Russian scientist formulated atomic. Atomic-molecular science

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The leading idea of atomic-molecular science, which forms the foundation of modern physics, chemistry and natural science, is the idea of discreteness (discontinuity of structure) of matter.

The first ideas that matter consists of individual indivisible particles, appeared in ancient times and were initially developed in line with general philosophical ideas about the world. For example, some schools of thought Ancient India(1st millennium BC) recognized not only the existence of primary indivisible particles of matter (anu), but also their ability to combine with each other, forming new particles. Similar teachings existed in other countries of the ancient world. The greatest fame and influence on the subsequent development of science was exerted by ancient Greek atomism, the creators of which were Leucippus (5th century BC) and Democritus (b. c. 460 BC - d. c. 370 BC. ). “The causes of all things,” wrote the ancient Greek philosopher and scientist Aristotle (384–322 BC), expounding the Democritus’ doctrine, “are certain differences in atoms. And there are three differences: form, order and position.” In the works of Aristotle himself there is an important concept of mixis - a homogeneous compound formed from various substances. Later, the ancient Greek materialist philosopher Epicurus (342–341 BC - 271–270 BC) introduced the concept of the mass of atoms and their ability to spontaneously deflect during movement.

It is important to note that, according to many ancient Greek scientists, a complex body is not a simple mixture of atoms, but a qualitatively new integral formation, endowed with new properties. However, the Greeks had not yet developed the concept of special “polyatomic” particles - molecules, intermediate between atoms and complex bodies, which would be the smallest carriers of the properties of bodies.

The Middle Ages saw a sharp decline in interest in ancient atomism. The Church accused ancient Greek philosophies of asserting that the world arose from random combinations of atoms, and not by the will of God, as required by Christian dogma.

In the XVI–XVII centuries. In an atmosphere of general cultural and scientific upsurge, the revival of atomism begins. During this period, advanced scientists different countries: G. Galileo (1564–1642) in Italy, P. Gassendi (1592–1655) in France, R. Boyle (1627–1691) in England and others - proclaimed the principle: not to seek truth in the Holy Scriptures, but “directly” read a book of nature

P. Gassendi and R. Boyle owe the main credit for the further development of ancient atomism. Gassendi introduced the concept of a molecule, by which he understood a qualitatively new formation, composed by combining several atoms. A broad program for the creation of a corpuscular philosophy of nature was proposed by R. Boyle. The world of corpuscles, their movement and “plexus”, according to the English scientist, is very complex. The world as a whole and its smallest particles are purposefully arranged mechanisms. Boyle's corpuscles are no longer the primary, unbreakable atoms of ancient philosophers, but a complex whole capable of changing its structure through movement.

“Ever since I read Boyle,” wrote M.V. Lomonosov, “I have been possessed by a passionate desire to explore the smallest particles.” The great Russian scientist M.V. Lomonosov (1711–1765) developed and substantiated the doctrine of material atoms and corpuscles. He attributed to atoms not only indivisibility, but also an active principle - the ability to move and interact. “Insensitive particles must differ in mass, shape, motion, inertial force or location.” Corpuscles of homogeneous bodies, according to Lomonosov, “consist of the same number of the same elements, connected in the same way... Corpuscles are heterogeneous when their elements are different or connected in various ways or in different numbers." Only because the study of mass relations at the beginning of the 18th century. just beginning, Lomonosov was unable to create a quantitative atomic-molecular theory.

This was done by the English scientist D. Dalton (1766–1844). He considered an atom as the smallest particle of a chemical element, differing from atoms of other elements primarily in mass. A chemical compound, according to his teaching, is a collection of “complex” (or “composite”) atoms containing certain numbers of atoms of each element, characteristic only for a given complex substance. The English scientist compiled the first table of atomic masses, but due to the fact that his ideas about the composition of molecules were often based on arbitrary assumptions based on the principle of “greatest simplicity” (for example, for water he accepted the formula OH), this table turned out to be inaccurate.

In addition, in the first half of the 19th century. many chemists did not believe in the possibility of determining true atomic masses and preferred to use equivalents that could be found experimentally. Therefore, different formulas were assigned to the same compound, and this led to the establishment of incorrect atomic and molecular masses.

One of the first who began the struggle for the reform of theoretical chemistry were the French scientists C. Gerard (1816–1856) and O. Laurent (1807–1853), who created the right system atomic masses and chemical formulas. In 1856, the Russian scientist D. I. Mendeleev (1834–1907), and then, independently of him, the Italian chemist S. Cannizzaro (1826 - 1910) proposed a method for calculating the molecular weight of compounds from the double density of their vapors relative to hydrogen. By 1860, this method was established in chemistry, which was crucial for the establishment of the atomic-molecular theory. In his speech at the International Congress of Chemists in Karlsruhe (1860), Cannizzaro convincingly proved the correctness of the ideas of Avogadro, Gerard and Laurent, the need for their adoption for the correct determination of atomic and molecular masses and the composition of chemical compounds. Thanks to the work of Laurent and Cannizzaro, chemists realized the difference between the form in which an element exists and reacts (for example, for hydrogen, it is H 2), and the form in which it is present in a compound (HCl, H 2 O, NH 3 and etc.). As a result, Congress adopted the following definitions of an atom and a molecule: molecule - “a quantity of a body that enters into reactions and determines chemical properties”; atom - “the smallest amount of an element included in particles (molecules) of compounds.” It was also accepted that the concept of “equivalent” should be considered empirical, not coinciding with the concepts of “atom” and “molecule”.

The atomic masses established by S. Cannizzaro served as the basis for D. I. Mendeleev in the discovery of the periodic law of chemical elements. The decisions of the congress had a beneficial effect on the development of organic chemistry, because the establishment of formulas of compounds opened the way for the creation of structural chemistry.

Thus, by the beginning of the 1860s. The atomic-molecular doctrine was formed in the form of the following provisions.

1. Substances consist of molecules. A molecule is the smallest particle of a substance that has its chemical properties. Many physical properties of a substance - boiling and melting points, mechanical strength, hardness, etc. - are determined by the behavior large number molecules and the action of intermolecular forces.

2. Molecules consist of atoms that are connected to each other in certain relationships (see Molecule; Chemical bond; Stoichiometry).

3. Atoms and molecules are in constant spontaneous motion.

4. Molecules simple substances consist of identical atoms (O 2, O 3, P 4, N 2, etc.); molecules of complex substances - from different atoms (H 2 O, HCl).

6. The properties of molecules depend not only on their composition, but also on the way in which the atoms are connected to each other (see Theory of chemical structure; Isomerism).

Modern science has developed the classical atomic-molecular theory, and some of its provisions have been revised.

It was established that the atom is not an indivisible structureless formation. However, many scientists in the last century also guessed about this.

It turned out that not in all cases the particles that form a substance are molecules. Many chemical compounds, especially in solid and liquid states, have ionic structures, such as salts. Some substances, such as noble gases, consist of individual atoms that weakly interact with each other even in liquid and solid states. In addition, a substance may consist of particles formed by the combination (association) of several molecules. Yes, chemically pure water formed not only by individual H 2 O molecules, but also by polymer molecules (H 2 O)n, where n = 2–16; At the same time, it contains hydrated H + and OH − ions. Special group compounds form colloidal solutions. And finally, when heated to temperatures of the order of thousands and millions of degrees, the substance transforms into special condition - plasma, which is a mixture of atoms, positive ions, electrons and atomic nuclei.

It turned out that the quantitative composition of molecules with the same qualitative composition can sometimes vary within wide limits (for example, nitrogen oxide can have the formula N 2 O, NO, N 2 O 3, NO 2, N 2 O 4, N 2 O 5, NO 3 ), while if we consider not only neutral molecules, but also molecular ions, then the boundaries possible compositions are expanding. Thus, the NO 4 molecule is unknown, but the NO 3− 4 ion was recently discovered; there is no CH 5 molecule, but the CH + 5 cation is known, etc.

So-called compounds of variable composition were discovered, in which per unit mass of a given element there is a different mass of another element, for example: Fe 0.89–0.95 O, TiO 0.7–1.3, etc.

The position that molecules consist of atoms was clarified. According to modern quantum mechanical concepts (see Quantum chemistry), only the core of atoms in a molecule remains more or less unchanged, i.e., the core and internal electron shells, while the nature of the movement of external (valence) electrons radically changes so that a new, molecular electron shell is formed, covering the entire molecule (see Chemical bond). In this sense, there are no unchanging atoms in molecules.

Taking into account these clarifications and additions, it should be borne in mind that modern science preserved the rational grain of the classical atomic-molecular teaching: ideas about the discrete structure of matter, the ability of atoms to give, by combining with each other in a certain order, qualitatively new and more complex formations, and the continuous movement of the particles that make up matter.

Lecture 1

SUBJECT AND IMPORTANCE OF CHEMISTRY

1. Chemistry subject. Among natural sciences, defining 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 to study chemical processes(reactions) of the formation and destruction of molecules (chemical bonds), as well as relationships and transitions between these processes and other forms of matter movements (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– 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, to put it differently, the crystal lattices of solids have defects (vacancies and inclusions of sites).

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 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. Molar mass 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. 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 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 connections important group form 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 global catastrophe, all the scientific knowledge accumulated by mankind were destroyed and only one phrase were passed down to future generations, then which statement, composed of the fewest words, would bring 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 passing electric current Through the rarefied gas enclosed in the tube, a stream of weak light emanates from the negative pole (cathode) - the cathode ray. 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 deviated towards the negative plate of the capacitor, b-rays (electron flow) are strongly deviated towards the positive plate of the capacitor, g-rays ( electromagnetic waves) are not deflected by either an electric or magnetic field.

And finally, the opening 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 high-temperature flame or other means) through a special optical device (prism, system of prisms or diffraction gratings), which decomposes the complex radiation into monochromatic components with a certain wavelength (l) and, accordingly, with a certain oscillation frequency 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, wave-particle 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 manifest themselves 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 Periodic table 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. His physical meaning as energy level numbers - 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.

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:

substances are made up of atoms;
when atoms interact, simple and complex molecules are formed;
during physical phenomena, molecules are preserved, their composition does not change; with chemicals - they are destroyed, their composition changes;
molecules of substances consist of atoms; in chemical reactions, atoms, unlike molecules, are preserved;
the atoms of one element are similar to each other, but different from the atoms of any other element;
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.

website, when copying material in full or in part, a link to the source is required.

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

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

substances are made up of atoms;
when atoms interact, simple and complex molecules are formed;
during physical phenomena, molecules are preserved, their composition does not change; with chemicals - they are destroyed, their composition changes;
molecules of substances consist of atoms; in chemical reactions, atoms, unlike molecules, are preserved;
the atoms of one element are similar to each other, but different from the atoms of any other element;
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.

blog.site, when copying material in full or in part, a link to the original source is required.

BASIC CONCEPTS AND LAWS OF CHEMISTRY

Substances and their properties. Chemistry subject

Let's look around. We ourselves and everything that surrounds us consists of substances. There are a lot of substances. Currently, scientists know about 10 million organic and about 100 thousand inorganic substances. And they are all characterized by certain properties. Properties of a substance are the characteristics by which substances differ from each other or are similar to each other..

Each individual type of matter, which under given conditions has certain physical properties, e.g. aluminum, sulfur, water, oxygen, called substance.

Chemistry studies the composition, structure, properties and transformation of substances. A deep knowledge of chemistry is absolutely necessary for specialists in all sectors of the national economy. Along with physics and mathematics, it forms the basis for the training of highly qualified specialists.

Various changes occur with substances, for example: evaporation of water, melting of glass, combustion of fuel, rusting of metals, etc. These changes with substances can be attributed to physical or to chemical phenomena.

Physical phenomena are those phenomena in which these substances do not transform into others, but usually only their state of aggregation or form changes

Chemical phenomena are those phenomena that result in the formation of other substances from given substances. Chemical phenomena are called chemical transformations or chemical reactions

In chemical reactions, starting substances are transformed into other substances that have different properties. This can be judged by external signs of chemical reactions: 1) release of heat (sometimes light); 2) color change; 3) the appearance of odor; 4) formation of sediment; 5) gas release.

Atomic-molecular science

In the XVIII – XIX centuries. As a result of the work of M.V. Lomonosov, Dalton, Avogadro and others, a hypothesis was put forward about the atomic-molecular structure of matter. This hypothesis is based on the idea of the real existence of atoms and molecules. In 1860, the International Congress of Chemists clearly defined the concepts atom and molecule. All scientists accepted the atomic-molecular doctrine. Chemical reactions began to be considered from the point of view of atomic-molecular theory. At the end of the 19th and beginning of the 20th centuries. Atomic-molecular teaching turned into a scientific theory. At this time, scientists proved experimentally that atoms and molecules exist objectively, independently of humans.

Currently, it is possible not only to calculate the sizes of individual molecules and their mass, but also to determine the order of connection of atoms in a molecule. Scientists determine the distance between molecules and even photograph some macromolecules. It is also now known that not all substances are made of molecules.

Basic provisions of atomic-molecular teaching can be formulated like this:

1. There are substances with molecular and non-molecular structure.

2. A molecule is the smallest particle of a substance that retains its chemical properties.

3. There are gaps between the molecules, the sizes of which depend on the state of aggregation and temperature. The greatest distances exist between gas molecules. This explains their easy compressibility. Liquids where the spaces between the molecules are much smaller are more difficult to compress. In solids, the spaces between molecules are even smaller, so they hardly compress.

4. Molecules are in continuous motion. The speed of movement of molecules depends on temperature. As temperature increases, the speed of molecular movement increases.

5. Between molecules there are forces of mutual attraction and repulsion. These forces are expressed to the greatest extent in solids, and to the least in gases.

6. Molecules are made up of atoms, which, like molecules, are in continuous motion.

7 Atoms are the smallest chemically indivisible particles.

8. Atoms of one type differ from atoms of another type in mass and properties. Each individual type of atom is called a chemical element.

9. During physical phenomena, molecules are preserved; during chemical phenomena, as a rule, they are destroyed. In chemical reactions, a rearrangement of atoms occurs.

Atomic-molecular theory is one of the main theories of natural sciences. This theory confirms the material unity of the world.

According to modern concepts, substances in gaseous and vaporous states are made up of molecules. In the solid (crystalline) state, only substances with a molecular structure consist of molecules, for example, organic substances, nonmetals (with a few exceptions), carbon monoxide (IV), and water. Most solid (crystalline) inorganic substances do not have a molecular structure. They do not consist of molecules, but of other particles (ions, atoms) and exist in the form of macrobodies. For example, many salts, oxides and sulfides of metals, diamond, silicon, metals.

In substances with a molecular structure, the chemical bonds between molecules are less strong than between atoms. Therefore, they have relatively low melting and boiling points. In substances with a non-molecular structure, the chemical bond between particles is very strong. Therefore, they have high melting and boiling points. Modern chemistry studies the properties of microparticles (atoms, molecules, ions, etc.) and macrobodies.

Molecules and crystals are made up of atoms. Each individual type of atom is called a chemical element.

In total, the existence of (92) different chemical elements has been established in nature (on Earth). Another 22 elements were obtained artificially using nuclear reactors and powerful accelerators.

All substances are divided into simple and complex.

Substances that consist of atoms of one element are called simple.

Sulfur S, hydrogen H2, oxygen O2, ozone O3, phosphorus P, iron Fe are simple substances.

Substances that consist of atoms of different elements are called complex.

For example, water H 2 O consists of atoms of different elements - hydrogen H and oxygen O; chalk CaCO 3 consists of atoms of the elements calcium Ca, carbon C and oxygen O . Water and chalk are complex substances.

The concept of “simple substance” cannot be identified with the concept of “chemical element”. A simple substance is characterized by a certain density, solubility, boiling and melting points, etc. A chemical element is characterized by a certain positive nuclear charge (ordinal number), oxidation state, 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.

The names of the elements coincide with the names of their corresponding simple substances, with the exception of carbon.

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

The phenomenon of allotropy is caused by two reasons: 1) a different number of atoms in the molecule, for example oxygen O 2 and ozone O 3; 2) the formation of various crystalline forms, such as diamond, graphite and carbine.

2. Stoichiometric laws

Stoichiometry- a branch of chemistry that deals with mass and volume relationships between reacting substances. Translated from Greek, the word “stoichiometry” means “component” and “I measure.”

The basis of stoichiometry is stoichiometric laws: conservation of mass of substances, constancy of composition, Avogadro’s law, law of volumetric ratios of gases, law of equivalents. They confirmed the atomic-molecular theory. In turn, atomic-molecular theory explains stoichiometric laws.

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