertolide formulas are changes in the composition of some of the elementary cells of the crystal (defects in the crystal structure), which do not entail a sharp change in the properties of the substance. For Daltonids, such a phenomenon is impossible, since a change in the composition of the molecule leads to the formation of a new chemical compound.

3. Law of equivalents (I. Richter, J. Dalton, 1792-1804).

The masses of reacting substances are directly proportional to their equivalent masses.

where E A and E B are the equivalent masses of the reacting substances.

The equivalent mass of a substance is the molar mass of its equivalent.

An equivalent is a real or conditional particle that donates or gains one hydrogen cation in acid-base reactions, one electron in redox reactions, or interacts with one equivalent of any other substance in exchange reactions. For example, when metallic zinc reacts with an acid, one zinc atom displaces two hydrogen atoms, giving up two electrons:

Zn + 2H + = Zn 2+ + H 2

Zn 0 - 2e - = Zn 2+

Therefore, the equivalent of zinc is 1/2 of its atom, i.e. 1/2 Zn (conditional particle).

The number showing which part of the molecule or formula unit of a substance is its equivalent is called the equivalence factor - f e. Equivalent mass, or molar mass of equivalent, is defined as the product of the equivalence factor and the molar mass:

For example, in a neutralization reaction, sulfuric acid gives up two hydrogen cations:

H 2 SO 4 + 2KOH = K 2 SO 4 + 2H 2 O

Accordingly, the equivalent of sulfuric acid is 1/2 H 2 SO 4, the equivalence factor is 1/2, and the equivalent mass is (1/2) × 98 = 49 g/mol. Potassium hydroxide binds one hydrogen cation, so its equivalent is the formula unit, the equivalence factor is equal to one, and the equivalent mass is equal to the molar mass, i.e. 56 g/mol.

From the examples considered, it is clear that when calculating the equivalent mass, it is necessary to determine the equivalence factor. There are a number of rules for this:

1. The equivalence factor of an acid or base is equal to 1/n, where n is the number of hydrogen cations or hydroxide anions involved in the reaction.

2. The salt equivalence factor is equal to the quotient of unity divided by the product of the valency (v) of the metal cation or acid residue and their number (n) in the salt (stoichiometric index in the formula):

For example, for Al 2 (SO 4) 3 - f e = 1/6

3. The equivalence factor of an oxidizing agent (reducing agent) is equal to the quotient of unity divided by the number of electrons attached (donated) by it.

Attention should be paid to the fact that the same compound may have a different equivalence factor in different reactions. For example, in acid-base reactions:

H 3 PO 4 + KOH = KH 2 PO 4 + H 2 O f e (H 3 PO 4) = 1

H 3 PO 4 + 2KOH = K 2 HPO 4 + 2H 2 O f e (H 3 PO 4) = 1/2

H 3 PO 4 + 3KOH = K 3 PO 4 + 3H 2 O f e (H 3 PO 4) = 1/3

or in redox reactions:

KMn 7+ O 4 + NaNO 2 + H 2 SO 4 ® Mn 2+ SO 4 + NaNO 3 + K 2 SO 4 + H 2 O

MnO 4 - + 8H + + 5e - ® Mn 2+ + 4H 2 O f e (KMnO 4) = 1/5

Lecture 1

SUBJECT AND IMPORTANCE OF CHEMISTRY

1. Chemistry subject. Among the natural sciences that determine the foundation of engineering knowledge, chemistry occupies a leading position due to its information significance. As is known, about a quarter of the total volume of scientific and technical information is chemical.

Modern definition chemistry: a system of chemical sciences (organic, inorganic, analytical, physical chemistry, etc.), the main task of which is the study of chemical processes (reactions) of the formation and destruction of molecules (chemical bonds), as well as the relationships and transitions between these processes and others forms and movements of matter (electromagnetic fields and radiation, etc.).

Chemistry studies the composition, structure of substances of organic and inorganic origin, the ability of substances to interact and the phenomenon of the transition of chemical energy into heat, electricity, light, etc.

The importance of chemistry in the existence and development of humanity is enormous. Suffice it to say that not a single branch of production can do without chemistry. If you look at what surrounds a person in everyday life or at work, these are all the gifts and deeds of chemistry. Whole books have been written about the importance of chemistry in various industries, agriculture and medicine. Famous English physicist W. Ramsay said: “That nation, that country, which surpasses others in the development of chemistry, will surpass them in general material well-being.”

Basic laws of chemistry

Atomic-molecular science is the theoretical foundation of chemistry.

Substance is one of the forms of existence of matter. Matter consists of individual tiny particles - molecules, atoms, ions, which in turn have a certain internal structure. In other words, every substance is not something continuous, but consists of individual very small particles; the basis of atomic-molecular teaching is the principle of discreteness (discontinuity of structure) of matter. The properties of substances are a function of the composition and structure of the particles that form it. For most substances, these particles are molecules.

Molecule – the smallest particle of a substance that has its chemical properties. Molecules, in turn, are made up of atoms. Atom – the smallest particle of an element that has its chemical properties.

It is necessary to distinguish between the concepts of “simple (elementary) substance” and “chemical element”. In fact, each simple substance is characterized by certain physical and chemical properties. When a simple substance undergoes a chemical reaction and forms a new substance, it loses most of its properties. For example, iron, when combined with sulfur, loses its metallic luster, malleability, magnetic properties etc. In the same way, hydrogen and oxygen, which are part of water, are contained in water not in the form of gaseous hydrogen and oxygen with their characteristic properties, but in the form of elements - hydrogen and oxygen. If these elements are in " free state", i.e. are not chemically bonded to any other element, they form simple substances. A chemical element can be defined as a type of atom characterized by a certain set of properties . When atoms of the same element combine with each other, simple substances are formed, while the combination of atoms of different elements gives either a mixture of simple substances or a complex substance.

The existence of a chemical element in the form of several simple substances is called allotropy. Various simple substances formed by the same element are called allotropic modifications of this element. The difference between a simple substance and an element becomes especially clear when one encounters several simple substances consisting of the same element. There are allotropy of composition and allotropy of form. Atoms of the same element, arranged in different geometric orders (shape allotropy) or combined into molecules of different compositions (composition allotropy), form simple substances with different physical properties with similar chemical properties. Examples include:
oxygen and ozone, diamond and graphite. 2. Stoichiometric laws. Chemical equivalent. The basis of atomic-molecular science is the basic laws of chemistry, discovered at the turn of the 18th and 19th centuries.

Law of conservation of masses and energies, is the basic law of natural science. It was first formulated and experimentally substantiated by M.V. Lomonosov in 1756-59, later it was discovered and confirmed by A.L. Lavoisier: the mass of the resulting reaction products is equal to the mass of the initial reagents. IN mathematical form this can be written:

Where i, j- whole numbers, equal to the number reagents and products.

IN modern form This law is formulated as follows: in an isolated system, the sum of masses and energies is constant. The law of conservation of mass is based on the study of reactions between individual substances and quantitative chemical analysis.

The law of the relationship between mass and energy (A. Einstein). Einstein showed that there is a relationship between energy and mass, quantified by the equation:

E = mc 2 or Dm = D E/c 2 (2.2)

where E is energy; m – mass; With - speed of light. The law is fair for nuclear reactions, in which a huge amount of energy is released with small changes in mass (atomic explosion).

Law of constancy of composition (J.L. Proust, 1801-1808): no matter how this chemically pure compound is obtained, its composition is constant. Thus, zinc oxide can be obtained as a result of a wide variety of reactions:

Zn + 1/2 O 2 = ZnO; ZnСO 3 = ZnO + CO 2; Zn(OH) 2 = ZnO + H 2 O.

But a chemically pure ZnO sample always contains 80.34% Zn and 19.66% O.

The law of constancy of composition is fully satisfied for gaseous, liquid and a number of solid substances ( colorblind people), however, many crystalline substances retain their structure with a variable (within certain limits) composition ( berthollides). These include compounds of certain metals with each other, individual oxides, sulfides, and nitrides. Consequently, this law is applicable only for substances that, regardless of their state of aggregation, have a molecular structure. In compounds of variable composition, this law has limits of applicability, especially for substances in the solid state, since the carrier of properties in a given state is not a molecule, but a certain set of ions of different signs, called a phase (a homogeneous part of a heterogeneous system, limited by an interface) , or, in other words, crystal lattices solids have defects (vacancies and inclusions of nodes).

Law of equivalents (Richter, 1792-1800): chemical elements combine with each other in mass ratios proportional to their chemical equivalents:

All stoichiometric calculations are carried out on the basis of this law.

Chemical equivalent of an element is the amount of it that combines with 1 mole (1.008 g) of hydrogen atoms or replaces the same number of hydrogen atoms in chemical compounds.

The concept of equivalents and equivalent masses also applies to complex substances. Equivalent to a complex substance is the amount of it that reacts without a residue with one equivalent of hydrogen or, in general, with one equivalent of any other substance.

Calculation of equivalents of simple and complex substances:

Where A r – atomic mass of the element; M A– molecular weight of the compound.

The law of multiple ratios (D. Dalton, 1808). If two elements form several chemical compounds with each other, then the amount of one of them, related to the same amount of the other, is related as small integers.

Avogadro's Law (1811). This is one of the basic laws of chemistry: equal volumes of gases under the same physical conditions (pressure and temperature) contain the same number of molecules.

A. Avogadro established that the molecules of gaseous substances are diatomic, not H, O, N, Cl, but H 2, O 2, N 2, Cl 2. However, with the discovery of inert gases (they are monatomic), exceptions were discovered.

First consequence: 1 mole of any gas under normal conditions has a volume equal to 22.4 liters.

Second consequence: the densities of any gases are related to their molecular masses: d 1 / d 2 = M 1 / M 2.

Avogadro's constant is the number of particles in 1 mole of a substance 6.02 × 10 23 mol -1.

The explanation of the basic laws of chemistry in the light of atomic-molecular theory lies in its postulates:

1) atoms are the smallest particles of matter that cannot be divided into their component parts (by chemical means) or converted into each other, or destroyed;

2) all atoms of one element are identical and have the same mass (if you do not take into account the existence of isotopes, see lecture 3);

3) atoms of different elements have different masses;

4) during a chemical reaction between two or a large number elements, their atoms are connected to each other in small integer ratios;

5) the relative masses of the elements that combine with each other are directly related to the masses of the atoms themselves, i.e. if 1 g of sulfur combines with 2 g of copper, this means that each copper atom weighs twice as much as a sulfur atom;

In a word, chemistry is “controlled” by integers, which is why all these laws are called stoichiometric. This is the triumph of atomic-molecular science.

3. Atomic and molecular masses. Mol. Let's consider in what units molecular and atomic masses are expressed. In 1961, a unified scale of relative atomic masses was adopted , which is based on 1/12 of the mass of an atom of the carbon isotope 12 C, called the atomic mass unit (amu). In accordance with this, at present the relative atomic mass (atomic mass) of an element is the ratio of the mass of its atom to 1/12 of the mass of a 12 C atom.

Similarly, the relative molecular weight (molecular weight) of a simple or complex substance is the ratio of the mass of its molecule
to 1/12 of the mass of a 12 C atom. Since the mass of any molecule is equal to the sum of the masses of its constituent atoms, the relative molecular mass is equal to the sum of the corresponding relative atomic masses. For example, the molecular weight of water, the molecule of which contains two hydrogen atoms and one oxygen atom, is equal to: 1.0079 × 2 + 15.9994 = 18.0152.

Along with units of mass and volume, in chemistry they also use a unit of quantity of a substance, called the mole. Mole – the amount of a substance containing as many molecules, atoms, ions, electrons or other structural units as there are atoms in 12 g of the carbon isotope 12 C.

The amount of a substance in moles is equal to the ratio of the mass of the substance m to its molecular weight M:

n= m/M. (2.8)

Molar mass ( M) is usually expressed in g/mol. The molar mass of a substance, expressed in g/mol, has the same numerical value as its relative molecular (atomic) mass. Thus, the molar mass of atomic hydrogen is 1.0079 g/mol, molecular hydrogen is 2.0158 g/mol.

Dependence of gas volume on pressure and temperature can be described equation of state of an ideal gas pV = RT, valid for one mole of gas, and taking into account the number of moles it becomes the famous equation
Clapeyron–Mendeleev:

pV= n RT (2.9)

Where R– universal gas constant (8.31 J/mol×K).

Using this equation and the second corollary of Avogadro's law, using simple measuring instruments (thermometer, barometer, scales), at the end of the 19th century. The molecular masses of many volatile simple and complex organic and inorganic substances were determined. In 1860, at the 1st International Congress of Chemists (Karlsruhe, Germany), classical definitions of basic concepts were adopted: atom, molecule, element, etc., systematics and classification of the main types of reactions and classes of chemical compounds were carried out.

4. Main classes of inorganic compounds. Classification of simple and complex chemical substances is based on consideration of the reagents and products of one of the main chemical reactions - the neutralization reaction. The foundations of this classification were laid by I.Ya. Berzelius in 1818, later it was significantly clarified and supplemented.

Alchemists also combined a number of simple substances with similar physical and chemical properties called metals . Typical metals are characterized by malleability, metallic luster, high thermal and electrical conductivity; in terms of their chemical properties, metals are reducing agents. The remaining simple substances were combined into the class non-metals (metalloids ). Nonmetals have more diverse physical and chemical properties. When simple substances interact with oxygen, they form oxides . Metals form basic oxides, non-metals – acidic . In the reaction of such oxides with water, respectively, grounds And acids . Finally, the neutralization reaction of acids and bases leads to the formation salts . Salts can also be obtained by the interaction of basic oxides with acidic oxides or acids, and acidic oxides with basic oxides or bases (Table 1).

Table 1

Chemical properties of the main classes of inorganic compounds

It should be emphasized that only those basic oxides that form water-soluble bases react directly with water - alkalis . Water-insoluble bases (for example, Cu(OH) 2) can be obtained from oxides only in two stages:

CuO + H 2 SO 4 = CuSO 4 + H 2 O, CuSO 4 + 2NaOH = Cu(OH) 2 ¯ + Na 2 SO 4.

The classification of oxides is not limited to basic and acidic. A number of oxides and their corresponding hydroxides exhibit dual properties: they react with acids as bases and with bases as acids (in both cases, salts are formed). Such oxides and hydroxides are called amphoteric :

Al 2 O 3 +6HCl=2AlCl 3 +3H 2 O, Al 2 O 3 +2NaOH=2NaAlO 2 +H 2 O (fusion of solids),

Zn(OH) 2 + 2HCl = ZnCl 2 + 2H 2 O, Zn(OH) 2 + 2NaOH = Na 2 (in solution).

Some oxides cannot be matched to their corresponding acid or base. Such oxides are called non-salt-forming , for example, carbon monoxide (II) CO, nitrogen oxide (I) N 2 O. They do not participate in acid-base interactions, but can enter into other reactions. So, N 2 O is a strong oxidizing agent, CO is a good reducing agent. Sometimes acidic, basic and amphoteric oxides are combined into a class salt-forming .

Among the acids stand out oxygen-free – for example, hydrogen chloride (hydrochloric) HCl, hydrogen sulfide H 2 S, hydrogen cyanide (hydrocyanide) HCN. In terms of acid-base properties, they do not differ from oxygen-containing acids There are also substances that have basic properties, but do not contain metal atoms, for example, ammonium hydroxide NH 4 OH - a derivative of ammonia NH 3.

The names of acids are derived from the element that forms the acid. In the case of oxygen-free acids, the suffix “o” and the word “hydrogen” are added to the name of the element (or group of elements, for example, CN - cyanogen) that forms the acid: H 2 S - hydrogen sulfide, HCN - hydrogen cyanide.

The names of oxygen-containing acids depend on the degree of oxidation of the acid-forming element. The maximum degree of oxidation of an element corresponds to the suffix “... n (th)” or “... ov (th)”, for example, HNO 3 - nitric acid, HClO 4 - perchloric acid, H 2 CrO 4 - chromic acid. As the oxidation state decreases, the suffixes change in the following sequence: “...ovat(aya)”, “...ist(aya)”, “...ovatist(aya)”; for example, HClO 3 is hypochlorous, HClO 2 is chlorous, HOCl is hypochlorous acid. If an element forms acids in only two oxidation states, then the suffix “...ist(aya)” is used to name the acid corresponding to the lowest oxidation state of the element; for example, HNO 2 is nitrous acid. Acids containing the group of atoms -O-O- in their composition can be considered as derivatives of hydrogen peroxide. They are called peroxoacids (or peracids). If necessary, after the prefix “peroxo”, a numerical prefix is ​​placed in the name of the acid indicating the number of atoms of the acid-forming element that are part of the molecule, for example: H 2 SO 5, H 2 S 2 O 8.

Among the compounds, an important group is formed grounds (hydroxides), i.e. substances containing hydroxyl groups OH - . The names of hydroxides are formed from the word “hydroxide” and the name of the element in the genitive case, after which, if necessary, the oxidation state of the element is indicated in Roman numerals in parentheses. For example, LiOH is lithium hydroxide, Fe(OH) 2 is iron (II) hydroxide.

Characteristic property bases is their ability to react with acids, acidic or amphoteric oxides to form salts, for example:

KOH + HCl = KCl + H 2 O,

Ba(OH) 2 + CO 2 = BaCO 3 + H 2 O

2NaOH + Al 2 O 3 = 2NaAlO 2 + H 2 O

From the point of view of the protolytic (proton) theory, bases are considered to be substances that can be proton acceptors, i.e. capable of attaching hydrogen ions. From this point of view, bases should include not only basic hydroxides, but also some other substances, for example ammonia, the molecule of which can add a proton, forming an ammonium ion:

NH 3 + H + = NH 4 +

Indeed, ammonia, like basic hydroxides, is capable of reacting with acids to form salts:

NH 3 + HCl = NH 4 Cl

Depending on the number of protons that can attach to the base, there are single-acid bases (for example, LiOH, KOH, NH 3), di-acid ones [Ca(OH) 2, Fe(OH) 2], etc.

Amphoteric hydroxides (Al(OH) 3, Zn(OH) 2) are capable of dissociating in aqueous solutions both as acids (with the formation of hydrogen cations) and as bases (with the formation of hydroxyl anions); they can be both donors and acceptors of protons. Therefore, amphoteric hydroxides form salts when reacting with both acids and bases. When interacting with acids, amphoteric hydroxides exhibit the properties of bases, and when interacting with bases, the properties of acids:

Zn(OH) 2 + 2HCl = ZnСl 2 + 2H 2 O,

Zn(OH) 2 + 2NaOH = Na 2 ZnO 2 + 2H 2 O.

There are compounds of elements with oxygen, which in composition belong to the class of oxides, but in their structure and properties belong to the class of salts. These are so-called peroxides, or peroxides. Peroxides are salts of hydrogen peroxide H 2 O 2, for example, Na 2 O 2, CaO 2. Characteristic feature The structure of these compounds is the presence in their structure of two interconnected oxygen atoms (“oxygen bridge”): -O-O-.

Salts upon electrolytic dissociation they form in aqueous solution cation K + and anion A – . Salts can be considered as products of complete or partial replacement of hydrogen atoms in an acid molecule with metal atoms or as products of complete or partial replacement of hydroxyl groups in a basic hydroxide molecule with acidic residues.

The neutralization reaction may not proceed completely. In this case, with an excess of acid, sour salts, with excess base - basic (salts formed at an equivalent ratio are called average ). It is clear that acid salts can be formed only by polyacid acids, basic salts - only by polyacid bases:

Ca(OH) 2 + 2H 2 SO 4 = Ca(HSO 4) 2 + 2H 2 O,

Ca(OH) 2 + H 2 SO 4 = CaSO 4 + 2H 2 O,

2Ca(OH) 2 + H 2 SO 4 = (CaOH) 2 SO 4 + 2H 2 O.

Among the diversity and huge number of chemical reactions, their classification has always been present. Thus, taking into account the development of chemistry, three main types of chemical reactions are distinguished:

1) acid-base balance, special cases - neutralization, hydrolysis, electrolytic dissociation of acids and bases;

2) redox with a change in the oxidation state of an atom, ion, molecule. In this case, the stages of oxidation and reduction are distinguished as parts of one process of electron loss and gain;

3) complex formation - the attachment of a certain number of molecules or ions to the central atom or ion of the metal, which is a complexing agent, and the former are ligands, the number of which is characterized by the coordination number (n).

According to these types of chemical reactions, chemical compounds are classified: acids and bases, oxidizing agents and reducing agents, complex compounds and ligands.

In a more modern interpretation, taking into account electronic structure atoms and molecules, reactions of the first type can be defined as reactions involving and transfer of a proton, reactions of the second type – with the transfer of an electron, reactions of the third type – with the transfer of a lone pair of electrons. The quantitative measure of reactions of the first type is, for example, pH, the second - potential (E, B), potential difference (Δφ, V), and the third - for example, the implementation of a certain coordination number (n) of chemical (donor-acceptor) bonds, energy stabilization of the ligand field of the central ion – complexing agent
(ΔG, kJ/mol), stability constant.

Atomic structure

1. Development of ideas about the structure of the atom. If, as a result of some kind of global catastrophe, all the accumulated scientific knowledge would be destroyed and only one phrase would be passed down to future generations, which statement, composed of the fewest words, would convey the most information? This question was posed by the famous American physicist, Nobel laureate Richard Feynman and he himself gave the following answer: this is the atomic hypothesis. All bodies consist of atoms - small bodies that are in continuous motion, attracted at a short distance, but repelled if one of them is pressed more closely to the other. However, the ancient Greek philosopher Democritus, who lived 400 years BC, could essentially agree with this statement. Modern people know more about atoms if, unlike the ancient Greeks, they were able to create atomic bombs and nuclear power plants based on their knowledge.

Until the end of the 19th century. believed the atom to be an indivisible and unchanging particle. But then phenomena were discovered that were inexplicable from this point of view. Electrochemical research G. Davy, M. Faraday showed that an atom can carry positive and negative charges as they are deposited at the cathode or anode of the electrolyser. Hence the corpuscular nature of the electric charge.

By improving methods of excitation of gases to obtain their spectra, W. Crooks discovered the so-called cathode rays (a phenomenon implemented in modern televisions). When an electric current passes through a rarefied gas enclosed in a tube, a stream of weak light—a cathode ray—emanates from the negative pole (cathode). The cathode ray imparts a negative charge to the bodies on which it falls and is deflected towards positively charged bodies close to the tube. Therefore, the cathode ray is a stream of negatively charged particles.

The phenomena of thermal emission and photoemission were also discovered ( A.G. Stoletov), consisting in knocking out negatively charged particles under the influence of temperature and light quanta, confirming the fact that the atom contains negatively charged particles. A.A. Becquerel discovered the phenomenon of radioactivity. Spouses Curie showed that the flow radioactive radiation is inhomogeneous and can be separated by electric and magnetic fields. The total radiation entering the capacitor is divided into three parts: a-rays (He 2+) are slightly deflected towards the negative plate of the capacitor, b-rays (electron flow) are strongly deflected towards the positive plate of the capacitor, g-rays (electromagnetic waves) are not deflected at all electric or magnetic field.

And finally, the discovery of X-rays Conrad Roentgen showed that the atom is complex and consists of positive and negative particles, the smallest of which H. Thomsen called the electron. Moreover, R.S. Mulliken measured its charge e= -1.6×10 -19 C (minimum possible, i.e. elementary) and found the mass of the electron m= 9.11×10 -31 kg.

The neutrality of an atom in the presence of electrons in it led to the conclusion that there was a region in the atom that carried a positive charge. The question remains open about the location or placement of electrons and supposed positive charges in atoms, i.e. question about the structure of the atom. Based on these studies, in 1903 H. Thomsen proposed a model of the atom, which was called “raisin pudding”, the positive charge in the atom is distributed evenly with a negative charge interspersed with it. But further research showed the inconsistency of this model.

E. Rutherford(1910) passed a stream of a-rays through a layer of material (foil), measuring the deflection of individual particles after passing through the foil. Summarizing the results of his observations, Rutherford established that a thin metal screen was partly transparent to alpha particles, which, passing through the sheet, either did not change their path or were deflected at small angles. Individual a-particles were thrown back, like a ball from a wall, as if they had encountered an insurmountable obstacle on their way. Since a very small number of a-particles passing through the foil were thrown back, this obstacle must occupy a volume in the atom, immeasurably smaller even in comparison with the atom itself, and it must have a large mass, since otherwise the a-particles from it would not ricochet. Thus, a hypothesis appeared about the nucleus of an atom, in which almost the entire mass of the atom and all the positive charge are concentrated. In this case, the deviations of the path of most alpha particles by small angles under the influence of electrostatic repulsion forces from the atomic nucleus become clear. Later it was found that the diameter of the nucleus is about 10 -5 nm, and the diameter of the atom is 10 -1 nm, i.e. the volume of the nucleus is 10 12 times less than the volume of the atom.

In the atomic model proposed by Rutherford, a positively charged nucleus is located at the center of the atom, and electrons move around it, the number of which is equal to the charge of the nucleus or serial number element, like the planets around the Sun (planetary model of the atom). The nuclear model developed by Rutherford was a major step forward in understanding the structure of the atom. It has been confirmed by a large number of experiments. However, in some respects the model contradicted well-established facts. Let us note two such contradictions.

First, Rutherford's planetary model of the atom could not explain the stability of the atom. According to the laws of classical electrodynamics, an electron, moving around a nucleus, inevitably loses energy. As the energy reserve of an electron decreases, the radius of its orbit must continuously decrease and, as a result, fall onto the nucleus and cease to exist. Physically, an atom is a stable system and can exist without destruction for an extremely long time.

Secondly, Rutherford's model led to incorrect conclusions about the nature of atomic spectra. The spectra of alkali metals turn out to be similar to the spectrum of atomic hydrogen, and their analysis led to the conclusion that the atoms of each alkali metal contain one electron, weakly bound to the nucleus compared to the remaining electrons. In other words, in an atom, electrons are not located at the same distance from the nucleus, but in layers.

Atomic spectra are obtained by passing the radiation of excited atoms (in a flame with high temperature or other means) through a special optical device(prism, system of prisms or diffraction gratings), which decomposes complex radiation into monochromatic components with a certain wavelength (l) and, accordingly, with a certain frequency of oscillations of electromagnetic radiation: n = With/l, where c– speed of light. Each monochromatic beam is registered at a specific location in the receiving device (photoplate, etc.). The result is a spectrum of this radiation. Atomic spectra consist of individual lines - these are line spectra.

Each type of atom is characterized by a strictly defined arrangement of lines in the spectrum that are not repeated in other types of atoms. This is the basis of the method of spectral analysis, with the help of which many elements were discovered. The linearity of atomic spectra contradicted the laws of classical electrodynamics, according to which the spectrum of atoms should be continuous as a result of the continuous emission of energy by the electron.

2. Model of the structure of the Bohr hydrogen atom. Since the laws of classical electrodynamics turned out to be inapplicable to describe the behavior of an electron in an atom, Niels Bohr first formulated postulates based on the laws of quantum mechanics.

1. There are orbits in the hydrogen atom, moving along which the electron does not emit. They are called stationary.

2. Emission or absorption of energy occurs as a result of the transition of an electron from one stationary orbit to another. Orbits distant from the nucleus are characterized by a large supply of energy. During the transition from lower to higher orbits, the atom goes into an excited state. But he may not remain in this state for long. It emits energy and returns to its original ground state. In this case, the energy of the radiation quantum is equal to:

h n= E n – Ek,

Where n And k- whole numbers.

3. Basic principles of wave (quantum) mechanics. The explanation of wave (spectral) properties arose simultaneously with quantum mechanical concepts in the theory of atomic structure. The premise was the theory Plank body radiation He showed that energy changes do not occur continuously (according to the laws of classical mechanics), but spasmodically, in portions that were called quanta. The quantum energy is determined by Planck’s equation: E = h n, where h – Planck's constant is equal to 6.63×10 –34 J×s,
n – radiation frequency. It turns out that the electron has corpuscular properties (mass, charge) and wave properties - frequency, wavelength.

Due to this Louis de Broglie put forward the idea of ​​particle-wave dualism . Moreover, particle-wave dualism is characteristic of all objects of the micro- and macroworld; only for macroscopic objects one of the sets of properties predominates, and they are spoken of as particles or waves, and for elementary particles both properties appear together. De Broglie's equation shows the relationship between particle momentum and wavelength: l = h/p = h/m u. Thus, an electron rotating around a nucleus can be assigned a certain wavelength.

According to these ideas, an electron is a cloud, smeared in the volume of an atom, having different densities. Consequently, to describe the position of an electron in an atom, it is necessary to introduce a probabilistic description of the electron density in an atom, taking into account its energy and spatial geometry.

4. Quantum numbers. Orbitals. Four quantum numbers have been proposed to explain the electronic structure of the hydrogen atom n, l, m l, s, characterizing the energy state and behavior of an electron in an atom. These numbers uniquely characterize the state of the electron of any atom of the Periodic Table of elements. For each electron, they collectively have different values.

Principal quantum number n characterizes the energy and size of electron clouds. It takes values ​​for the ground states of atoms 1-8 and, in principle, to infinity. Its physical meaning as an energy level number is the energy value of an electron in an atom and, as a consequence, the size of the atom. At P=1 electron is in the first energy level with total minimum energy, etc. When increasing P total energy increases. The energy of each energy level can be estimated using the formula: E = - 1 / 13.6 ×n 2. Energy levels are usually designated by letters as follows:

Meaning ( n)
Designations	K	L	M	N	Q

Side, orbital(or azimuthal)quantum number l characterizes the shape of electron orbitals (clouds) around an atom and determines the change in energy within the energy level, i.e. characterizes energy sublevel. Each shape of the electron cloud corresponds to a certain value of the mechanical momentum of the electron, determined by the side quantum number l, which vary from 0 to P–1: P=1, l=0; P=2, l=0, l=1; P=3, l=0,l=1, l=2, etc. Energy sublevels depending on l denoted by letters:

Values ​​( l)
Notation ( V)	s	p	d	f	g	h

Those electrons that are in the s level are called s- electrons,
on p level – p- electrons, on d level – d- electrons.

The energy of electrons depends on the external magnetic field. This dependence is described by the magnetic quantum number. Magnetic quantum number m l indicates orientation in space electron orbital(cloud). An external electric or magnetic field changes the spatial orientation of electron clouds, and energy splitting occurs.
sublevels. Number m l varies from – l, 0, +l and may have (2× l+1) values:

The combination of three quantum numbers uniquely describes the orbital. It is designated as a “square” - . An electron as a particle experiences rotation around its own axis - clockwise and counterclockwise. It is described spin quantum number s(m s), which takes values ​​±1/2. The presence of electrons in an atom with oppositely directed spins is indicated as “arrows”. So the four sets of quantum numbers describe the energy of electrons.

5. Multielectron atoms. Determination of the number of electrons at levels and sublevels. In multi-electron atoms, the electron arrangement in accordance with a set of quantum numbers is governed by two postulates.

Pauli principle: in an atom there cannot be two electrons that have four identical quantum numbers (otherwise they are indistinguishable, the minimum energy difference is in the spins). As a consequence, in one electron cell in an orbital there can be no more than two electrons with oppositely directed spins.

Filling of cells with electrons is carried out in accordance with Hund's rule. Electrons fill s-, p-, d-, f- orbitals in such a way that the total spin is maximum, or, in other words, electrons tend to fill vacant (empty) orbitals, and only then pair (according to Pauli):

Taking into account the principles of quantum chemistry, it is possible to construct the electronic configuration of any atom, as follows from table. 2, from which we derive formulas for determining the number of electrons at the 2n 2 level, at the 2(2 l+1). The number of orbitals is equal to the number of values ​​of m (m=1, m=2, m=3).

The filling of sublevels with electrons is carried out in accordance with Klechkovsky's rule. The filling of energy levels occurs in increasing order of the sum of the main and secondary quantum numbers n+l.

If this amount has same values, then filling is carried out in ascending order n. Sublevels are filled in order of increasing energy:

1s<< 2s << 2p << 3s << 3p << 4s £ 3d << 4p << 5s £ 4d << 5p << 6s £ 4f £ 5d…

Table 2 - Electronic configurations of atoms

Which level is filled next? 4s»3d in energy. 4s n=3, d=2, sum is 5, n=4, s=0, sum = 4, i.e. 4s are being filled, etc. Energy 5s » 4d, the sum is 5 and 6, therefore 5s is filled first, then 4d. The energy is 6s » 5d » 4f, the sum is 6, 7 and 7. 6s is filled in at the beginning. The main quantum number is smaller for 4f, therefore, this sublevel is filled further, followed by 5d.

The electronic configuration of an atom is written as a formula, where the number of electrons in a sublevel is indicated by a superscript. For example, for aluminum you can write the electron configuration formula as 1s 2 2s 2 2p 6 3s 2 3p 1. This means that there are 2, 2, 6, 2, 1 electrons in the 1s, 2s, 2p, 3s, 3p sublevels.

In a multielectron unexcited atom, electrons occupy orbitals with minimal energies. They interact with each other: electrons located on the internal energy levels screen (obscure) electrons located on the external levels from the action of the positive nucleus. This influence determines the change in the sequence of increasing orbital energy compared to the sequence of increasing orbital energy in the hydrogen atom.

It should be noted that for elements with fully or half filled d- And f-deviations from this rule are observed at sublevels. For example, in the case of the copper atom Cu. The electronic configuration [Аr] 3d 10 4s 1 corresponds to lower energy than the configuration [Аr] 3d 9 4s 2 (the symbol [Аr] means that the structure and filling of internal electronic levels is the same as in argon). The first configuration corresponds to the ground state, and the second to the excited state.

Chemical bond

1. The nature of the chemical bond. Theories to explain chemical bonding are based on Coulomb, quantum and wave interactions of atoms. First of all, they must explain the gain in energy during the formation of molecules, the mechanism of the formation of a chemical bond, its parameters and the properties of the molecules.

The formation of a chemical bond is an energetically favorable process and is accompanied by the release of energy. This is confirmed by a quantum mechanical calculation of the interaction of two hydrogen atoms during the formation of a molecule (Heitler, London). Based on the calculation results, the dependence of the potential energy of the system is derived E on the distance between hydrogen atoms r(Fig. 4).

Rice. 4. Dependence of energy on internuclear distance.

When atoms come closer together, electrostatic forces of attraction and repulsion arise between them. If atoms with antiparallel spins come together, the attractive forces initially predominate, so the potential energy of the system decreases (curve 1). Repulsive forces begin to dominate at very small distances between atoms (nuclear interactions). At a certain distance between the atoms r 0, the energy of the system is minimal, so the system becomes most stable, a chemical bond occurs and a molecule is formed. Then r 0 is the internuclear distance in the H2 molecule, which is the length of the chemical bond, and the decrease in the energy of the system at r 0 is the energy gain during the formation of a chemical bond (or the energy of a chemical bond E sv). It should be noted that the energy of dissociation of a molecule into atoms is equal to E sv in magnitude and opposite in sign.

For a quantum mechanical description of a chemical bond, two complementary methods are used: the valence bond (VB) method and the molecular orbital (MO) method.

2. Valence bond (VB) method. Covalent bond. The main universal type of chemical bond is a covalent bond. Let us consider the mechanism of formation of a covalent bond using the BC method (using the example of the formation of a hydrogen molecule):

1. A covalent bond between two interacting atoms is carried out by the formation of a common electron pair. Each atom contributes one unpaired electron to form a common electron pair:

N·+·N ® N : N

Thus, according to the BC method, the chemical bond is two-center and two-electron.

2. A common electron pair can only be formed through the interaction of electrons with antiparallel spins:

Н+¯Н ® Н¯Н.

3. When a covalent bond is formed, electron clouds overlap:

This is confirmed by the experimentally determined value of the internuclear distance in the H 2 molecule, r = 0.074 nm, which is significantly less than the sum of the radii of two free hydrogen atoms, 2r = 0.106 nm.

In the region of cloud overlap, the electron density is maximum, i.e. the probability of two electrons being in the space between nuclei is much greater than in other places. A system arises in which two nuclei electrostatically interact with a pair of electrons. This leads to a gain in energy, and the system becomes more stable, and a molecule is formed. The more the electron clouds overlap, the stronger the covalent bond.

Donor-acceptor mechanism of covalent bonds. The formation of a covalent bond can occur due to the own lone pair of electrons of one atom (ion) - donor and a free atomic orbital of another atom (ion) – acceptor. This mechanism of covalent bond formation is called donor-acceptor.

The formation of the ammonia molecule NH 3 occurs by sharing three unpaired electrons of a nitrogen atom and one unpaired electron of three hydrogen atoms to form three common electron pairs. In the ammonia molecule NH 3, the nitrogen atom has its own lone pair of electrons. The 1s atomic orbital of the hydrogen ion H + does not contain electrons (vacant orbital). When the NH 3 molecule and the hydrogen ion approach each other, the lone electron pair of the nitrogen atom and the vacant orbital of the hydrogen ion interact to form a chemical bond via the donor-acceptor mechanism and the NH 4 + cation. Due to the donor-acceptor mechanism, the valence of nitrogen is B = 4.

The formation of chemical bonds by the donor-acceptor mechanism is a very common phenomenon. Thus, a chemical bond in coordination (complex) compounds is formed according to the donor-acceptor mechanism (see lecture 16).

Let us consider, within the framework of the BC method, the characteristic properties of a covalent bond: saturation and directionality.

Saturation Bonding is the ability of an atom to participate in only a certain number of covalent bonds. Saturation is determined by the valency of the atom. Saturation characterizes the number (number) of chemical bonds formed by an atom in a molecule, and this number is called covalency (or, as in the MO method, bond order).

The valency of an atom is a concept widely used in the study of chemical bonds. Valency refers to affinity, the ability of an atom to form chemical bonds. Quantitative assessment of valency may differ for different ways of describing a molecule. According to the BC method, the valence of an atom (B) is equal to the number of unpaired electrons. For example, from the electron cell formulas of oxygen and nitrogen atoms it follows that oxygen is divalent (2s 2 2p 4), and nitrogen is trivalent (2s 2 2p 3).

Excited state of atoms (v.s.). Paired electrons of the valence level, when excited, can be unpaired and transferred to free atomic orbitals (AO) of a higher sublevel within a given valence level. For example, for beryllium in an unexcited state (n.s.) B = 0, because There are no unpaired electrons in the outer level. In the excited state (ES), paired electrons 2s 2 occupy 2s 1 and 2p 1 sublevels, respectively - B = 2.

The valence capabilities of p-elements of the same group may not be the same. This is due to the unequal number of AOs in the valence level of atoms of elements located in different periods. For example, oxygen exhibits a constant valency B = 2, since its valence electrons are at energy level 2, where there are no vacant (free) AOs. Sulfur in an excited state has a maximum B=6. This is explained by the presence of vacant 3d orbitals at the third energy level.

Direction of covalent bond. Spatial structure of molecules. The strongest chemical bonds arise in the direction of maximum overlap of atomic orbitals (AO). Since AOs have a certain shape and energy, their maximum overlap is possible with the formation of hybrid orbitals. AO hybridization makes it possible to explain the spatial structure of molecules, therefore the covalent bond is characterized by directionality.

3. Hybridization of atomic orbitals and spatial structure
molecules. Atoms often form bonds with electrons of different energy states. Thus, the atoms of beryllium Be (2s12р1), boron B (2s12р2), carbon C (2s12р3) take part in the formation of bonds s- And R-electrons. Although s- And R-clouds differ in shape and energy, the chemical bonds formed with their participation turn out to be equivalent and located symmetrically. The question arises of how electrons of unequal initial state form equivalent chemical bonds. The answer to this gives insight into the hybridization of valence orbitals.

According to hybridization theories chemical bonds are formed by electrons not of “pure” ones, but of “mixed” ones, the so-called hybrid orbitals. During hybridization, the original shape and energy of the orbitals (electron clouds) change and AOs of a new, but identical shape and energy are formed. Wherein the number of hybrid orbitals is equal to the number of atomic orbitals, from which they were formed.

Rice. 5. Types of hybridization of valence orbitals.

The nature of the hybridization of the valence orbitals of the central atom and their spatial arrangement determine the geometry of the molecules. Yes, when sp hybridization In beryllium Be AOs, two sp-hybrid AOs arise, located at an angle of 180° (Fig. 5), hence the bonds formed with the participation of hybrid orbitals have a bond angle of 180°. Therefore, the BeCl 2 molecule has a linear shape. At sp 2 -hybridization boron B, three sp 2 hybrid orbitals are formed, located at an angle of 120°. As a result, the BCl 3 molecule has a trigonal shape (triangle). At sp 3 -hybridization AO carbon C, four hybrid orbitals arise, which are symmetrically oriented in space to the four vertices of the tetrahedron, therefore the CCl 4 molecule has
also tetrahedral shape. The tetrahedral shape is characteristic of many tetravalent carbon compounds. Due to sp 3 -hybridization of the orbitals of nitrogen and boron atoms, NH 4 + and BH 4 – also have a tetrahedral shape.

The fact is that the central atoms of these molecules, respectively, the C, N and O atoms, form chemical bonds due to sp 3 hybrid orbitals. The carbon atom has four unpaired electrons per four sp 3 hybrid orbitals. This determines the formation of four C-H bonds and the arrangement of hydrogen atoms at the vertices of a regular tetrahedron with a bond angle of 109°28¢. The nitrogen atom has one lone electron pair and three unpaired electrons per four sp 3 hybrid orbitals. The electron pair turns out to be nonbonding and occupies one of the four hybrid orbitals, so the H 3 N molecule has the shape of a trigonal pyramid. Due to the repulsive effect of the non-bonding electron pair, the bond angle in the NH 3 molecule is less than the tetrahedral one and amounts to 107.3°. The oxygen atom has two nonbonding electron pairs and two unpaired electrons per four sp 3 hybrid orbitals. Now two of the four hybrid orbitals are occupied by nonbonding electron pairs, so the H 2 O molecule has an angular shape. The repulsive effect of two non-bonding electron pairs is manifested to a greater extent, therefore the bond angle is distorted against the tetrahedral one even more strongly and in a water molecule is 104.5° (Fig. 6).

Rice. 6. Effect of non-bonding electron pairs
central atom on the geometry of molecules.

Thus, the BC method well explains the saturation and direction of chemical bonds, such quantitative parameters as energy ( E), length of chemical bonds ( l) and bond angles (j) between chemical bonds (structure of molecules). This is conveniently and clearly demonstrated using ball-and-stick models of atoms and molecules. The BC method also explains well the electrical properties of molecules, characterized by the electronegativity of atoms and the dipole moment of molecules. Electronegativity of atoms refers to their ability to be more positive or negative when forming a chemical bond, or in other words, the ability to attract or donate electrons, forming anions and cations. The first is quantitative
characterized by ionization potential ( E P.I), the second is the energy of electron affinity ( E S.E).

Table 3

Spatial configuration of molecules and complexes AB n

Type of hybridization of the central atom A	Number of electron pairs of atom A	Molecule type	Spatial configuration	Examples
connecting	non-binding
sp			AB 2	Linear	BeCl 2 (g), CO 2
sp 2			AB 3	Triangular	BCl 3 , CO 3 2–
			AB 2	Corner	O 3
sp 3			AB 4	Tetrahedral	CCl4, NH4, BH4
			AB 3	Trigonal-pyramidal	H3N,H3P
			AB 2	Corner	H2O
sp 3 d			AN 5	Trigonal bipyramidal	PF5, SbCl5
			AB 4	Distorted tetrahedral	SF 4
			AB 3	T-shaped	ClF 3
			AB 2	Linear	XeF 2
sp 3 d 2			AB 6	Octahedral	SF 6, SiF 6 2–
			AB 5	Square-pyramidal	IF 5

Chemical thermodynamics

1. Basic concepts and definitions.Thermodynamics – is a science that studies the general patterns of processes accompanied by the release, absorption and transformation of energy. Chemical thermodynamics studies the mutual transformations of chemical energy and its other forms - thermal, light, electrical, etc., establishes the quantitative laws of these transitions, and also makes it possible to predict the stability of substances under given conditions and their ability to enter into certain chemical reactions. Thermochemistry, which is a branch of chemical thermodynamics, studies the thermal effects of chemical reactions.

Hess's law. In chemical thermodynamics, the first law is transformed into Hess's law, which characterizes the thermal effects of chemical reactions. Heat, like work, is not a function of state. Therefore, to give the thermal effect the property of a state function, enthalpy (D H), the directional change of which is D H=D U+P D V at constant pressure. Let us note that P D V= A – expansion work, and D H = –Q(with reverse sign) . Enthalpy is characterized by the heat content of the system so that the exothermic reaction lowers D H. Please note that the release of heat in a chemical reaction ( exothermic) corresponds to D H < 0, а поглощению (endothermic) D H> 0. In the old chemical literature it was accepted opposite system of signs (!) ( Q> 0 for exothermic reactions and Q < 0 для эндотермических).

The change in enthalpy (thermal effect) does not depend on the reaction path, but is determined only by the properties of the reactants and products (Hess’s law, 1836)

Let's show this with the following example:

C(graphite) + O 2 (g) = CO 2 (g) D H 1 = –393.5 kJ

C(graphite) + 1/2 O 2 (g) = CO(g) D H 2 = –110.5 kJ

CO (g.) + 1 / 2 O 2 (g.) = CO 2 (g.) D H 3 = –283.0 kJ

Here, the enthalpy of formation of CO 2 does not depend on whether the reaction proceeds in one stage or in two, with the intermediate formation of CO (D H 1 = D H 2+D H 3). Or in other words, the sum of the enthalpies of chemical reactions in the cycle is zero:

Where i– number of reactions in a closed cycle.

In any process where the final and initial states of substances are the same, the sum of all heats of reaction is zero.

For example, we have a sequence of several chemical processes that ultimately lead to the original substance and are each characterized by its own enthalpy, i.e.

and according to Hess's law,

D H 1+D H 2+D H 3+D H 4 = 0, (7.4)

The resulting thermal effect is zero because heat is released at some stages and absorbed at others. This leads to mutual compensation.

Hess's law allows us to calculate the thermal effects of those reactions for which direct measurement is impossible. For example, consider the reaction:

H 2 (g.) + O 2 (g.) = H 2 O 2 (l.) D H 1 = ?

The following thermal effects can be easily measured experimentally:

H 2 (g.) + 1/2 O 2 (g.) = H 2 O (l.) D H 2 = –285.8 kJ,

H 2 O 2 (l.) = H 2 O (l.) + 1 / 2 O 2 (g.) D H 3 = –98.2 kJ.

Using these values, you can get:

D H 1 = D H 2 – D H 3 = –285.8 + 98.2 = –187.6 (kJ/mol).

Thus, it is sufficient to measure the thermal effects of a limited number of reactions in order to then theoretically calculate the thermal effect of any reaction. In practice tabulated standard enthalpies of formation D Hf° 298 measured at T=298.15 K (25°C) and pressure p= 101.325 kPa (1 atm), i.e. at standard conditions. (Do not confuse standard conditions with normal conditions!)

Standard enthalpy of formation D Hf° is the change in enthalpy during the reaction of the formation of 1 mole of a substance from simple substances:

Ca (solid) + C (graphite) + 3 / 2 O 2 (g) = CaCO 3 (solid) D H° 298 =–1207 kJ/mol.

Please note that the thermochemical equation indicates the aggregative states of substances. This is very important, since transitions between states of aggregation ( phase transitions) are accompanied by the release or absorption of heat:

H 2 (g.) + 1/2 O 2 (g.) = H 2 O (l.) D H° 298 = –285.8 kJ/mol,

H 2 (g.) + 1/2 O 2 (g.) = H 2 O (g.) D H° 298 = –241.8 kJ/mol.

H 2 O (g.) = H 2 O (l.) D H° 298 = –44.0 kJ/mol.

The standard enthalpies of formation of simple substances are assumed to be zero. If a simple substance can exist in the form of several allotropic modifications, then D H° = 0 is assigned to the most stable form under standard conditions, for example, oxygen, and not ozone, graphite, and not diamond:

3 / 2 O 2 (g.) = O 3 (g.) D H° 298 = 142 kJ/mol,

C (graphite) = C (diamond) D H° 298 = 1.90 kJ/mol.

A consequence of Hess’s law, taking into account the above, is that the change in enthalpy during the reaction will be equal to the sum of the enthalpies of formation of the products minus the sum of the enthalpies of formation of the reactants, taking into account the stoichiometric coefficients of the reaction:

Related information.

§ 1 M.V. Lomonosov as the founder of atomic-molecular science

Since the 17th century, science has had molecular teaching, which has been used to explain physical phenomena. The practical application of molecular theory in chemistry was limited by the fact that its provisions could not explain the essence of the occurrence of chemical reactions or answer the question of how new substances are formed from some substances during a chemical process.

The solution to this issue turned out to be possible on the basis of atomic-molecular theory. In 1741, in the book “Elements of Mathematical Chemistry,” Mikhail Vasilyevich Lomonosov actually formulated the foundations of atomic-molecular science. The Russian scientist-encyclopedist considered the structure of matter not as a specific combination of atoms, but as a combination of larger particles - corpuscles, which, in turn, consist of smaller particles - elements.

Lomonosov's terminology underwent changes over time: what he called corpuscles began to be called molecules, and the term element was replaced by the term atom. However, the essence of the ideas and definitions he expressed brilliantly stood the test of time.

§ 2 History of the development of atomic-molecular science

The history of the development and establishment of atomic-molecular science in science turned out to be very difficult. Working with objects of the microworld caused enormous difficulties: atoms and molecules were impossible to see and, thus, verify their existence, and attempts to measure atomic masses often ended in obtaining erroneous results. 67 years after Lomonosov's discovery, in 1808, the famous English scientist John Dalton put forward the atomic hypothesis. According to it, atoms are the smallest particles of matter that cannot be divided into their component parts or converted into each other. According to Dalton, all atoms of one element have exactly the same weight and are different from the atoms of other elements. By combining the theory of atoms with the theory of chemical elements developed by Robert Boyle and Mikhail Vasilyevich Lomonosov, Dalton provided a solid foundation for further theoretical research in chemistry. Unfortunately, Dalton denied the existence of molecules in simple substances. He believed that only complex substances consist of molecules. This did not contribute to the further development and application of atomic-molecular teaching.

The conditions for the dissemination of the ideas of atomic-molecular science in natural sciences developed only in the second half of the 19th century. In 1860, at the International Congress of Natural Scientists in the German city of Karlsruhe, scientific definitions of the atom and molecule were adopted. There was no study of the structure of substances at that time, so it was accepted that all substances consist of molecules. It was believed that simple substances, such as metals, consist of monatomic molecules. Subsequently, such a complete extension of the principle of molecular structure to all substances turned out to be erroneous.

§ 3 Basic provisions of atomic-molecular teaching

1. A molecule is the smallest part of a substance that retains its composition and most important properties.

2. Molecules are made up of atoms. Atoms of one element are similar to each other, but different from atoms of other chemical elements.

Atomic-molecular science- a set of provisions, axioms and laws that describe all substances as a set of molecules consisting of atoms.

Ancient Greek philosophers Long before the beginning of our era, they already put forward the theory of the existence of atoms in their works. Rejecting the existence of gods and otherworldly forces, they tried to explain all incomprehensible and mysterious natural phenomena by natural causes - the connection and separation, interaction and mixing of particles invisible to the human eye - atoms. But for many centuries, church ministers persecuted adherents and followers of the doctrine of atoms and subjected them to persecution. But due to the lack of necessary technical devices, ancient philosophers could not scrupulously study natural phenomena, and under the concept of “atom” they hid the modern concept of “molecule”.

Only in the middle of the 18th century the great Russian scientist M.V. Lomonosov substantiated atomic-molecular concepts in chemistry. The main provisions of his teaching are set out in the work “Elements of Mathematical Chemistry” (1741) and a number of others. Lomonosov named the theory corpuscular-kinetic theory.

M.V. Lomonosov clearly distinguished between two stages in the structure of matter: elements (in the modern sense - atoms) and corpuscles (molecules). The basis of his corpuscular-kinetic theory (modern atomic-molecular teaching) is the principle of discontinuity of the structure (discreteness) of matter: any substance consists of individual particles.

In 1745 M.V. Lomonosov wrote:“An element is a part of a body that does not consist of any smaller and different bodies... Corpuscles are a collection of elements into one small mass. They are homogeneous if they consist of the same number of the same elements connected in the same way. Corpuscles are heterogeneous when their elements are different and connected in different ways or in different numbers; the infinite variety of bodies depends on this.

Molecule is the smallest particle of a substance that has all its chemical properties. Substances having molecular structure, consist of molecules (most non-metals, organic substances). A significant part of inorganic substances consists of atoms(atomic crystal lattice) or ions (ionic structure). Such substances include oxides, sulfides, various salts, diamond, metals, graphite, etc. The carrier of chemical properties in these substances is a combination of elementary particles (ions or atoms), that is, a crystal is a giant molecule.

Molecules are made up of atoms. Atom- the smallest, further chemically indivisible component of the molecule.

It turns out that molecular theory explains the physical phenomena that occur with substances. The study of atoms comes to the aid of molecular theory in explaining chemical phenomena. Both of these theories - molecular and atomic - are combined into the atomic-molecular theory. The essence of this doctrine can be formulated in the form of several laws and regulations:

1. substances are made up of atoms;
2. when atoms interact, simple and complex molecules are formed;
3. during physical phenomena, molecules are preserved, their composition does not change; with chemicals - they are destroyed, their composition changes;
4. molecules of substances consist of atoms; in chemical reactions, atoms, unlike molecules, are preserved;
5. the atoms of one element are similar to each other, but different from the atoms of any other element;
6. chemical reactions involve the formation of new substances from the same atoms that made up the original substances.

Thanks to its atomic-molecular theory M.V. Lomonosov is rightfully considered the founder of scientific chemistry.

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