Physical chemistry basic formulas and concepts. Physical chemistry

Ministry of Education of the Russian Federation Tomsk Polytechnic University ________________________________________________________________________________ N. A. Kolpakova, V. A. Kolpakov, S. V. Romanenko PHYSICAL CHEMISTRY Textbook Part I Tomsk 2004 UDC 541.1 Physical chemistry. Textbook / N.A. Kolpakova, V.A. Kolpakov, S.V. Romanenko. - Tomsk: Ed. TPU, 2004. - Part 1. - 168 p. The textbook covers the following sections of "Physical Chemistry": the basic laws of thermodynamics, chemical and phase equilibrium, thermodynamics of non-electrolyte solutions. The manual was prepared at the Department of Physical and Analytical Chemistry of TPU and is intended for students of correspondence courses in chemical specialties. Published by order of the Editorial and Publishing Council of Tomsk Polytechnic University Reviewers: Kurina L.N. – Prof. Department of Physical Chemistry, TSU, Doctor of Chem. sciences; Buinovsky A.S. - Head. cafe Chemistry TPU STU, doctor of chem. Sciences. © Tomsk Polytechnic University, 2004 © Authors, 2004 CHAPTER 1 . INTRODUCTION TO PHYSICAL CHEMISTRY 1.1. BRIEF HISTORICAL OUTLINE OF THE DEVELOPMENT OF PHYSICAL CHEMISTRY The name and definition of the content of physical chemistry was first given by M.V. Lomonosov (1752): “Physical chemistry is a science that, on the basis of the positions and experiments of physical scientists, must explain the reason for what happens through chemical operations in complex bodies” . The teaching of physical chemistry in Russia as an independent science was introduced by prof. N. N. Beketov in 1860 at Kharkov University. Lomonosov's most important theoretical and experimental studies led him to discoveries that have not lost their significance even now. Lomonosov came close to the correct definition of the principle of conservation of matter and motion, the kinetic nature of heat, and also noted the impossibility of a spontaneous transfer of heat from a colder body to a warmer one, which is currently one of the formulations of the second law of thermodynamics. Over the next century, research was carried out, on the basis of which many important discoveries and generalizations were made. K. V. Scheele in Sweden (1773) and Fontana in France (1777) discovered the adsorption of gases; T. E. Lovits in Russia (1785) discovered adsorption from solutions. A. L. Lavoisier and P. S. Laplace in France (1779–1784) studied the heat capacities of substances and the heat effects of reactions. At the beginning of the XIX century. G. Davy in England and L. J. Tenard in France discovered catalytic reactions, and J. J. Berzelius in Sweden (1835) further developed the idea of ​​catalysis. The foundations of electrochemistry were laid by research on galvanic cells, electrolysis, and current transfer in electrolytes. Galvani and A. Volta in Italy created in 1799 a galvanic cell. VV Petrov in Russia (1802) discovered the phenomenon of an electric arc. T. Grotgus in Russia in 1805 laid the foundations for the theory of electrolysis. In 1800, G. Davy advanced the electrochemical theory of the interaction of substances: he widely used electrolysis for chemical research. M. Faraday, a student of Davy, in 1833-1834 formulated the quantitative laws of electrolysis. B. S. Jacobi in Russia, solving the problems of the practical use of the electrolysis process, discovered in 1836 galvanoplasty. In the first half of the XIX century. thanks to the works of D. Dalton in England (1801–1803), J. L. Gay-Lussac in France (1802) and A. Avogadro in Italy (1811), who discovered the most important laws of the gaseous state, atomistic ideas were widely developed. The works of G. I. Hess (1802–1856) on thermochemistry belong to the same period. C. Guldberg and P. Waage in Norway (1864–1867), J. W. Gibbs in the USA (1873–1878) developed the thermodynamic doctrine of chemical equilibrium, and A. L. Le Chatelier in France (1884) discovered the general principle of displacement equilibrium under changing external conditions. In the works of the Dutch chemist J. H. van't Hoff, the thermodynamic theory of chemical equilibrium was developed. He also developed the quantitative theory of dilute solutions (1885–1889). The transfer of electricity in solutions was studied in Germany by I. V. Gittorf and F. V. G. Kohlrausch. The Swedish scientist S. A. Arrhenius developed in 1883–1887. theory of electrolytic dissociation. A. M. Butlerov, who created the theory of the structure of organic compounds, left a deep mark on the development of physical chemistry. The great Russian chemist D. I. Mendeleev (1834–1907) discovered the existence of a critical temperature (1860), derived the general equation of state for gases (1874) and developed the chemical theory of solutions (1887). D. P. Konovalov (1889), a student of Mendeleev, is one of the founders of the theory of solutions. At the end of the XIX century. a number of major discoveries were made in the field of the doctrine of the structure of matter, which proved the complexity of the structure of the atom and played a huge role in the development of physical chemistry. These include the discoveries of the electron by J. B. Perrin (1895) and J. Thomson (1897), the quantum nature of light by R. Planck (1900), the existence of light pressure by P. N. Lebedev (1899), the study (since 1898 of ) phenomena of radioactivity P. Curie and M. Sklodowska-Curie. By the beginning of the XX century. physical chemistry was defined as the science that studies the structure of matter, chemical thermodynamics, including thermochemistry and the theory of equilibrium, solutions, chemical kinetics and electrochemistry. New theoretical methods were applied, and studies of the structure of atoms, molecules, and crystals came to the fore. The doctrine of the structure of matter, especially the structure of atoms and molecules, developed most rapidly in the 20th century. A major achievement in this area was the nuclear theory of the atom, proposed by E. Rutherford (1911) and developed in the first quantitative theory of the hydrogen atom, developed by the Danish physicist N. Bohr (1913). The study of the nature of the chemical bond and the structure of molecules developed in parallel with the study of the structure of the atom. By the early 1920s, W. Kossel and G. N. Lewis had developed the fundamentals of the electronic theory of chemical bonding. VG Geitler and F. London (1927) developed the quantum-mechanical theory of chemical bonding. Based on the largest discoveries of physics in the field of atomic structure and using the theoretical methods of quantum mechanics and statistical physics, as well as new experimental methods, such as X-ray analysis, spectroscopy, mass spectroscopy, magnetic methods, the method of labeled atoms and others , physicists and physical chemists have made great strides in studying the structure of molecules and crystals and in understanding the nature of the chemical bond. The theory of the rates of chemical reactions, i.e., chemical kinetics, has been greatly developed, and is now associated specifically with studies of the structure of molecules and the strength of bonds between atoms in a molecule. New branches of physical chemistry have arisen and are successfully developing: magnetochemistry, radiation chemistry, physical chemistry of high polymers, physical chemistry of silicates, gas electrochemistry, etc. Like other sciences, physical chemistry and its individual branches arose or began to develop especially successfully in periods when one or another practical need necessitated the rapid development of some branch of industry, and this development required a solid theoretical foundation. Here it is necessary to note the major studies of N. S. Kurnakov on physicochemical analysis, the work in the field of electrochemistry by A. N. Frumkin, the creation of the theory of chain reactions by N. N. Semenov, and the development of the theory of heterogeneous catalysis by A. A. Balandin. Physical chemistry plays a leading role in solving numerous problems facing chemical science and practice. At present, physical chemistry is an independent discipline with its own research methods and is the theoretical basis for applied chemical engineering disciplines. 1.2. SUBJECT AND OBJECTIVES OF PHYSICAL CHEMISTRY Physical chemistry is the science of regularities of chemical processes and physical phenomena. The main task of physical chemistry is the study and explanation of the main regularities that determine the direction of chemical processes, their speed, the influence of the medium, impurities, radiation, and the conditions for obtaining the maximum yield of a useful product. The study of physical chemistry makes it possible to understand the laws of chemistry, as well as to predict and control chemical phenomena. Modern physical chemistry makes it possible to solve the problems of efficient production control, intensification and automation of production processes. It serves as the theoretical foundation of chemical technology. Such important production processes in chemical technology as the synthesis and oxidation of ammonia, the contact production of sulfuric acid, the production of ethanol from natural gas, oil cracking, and many others are based on the results of physicochemical studies of the reactions underlying these processes. 5 processes. Without physical chemistry, it is impossible to solve the problem of creating substances with desired properties, develop new current sources, and many other issues of efficient production. Therefore, knowledge of physical chemistry for future process engineers opens up great opportunities for solving various problems encountered in the practical activities of an engineer at factories and research institutes. The name of the science - "physical chemistry" - reflects both the history of its emergence at the junction of two sciences - physics and chemistry, as well as the fact that it widely uses the theoretical laws and experimental methods of physics in the study of chemical phenomena. 1.3. CLASSIFICATION OF METHODS OF PHYSICAL CHEMISTRY Several theoretical methods are used in physical chemistry.  The quantum chemical method uses the properties of elementary particles to describe chemical transformations. Using the laws of quantum mechanics, the properties and reactivity of molecules are described, as well as the nature of the chemical bond based on the properties of the elementary particles that make up the molecules.  The thermodynamic (phenomenological) method is based on several laws (postulates), which are a generalization of experimental data. It makes it possible, on their basis, to find out the energy properties of the system, to predict the course of the chemical process and its result by the moment of equilibrium.  The quantum-statistical method explains the properties of substances on the basis of the properties of the molecules that make up these substances.  The kinetic method allows you to establish the mechanism and create a theory of chemical processes by studying the change in the rate of chemical reactions from various factors. Physical chemistry is characterized by the widespread use of mathematics, which not only makes it possible to most accurately express theoretical laws, but is also a necessary tool for establishing them. 6 CHAPTER 2 . BASIC LAWS OF THERMODYNAMICS The word "thermodynamics" comes from the Greek therme - heat and dynamis - force. Thermodynamics is the science of the transformation of various types of energy from one into another. Chemical thermodynamics studies the transformation of various types of energy occurring during the course of chemical reactions. 2.1. BASIC CONCEPTS OF CHEMICAL THERMODYNAMICS A system is a separate body or a group of bodies interacting and separated from the environment by a real or imaginary shell (boundary). An open system is a system that exchanges substances (mass) and energy (for example, heat) with the external environment. An isolated system (or closed system) is a system that does not exchange heat and work with the environment. The energy and volume of an isolated system are constant in time. An example of such a system is, for example, a thermos. If the boundary does not pass heat, then the process occurring in the system is called adiabatic. When a system exchanges heat and work with the environment, changes occur both in the system and in the environment. Thermodynamic systems can be homogeneous or heterogeneous. If there are no interfaces inside the system separating parts of the system with different composition or structure, then this system is called homogeneous. Accordingly, a system consisting of various parts differing in structure or chemical composition is called heterogeneous. These parts are called phases. Thus, a phase is a part of a heterogeneous system limited by the interface and characterized by the same physical and chemical properties at all points. Each system consists of one or more substances. Individual chemicals that can be isolated from the system and exist outside of it on their own as a separate phase are called constituent substances of the system. For example, in a glass there is water in which a platinum plate is lowered. Above the glass is a mixture of gases: oxygen, hydrogen and nitrogen. This system is three-phase, it contains five constituent substances. 7 The thermodynamic state of a system is a set of values ​​of independent variables (system parameters) that determine its properties. Any property of a system can be called a thermodynamic state parameter if it is considered as one of the independent variables that determine the state of the system. Thermodynamics considers matter as a continuous medium and uses for research such thermodynamic parameters that are the result of the action of a large number of particles (macroparameters). For example, the macroparameters of a chemical reaction that proceeds even under “normal conditions” are temperature, pressure, volume, concentration, strength of gravitational, magnetic, electric and electromagnetic fields, etc. “Normal conditions” is a temperature of 20– 25 °C, atmospheric pressure, i.e. about 101 kPa, acceleration of gravity - on average about 9.8 m/s2, magnetic field strength - on average about 40 A/m, electric field strength - on average about 130 V/m, visible light illumination - about 500 lux on average. To characterize the thermodynamic state of a system, it is necessary to know not all properties, but only the smallest number of them, the so-called independent parameters of the system. As a rule, when describing a chemical process occurring on the Earth, we do not indicate the characteristics of the field, since they are constant and therefore do not affect the composition and yield of the reaction products. If the chemical process is carried out under conditions of strong magnetic or electric fields, or under intense irradiation with ultraviolet, X-rays, or even visible light, then the field parameters will have a significant effect on the composition and yield of the reaction products. In this case, the field parameters must be specified. Thermodynamic parameters are divided into extensive and intensive. Quantities proportional to the mass (or amount of substance) of the considered working fluid or thermodynamic system are called extensive, they are volume, internal energy, enthalpy, etc. Intensive quantities do not depend on the mass of the thermodynamic system. These are, for example, temperature and pressure. Pressure is a physical quantity equal to the ratio of a force uniformly distributed over the surface of a body to the surface area located perpendicular to the force: p \u003d S The unit of pressure in SI - pascal (Pa) is the pressure caused by a force of 1 N, uniformly distributed on a surface of 1 m2 located perpendicular to the direction of force: 1 N/m2 = 1 Pa. In practice, multiple and sub-multiple units of pressure are used: kilopascal 8 (103 Pa = 1 kPa); megapascal (106 Pa = 1 MPa); hectapascal (102 Pa = 1 hPa), as well as an off-system unit - bar (1 bar = 105 Pa). According to the conclusions of the molecular-kinetic theory, the pressure of a gas is the result of impacts of randomly continuously moving molecules against the vessel wall. The simplest relationships between the parameters and the behavior of molecules were obtained for an ideal gas. An ideal gas is understood as a gas consisting of elastic molecules, between which there are no interaction forces, which have a negligibly small intrinsic volume compared to the volume occupied by the gas. Any real gas at a relatively low pressure (close to atmospheric) behaves practically like an ideal one (strictly at p → 0). The equation of state of an ideal gas - the Mendeleev - Clapeyron equation has the form: pV = nRT, where p is the gas pressure, Pa; V - volume, m3; n is the amount of gas, mol; R is the universal gas constant equal to 8.314 J/(mol K); T is the absolute temperature, K. The temperature characterizes the thermal state of the system. Experimentally, the concepts of a warmer and colder body can be established, but the temperature cannot be measured directly. It is determined from the numerical values ​​of other physical parameters that depend on temperature, which is the basis for constructing empirical temperature scales. Various physical quantities can serve as such parameters (thermometric parameters). Among them are the volume of a body at constant pressure, pressure at a constant volume, electrical conductivity, thermoelectromotive force, geometric parameters of bodies, brightness of the glow, etc. A device for measuring temperature is called a thermometer. To build any empirical temperature scale, three assumptions are used: 1) the size of a degree is set by choosing the numerical value of ∆T between two reference temperature points - temperature standards; 2) the position of the temperature zero in empirical scales is arbitrary; 3) it is assumed that the thermometric function is linear in a given temperature range. The phase transitions of pure substances are used as reference points. For example, for the empirical Celsius scale, the melting and boiling points of water at atmospheric pressure (0 and 100 degrees, respectively) are taken as reference points. The interval between these temperatures is divided into one hundred equal parts (degrees Celsius - °C). Although an objective temperature scale can be constructed using any theoretically defined thermometric function, thermodynamics uses the ideal gas equation of state as such a function. The gas thermometer makes it possible to carry out the most accurate (close to the absolute temperature scale - the Kelvin scale) temperature measurements. However, determining the temperature on the scale of a gas thermometer is a rather difficult job, which is carried out only to establish the absolute temperatures of a few reference points of phase transitions taken as reference ones. Intermediate temperatures are usually determined by empirical thermometric methods. The International Practical Temperature Scale (IPTS), adopted in 1954, is the most accurate approximation to the absolute temperature scale at the present stage. In contrast to empirical scales, the MPSH uses one experimental reference temperature point. The temperature of the triple point of water (when ice, water and water vapor are in equilibrium at the same time) was used as such a point. The temperature of the triple point of water is taken in the IPTS as 273.16 K (exactly). At atmospheric pressure, ice melts 0.01° lower. The reference point on the Celsius scale - 0 °C - corresponds to 273.15 K. The numerical value of temperatures for all other reference points (except for the triple point of water) is continuously refined as the accuracy of working with a gas thermometer increases. In 1968, twelve reference points were recommended as reference temperature points, covering the interval from the hydrogen triple point to the melting point of gold. Currently, Celsius temperature (t) is expressed as a relationship with absolute temperature (T), which is: T = 273.15 + t. The properties of a system that can be unambiguously expressed as functions of temperature, pressure, and concentration of the substances that make up the system are called thermodynamic functions. For example, heat capacity, internal energy, entropy, etc. If the change in the thermodynamic function depends only on the initial and final state of the system and does not depend on the path of the process, then such a function is called the state function of the system. A thermodynamic process is any change in a system associated with a change in at least one of the thermodynamic parameters. A circular process or cycle is a process in which a thermodynamic system, having left some initial state and undergoing a series of changes, returns to the same state; in this process, the change in any state parameter is equal to zero. ten

The beginning of physical chemistry was laid in the middle of the 18th century. The term "Physical chemistry", in the modern understanding of the methodology of science and questions of the theory of knowledge, belongs to M. V. Lomonosov, who for the first time read the "Course of True Physical Chemistry" to students of St. Petersburg University. In the preamble to these lectures, he gives the following definition: "Physical chemistry is a science that must, on the basis of the provisions and experiments of physical scientists, explain the reason for what happens through chemical operations in complex bodies." The scientist in the works of his corpuscular-kinetic theory of heat deals with issues that fully meet the above tasks and methods. This is precisely the nature of the experimental actions that serve to confirm individual hypotheses and provisions of this concept. M. V. Lomonosov followed these principles in many areas of his research: in the development and practical implementation of the “science of glass” founded by him, in various experiments devoted to confirming the law of conservation of matter and force (motion); - in works and experiments related to the doctrine of solutions - he developed an extensive program of research on this physical and chemical phenomenon, which is in the process of development to the present day.

This was followed by a break of more than a hundred years, and one of the first physicochemical studies in Russia in the late 1850s was started by D. I. Mendeleev.

The next course in physical chemistry was taught by N. N. Beketov at Kharkov University in 1865.

The first in Russia Department of Physical Chemistry was opened in 1914 at the Faculty of Physics and Mathematics of St. Petersburg University, in the fall, a student of D.P. Konovalov, M.S. Vrevsky, began to read the compulsory course and practical classes in physical chemistry.

The first scientific journal intended to publish articles on physical chemistry was founded in 1887 by W. Ostwald and J. van't Hoff.

The subject of physical chemistry

Physical chemistry is the main theoretical foundation of modern chemistry, using the theoretical methods of such important sections of physics as quantum mechanics, statistical physics and thermodynamics, nonlinear dynamics, field theory, etc. It includes the doctrine of the structure of matter, including: the structure of molecules, chemical thermodynamics, chemical kinetics and catalysis. As separate sections in physical chemistry, electrochemistry, photochemistry, physical chemistry of surface phenomena (including adsorption), radiation chemistry, the theory of metal corrosion, physical chemistry of macromolecular compounds (see polymer physics), etc. are also distinguished. Very closely adjacent to physical chemistry and are sometimes considered as its independent sections of colloid chemistry, physico-chemical analysis and quantum chemistry. Most sections of physical chemistry have fairly clear boundaries in terms of objects and methods of research, in terms of methodological features and the apparatus used.

The difference between physical chemistry and chemical physics

The content of the article

CHEMISTRY PHYSICAL, a branch of chemistry that studies the chemical properties of substances based on the physical properties of their constituent atoms and molecules. Modern physical chemistry is a broad interdisciplinary field bordering on various branches of physics, biophysics, and molecular biology. It has many points of contact with such branches of chemical science as organic and inorganic chemistry.

A distinctive feature of the chemical approach (as opposed to the physical and biological one) is that, along with the description of macroscopic phenomena, their nature is explained based on the properties of individual molecules and interactions between them.

New instrumental and methodological developments in the field of physical chemistry are used in other branches of chemistry and related sciences, such as pharmacology and medicine. Examples include electrochemical methods, infrared (IR) and ultraviolet (UV) spectroscopy, laser and magnetic resonance techniques, which are widely used in therapy and for the diagnosis of various diseases.

The main sections of physical chemistry are traditionally considered: 1) chemical thermodynamics; 2) kinetic theory and statistical thermodynamics; 3) questions of the structure of molecules and spectroscopy; 4) chemical kinetics.

Chemical thermodynamics.

Chemical thermodynamics is directly related to the application of thermodynamics - the science of heat and its transformations - to the problem of chemical equilibrium. The essence of the problem is formulated as follows: if there is a mixture of reagents (system) and the physical conditions in which it is located (temperature, pressure, volume) are known, then what spontaneous chemical and physical processes can bring this system to equilibrium? The first law of thermodynamics states that heat is a form of energy and that the total energy of a system (together with its environment) remains unchanged. Thus, this law is one of the forms of the law of conservation of energy. According to the second law, a spontaneously occurring process leads to an increase in the total entropy of the system and its environment. Entropy is a measure of the amount of energy that a system cannot use to do useful work. The second law indicates the direction in which the reaction will go without any external influences. To change the nature of the reaction (for example, its direction), you need to expend energy in one form or another. Thus, it imposes strict limits on the amount of work that can be done as a result of the conversion of heat or chemical energy released in a reversible process.

We owe important achievements in chemical thermodynamics to J. Gibbs, who laid the theoretical foundation of this science, which made it possible to combine the results obtained by many researchers of the previous generation into a single whole. The approach developed by Gibbs does not make any assumptions about the microscopic structure of matter, but considers the equilibrium properties of systems at the macro level. This is why one can think that the first and second laws of thermodynamics are universal and will remain valid even when we learn much more about the properties of molecules and atoms.

Kinetic theory and statistical thermodynamics.

Statistical thermodynamics (as well as quantum mechanics) allows one to predict the equilibrium position for some reactions in the gas phase. With the help of the quantum mechanical approach, it is possible to describe the behavior of complex molecules of a number of substances that are in liquid and solid states. However, there are reactions whose rates cannot be calculated either within the framework of the kinetic theory or with the help of statistical thermodynamics.

A real revolution in classical statistical thermodynamics took place in the 1970s. New concepts such as universality (the notion that members of some broad classes of compounds have the same properties) and the principle of similarity (estimation of unknown quantities based on known criteria) have led to a better understanding of the behavior of liquids near the critical point, when the distinction between liquid and gas. Using a computer, the properties of simple (liquid argon) and complex (water and alcohol) liquids in a critical state were simulated. More recently, the properties of liquids such as liquid helium (whose behavior is perfectly described in terms of quantum mechanics) and free electrons in molecular liquids have been comprehensively investigated using computer simulations (SUPERCONDUCTIVITY). This allowed a better understanding of the properties of ordinary liquids. Computer methods combined with the latest theoretical developments are intensively used to study the behavior of solutions, polymers, micelles (specific colloidal particles), proteins and ionic solutions. To solve problems of physical chemistry, in particular, to describe some properties of systems in a critical state and to study issues of high energy physics, the mathematical method of the renormalization group is increasingly being used.

The structure of molecules and spectroscopy.

Organic chemists of the 19th century. developed simple rules for determining the valency (ability to combine) of many chemical elements. For example, they found that the valence of carbon is 4 (one carbon atom can attach four hydrogen atoms to form a methane molecule CH 4), oxygen - 2, hydrogen - 1. Based on empirical ideas based on experimental data, assumptions were made about the spatial arrangement atoms in molecules (for example, the methane molecule has a tetrahedral structure, while the carbon atom is in the center of a triangular pyramid, and hydrogen is in its four vertices). However, this approach did not allow revealing the mechanism of formation of chemical bonds, and therefore, to estimate the size of molecules, to determine the exact distance between atoms.

Using spectroscopic methods developed in the 20th century, the structure of water molecules (H 2 O), ethane (C 2 H 6), and then much more complex molecules, such as proteins, was determined. The methods of microwave spectroscopy (EPR, NMR) and electron diffraction made it possible to establish the bond lengths, the angles between them (valence angles) and the mutual arrangement of atoms in simple molecules, and X-ray diffraction analysis - similar parameters for larger molecules that form molecular crystals. The compilation of catalogs of molecular structures and the use of simple concepts of valency laid the foundations of structural chemistry (L. Pauling was its pioneer) and made it possible to use molecular models to explain complex phenomena at the molecular level. If the molecules did not have a definite structure, or if the parameters of the C–C and C–H bonds in chromosomes were very different from those in the molecules of methane or ethane, then with the help of simple geometric models, J. Watson and F. Crick would not be able to build at the beginning 1950s for his famous double helix model of deoxyribonucleic acid (DNA). By studying the vibrations of atoms in molecules using IR and UV spectroscopy, it was possible to establish the nature of the forces that hold atoms in the composition of molecules, which, in turn, led to the idea of ​​the presence of intramolecular motion and made it possible to study the thermodynamic properties of molecules ( see above). This was the first step towards determining the rates of chemical reactions. Further, spectroscopic studies in the UV region helped to establish the mechanism of chemical bond formation at the electronic level, which made it possible to describe chemical reactions based on the idea of ​​the transition of reactants to an excited state (often under the action of visible or UV light). There was even a whole scientific field - photochemistry. Nuclear magnetic resonance (NMR) spectroscopy has made it possible for chemists to study individual stages of complex chemical processes and to identify active centers in enzyme molecules. This method also made it possible to obtain three-dimensional images of intact cells and individual organs. PHOTOCHEMISTRY.

Valency theory.

Using the empirical rules of valency developed by organic chemists, the periodic system of elements and Rutherford's planetary model of the atom, G. Lewis found that the key to understanding the chemical bond is the electronic structure of matter. Lewis came to the conclusion that a covalent bond is formed as a result of the socialization of electrons belonging to different atoms; in doing so, he proceeded from the idea that binding electrons are located on strictly defined electron shells. Quantum theory makes it possible to predict the structure of molecules and the nature of the covalent bonds formed in the most general case.

Our ideas about the structure of matter, which were formed due to the successes of quantum physics in the first quarter of the 20th century, can be summarized as follows. The structure of an atom is determined by the balance of electrical forces of repulsion (between electrons) and attraction (between electrons and a positively charged nucleus). Almost all the mass of an atom is concentrated in the nucleus, and its size is determined by the amount of space occupied by the electrons that revolve around the nuclei. Molecules consist of relatively stable nuclei held together by fast moving electrons, so that all chemical properties of substances can be explained in terms of the electrical interaction of elementary particles that make up atoms and molecules. Thus, the main provisions of quantum mechanics, concerning the structure of molecules and the formation of chemical bonds, create the basis for an empirical description of the electronic structure of matter, the nature of the chemical bond, and the reactivity of atoms and molecules.

With the advent of high-speed computers, it was possible to calculate (with a low but sufficient accuracy) the forces acting between atoms in small polyatomic molecules. The theory of valence, based on computer simulation, is currently a working tool for studying the structures, nature of chemical forces and reactions in cases where experiments are difficult or time consuming. This refers to the study of free radicals present in the atmosphere and flames or formed as reaction intermediates. There is hope that someday a theory based on computer calculations will be able to answer the question: how do chemical structures “calculate” their most stable state in a time of the order of picoseconds, while obtaining the corresponding estimates, at least in some approximation, requires a huge amount of machine time.

Chemical kinetics

deals with the study of the mechanism of chemical reactions and the determination of their rates. At the macroscopic level, the reaction can be represented as successive transformations, during which others are formed from one substance. For example, the seemingly simple transformation

H 2 + (1/2) O 2 → H 2 O

actually consists of several successive stages:

H + O 2 → OH + O

O + H 2 → HO + H

H + O 2 → HO 2

HO 2 + H 2 → H 2 O + OH

and each of them is characterized by its own rate constant k. S. Arrhenius suggested that the absolute temperature T and reaction rate constant k related by the ratio k = A exp(- E Act)/ RT, where BUT– pre-exponential factor (so-called frequency factor), E act - activation energy, R is the gas constant. For measuring k and T instruments are needed to track events that occur over a period of about 10–13 s, on the one hand, and over decades (and even millennia), on the other (geological processes); it is also necessary to be able to measure negligible concentrations of extremely unstable reagents. The task of chemical kinetics also includes the prediction of chemical processes occurring in complex systems (we are talking about biological, geological, atmospheric processes, combustion and chemical synthesis).

To study gas-phase reactions "in pure form" the method of molecular beams is used; in this case, molecules with strictly defined quantum states react with the formation of products that are also in certain quantum states. Such experiments provide information about the forces that cause certain reactions to occur. For example, in a molecular beam setup, even such small molecules as CH 3 I can be oriented in a given way and the collision rates in two “different” reactions can be measured:

K + ICH 3 → KI + CH 3

K + CH 3 I → KI + CH 3

where the CH 3 group is oriented differently with respect to the approaching potassium atom.

One of the issues that physical chemistry (as well as chemical physics) deals with is the calculation of reaction rate constants. Here, the transition state theory developed in the 1930s, which uses thermodynamic and structural parameters, is widely used. This theory, combined with the methods of classical physics and quantum mechanics, makes it possible to simulate the course of a reaction as if it were occurring under the conditions of an experiment with molecular beams. Experiments are being carried out on laser excitation of certain chemical bonds, which make it possible to test the correctness of the statistical theories of the destruction of molecules. Theories are being developed that generalize modern physical and mathematical concepts of chaotic processes (for example, turbulence). We are not so far from fully understanding the nature of both intra- and intermolecular interactions, revealing the mechanism of reactions occurring on surfaces with desired properties, and establishing the structure of the catalytic centers of enzymes and transition metal complexes. At the microscopic level, works on the formation kinetics of such complex structures as snowflakes or dendrites (crystals with a tree structure) can be noted, which stimulated the development of computer simulations based on simple models of the theory of nonlinear dynamics; this opens up prospects for creating new approaches to describing the structure and development of complex systems.

3rd ed., rev. - M.: Higher School, 2001 - 512 p., 319 p.

The textbook is compiled in accordance with the program in physical chemistry.

The following sections of the course are detailed in the first book: quantum mechanical foundations of the theory of chemical bonding, the structure of atoms and molecules, spectral methods for studying molecular structure, phenomenological and statistical thermodynamics, thermodynamics of solutions and phase equilibria.

In the second part of the section of the course of physical chemistry, electrochemistry, chemical kinetics and catalysis are presented on the basis of the concepts developed in the first part of the book - the structure of matter and statistical thermodynamics. The `Catalysis` section reflects the kinetics of heterogeneous and diffusion processes, adsorption thermodynamics and questions of reactivity.

For university students enrolled in chemical engineering specialties.

Book 1.

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Book 2.

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TABLE OF CONTENTS Book 1.
Preface. 3
Introduction 6
Section one. Quantum-mechanical substantiation of the theory of molecular structure and chemical bond
Chapter 1. The structure of the atom 9
§ 1.1. Quantum mechanical features of microparticles 9
§ 1.2. Hydrogen atom 11
§ 1.3. Atomic orbitals of a hydrogen-like atom 14
§ 1.4. Electron spin 21
§ 1.5. Multielectron atoms 23
§ 1.6. Pauli Principle 26
§ 1.7. Electronic configurations of atoms 28
Chapter 2. Molecules. Theoretical methods used in the study of the structure of molecules and chemical bonding 34
§ 2.1. Molecule. potential surface. Equilibrium configuration 34
§ 2.2. Theory of chemical bond and its problems. Schrödinger equation for molecules 39
§ 2.3. Variational method for solving the Schrödinger equation 42
§ 2.4. Two main methods of the theory of the structure of molecules. Valence bond method and molecular orbital method 44
§ 2.5. Basic ideas of the molecular orbital method 49
§ 2.6. Approximate description of the molecular orbital in the MO LCAO 50 method
§ 2.7. The II molecule in the MO LCAO method. Calculation of energy and wave function by the variational method 53
§ 2.8. Molecule H in the MO LCAO method. Covalent bond 58
Chapter 3. Diatomic molecules in the MO LCAO method 62
§ 3.1. Molecular orbitals of homonuclear diatomic molecules 62
§ 3.2. Electronic configurations and properties of homonuclear molecules formed by atoms of elements of the first and second periods 65
§ 3.3. Heteronuclear diatomic molecules 73
§ 3.4. polar connection. Electric dipole moment of a molecule 78
§ 3.5. Saturation of a covalent bond 81
§ 3.6. Donor-acceptor bond 82
§ 3.7. Ionic bond. The degree of polarity of the chemical bond 84
Chapter 4. Polyatomic molecules in the MO method 88
§ 4.1. Molecular orbitals in polyatomic molecules. Orbital symmetry. Delocalized and localized orbitals. HgO 88 molecule
§ 4.2. Description of the methane molecule. Delocalized and localized MOs. Hybridization of orbitals 95
§ 4.3. On the prediction of equilibrium configurations of molecules 99
§ 4.4. Nonrigid Molecules 101
§ 4.5. Molecules with multiple bonds in the MO LCAO method 104
§ 4.6. Hückel method 108
§ 4.7. Description of aromatic systems in the MOX 110 method
§ 4.8. Chemical bond in coordination compounds. Ligand field theory 117
§ 4.9. Ionic bonding in a crystal 126
Chapter 5. Intermolecular interaction 129
§ 5.1. Van der Waals forces. Other types of non-specific interactions 129
§ 5.2. Hydrogen bond 136
Section two. Spectral methods for studying the structure and energy states of molecules
Chapter 6. General information about molecular spectra. Elements of the theory of molecular spectra 141
§ 6.1. Intramolecular motion and electromagnetic spectrum. 141
§ 6.2. Molecular spectra of emission, absorption and Raman scattering. EPR and NMR spectra 145
§ 6.3. Rotational spectrum of a diatomic molecule (rigid rotator approximation) 150
§ 6.4. Vibrational-rotational spectrum of a diatomic molecule. Harmonic Oscillator Approximation 156
§ 6.5. The molecule is an anharmonic oscillator. Structure of the vibrational spectrum 162
§ 6.6. Electronic spectra. Determination of the dissociation energy of diatomic molecules 169
§ 6.7. Rotational spectra and strict polyatomic molecules.... 171
§ 6.8. Vibrations, spectrum and structure of polyatomic molecules 175
§ 6.9. Use of vibrational spectra to determine the structure of molecules 180
§ 6.10. Influence of the intermolecular interaction of the medium and state of aggregation on the vibrational spectrum 183
Section three. Chemical thermodynamics
Chapter 7. General concepts. The first law of thermodynamics and its application 186
§ 7.1. Subject and tasks of chemical thermodynamics 186
§ 7.2. Basic concepts and definitions of chemical thermodynamics 188
§ 7.3. First law of thermodynamics. Non-circular processes 199
§ 7.4. Heat capacity 202
§ 7.5. Influence of temperature on heat capacity. Temperature series.. 208
§ 7.6. Quantum theory of heat capacity of crystalline matter 211
§ 7.7. Quantum-statistical theory of the heat capacity of a gaseous substance 215
§ 7.8. thermal effects. Hess Law 217
§ 7.9. Application of Hess' law to the calculation of thermal effects 220
§ 7.10. Dependence of thermal effect on temperature. Kirchhoff equation 227
Chapter 8. The second law of thermodynamics and its application 235
§ 8.1. Spontaneous and non-spontaneous processes. The Second Law of Thermodynamics 235
§ 8.2. Entropy 236
§ 8.3. Entropy change in non-static processes 239
§ 8.4. Entropy change as a criterion of directionality and equilibrium in an isolated "system 240
§ 8.5. Characteristic functions. Thermodynamic potentials 241
§ 8.6. Criteria for the possibility of a spontaneous process and equilibrium in closed systems 249
§ 8.7. Entropy change in some processes 251
§ 8.8. Gibbs energy of a mixture of ideal gases. Chemical potential 261
§ 8.9. General conditions of chemical equilibrium 265
§ 8.10. The law of active masses. Equilibrium constant for gas phase reactions 266
§ 8.11. Reaction isotherm equation 271
§ 8.12. Using the law of mass action to calculate the composition of an equilibrium mixture 273
§ 8.13. Effect of temperature on chemical equilibrium. Reaction isobar equation 282
§ 8.14. Integral form of dependence of Gibbs energy and equilibrium constant on temperature 284
§ 8.15. Chemical equilibrium in heterogeneous systems 286
Chapter 9. The Third Law of Thermodynamics and the Calculation of Chemical Equilibrium 289
§ 9.1. Thermal Nernst theorem. Third law of thermodynamics 289
§ 9.2. Calculation of the change in standard Gibbs energy and equilibrium constant by the method of Temkin - Schwartzman 294
§ 9.3. Calculation of the change in the standard Gibbs energy and the equilibrium constant using the functions of the reduced Gibbs energy 297
§ 9.4. Adiabatic reactions 299
Chapter 10. Chemical equilibrium in real systems 303
§ 10.1. Fugacity and coefficient of fugacity of gases 303
§ 10.2. Calculation of chemical equilibrium in a real gas system at high pressures 312
§ 10.3. Calculation of chemical equilibrium in systems in which several reactions occur simultaneously 314
Chapter 11. Introduction to statistical thermodynamics 320
§ 11.1. Statistical physics and statistical thermodynamics. Macroscopic and microscopic description of the state of the system 320
§ 11.2. Microscopic description of the state by the method of classical mechanics 323
§ 11.3. Microscopic description of the state by the method of quantum mechanics. Quantum statistics 324
§ 11.4. Two types of averages (microcanonical and canonical averages) 325
§ 11.5. Relationship between entropy and statistical weight. Statistical nature of the second law of thermodynamics 326
§ 11.6. Thermostat system. Canonical Gibbs distribution. 330
§ 11.7. The sum over the states of the system and its connection with energy. Helmholtz 335
§ 11.8. Sum over particle states 337
§ 11.9. Expression of thermodynamic functions in terms of the sum over the states of the system 340
§ 11.10. The sum over the states of a system of one-dimensional harmonic oscillators. Thermodynamic properties of a monatomic solid according to Einstein's theory 343
§ 11.11. Boltzmann quantum statistics. Maxwell's law of molecular velocity distribution 346
§ 11.12. Fermi - Dirac and Bose - Einstein statistics 352
§ 11.13. General formulas for calculating thermodynamic functions from molecular data 353
§ 11.14 Calculation of the thermodynamic functions of an ideal gas under the assumption of rigid rotation and harmonic vibrations of molecules 357
Section four. Solutions
Chapter 12. General characteristics of solutions 365
§ 12.1. Classification of mortars 365
§ 12.2. Concentration of solutions 367
5 12.3. Specificity of solutions. The role of intermolecular and chemical interactions, the concept of solvation 368
§ 12.4. The main directions in the development of the theory of solutions 372
§ 12.5. Thermodynamic conditions for the formation of solutions 374
§ 12.6. Partial molar values ​​375
§ 12.7. Basic Methods for Determining Partial Molar Values ​​379
§ 12.8. Partial and relative partial molar enthalpies 381
§ 12.9. Heats of dissolution and dilution 382
§ 12.10. Thermodynamic properties of ideal liquid solutions 386
§ 12.11.3 Raoult law 390
§ 12.12. Boiling point of an ideal solution 392
§ 12.13. Freezing point of an ideal solution 395
§ 12.14.0 smotic pressure of an ideal solution 397
§ 12.15 Non-ideal solutions 400
§ 12.16. Extremely dilute, regular and athermal solutions 402
§ 12.17. Activity. Activity coefficient. Standard state 404
§ 12.18.0smotic coefficient 407
§ 12.19. Methods for determining activities 409
§ 12.20. Relationship of the activity and activity coefficient with the thermodynamic properties of the solution and excess thermodynamic functions 412
Section Five. Phase Equilibria
Chapter 13. Thermodynamic theory of phase equilibria 415
§ 13.1. Basic concepts 415
§ 13.2. Phase equilibrium conditions 418
§ 13.3. Gibbs phase rule 419
Chapter 14 Single Component Systems 421
§ 14.1. Application of the Gibbs phase rule to one-component systems 421
§ 14.2. Phase transitions of the first and second kind 422
§ 14.3. Equation of Clapeyron - Clausius 425
§ 14.4. Saturated steam pressure 423
§ 14.5. State diagrams of one-component systems 429
§ 14.6. Carbon dioxide state diagram 431
§ 14.7. Water Status Diagram 432
§ 14.8. Sulfur state diagram 433
§ 14.9. Enantiotropic and monotropic phase transitions 435
Chapter 15. Two-component systems 436
§ 15.1. Physical and chemical analysis method 436
§ 15.2. Application of the Gibbs phase rule to two-component systems 437
§ 15.3. Equilibrium gas - liquid solution in two-component systems 438
§ 15.4. Equilibrium liquid - liquid in two-component systems 442
§ 15.5. Equilibrium vapor - liquid solution in two-component systems 444
§ 15.6. Physical and chemical bases of solution distillation 453
§ 15.7. Equilibrium crystals - liquid solution in two-component systems 457
§ 15.8. Equilibrium liquid - gas and crystals - gas (steam) in two-component systems 476
§ 15-9. State Diagram Calculations 476
Chapter 16. Three-component systems 482
§ 16.1. Application of the Gibbs phase rule to three-component systems 482
§ 16.2. Graphical representation of the composition of a three-component system 482
§ 16.3. Equilibrium crystals - liquid solution in three-component systems 484
§ 16.4. Equilibrium liquid - liquid in three-component systems 489
§ 16.5. Distribution of a solute between two liquid phases. Extraction 491
Appendix 495
Index 497

TABLE OF CONTENTS Book 2.
Preface 3
Section six. Electrochemistry
Chapter 17. Solutions, electrolytes 4
§ 17.1. Electrochemistry subject 4
§ 17.2. Specificity of electrolyte solutions 5
§ 17.3. Electrolytic dissociation in solution 6
§ 17.4. Average ionic activity and activity factor 10
§ 17.5. Basic concepts of the electrostatic theory of strong electrolytes Debye and Hückel 13
§ 17.6. Basic concepts of ion association theory 22
§ 17.7. Thermodynamic properties of ions 24
§ 17.8. Thermodynamics of ionic solvation 28
Chapter 18. Non-equilibrium phenomena in electrolytes. Electrical conductivity of electrolytes 30
§ 18.1. Basic concepts. Faraday's laws 30
§ 18.2. Movement of ions in an electric field. Ion transport numbers. 32
§ 18.3. Electrical conductivity of electrolytes. Electrical conductivity 37
§ 18.4. Electrical conductivity of electrolytes. Molar electrical conductivity 39
§ 18.5. Molar electrical conductivity of hydronium and hydroxide ions 43
§ 18.6. Electrical conductivity of non-aqueous solutions 44
§ 18.7. Electrical conductivity of solid and molten electrolytes 46
§ 18.8. Conductometry 47
Chapter 19. Equilibrium electrode processes 49
§ 19.1. Basic concepts 49
§ 19.2. EMF of an electrochemical system. Electrode potential 51
§ 19.3. Occurrence of a potential jump at the solution-metal interface 53
§ 19.4. Diffusion potential 55
§ 19.5. The structure of the electrical double layer at the solution-metal interface 56
§ 19.6. Thermodynamics of reversible electrochemical systems 60
§ 19.7. Classification of reversible electrodes 64
§ 19.8. Electrode potentials in non-aqueous solutions 74
§ 19.9. Electrochemical circuits 75
§ 19.10. Application of the theory of electrochemical systems to the study of equilibrium in solutions 82
§ 19.11. Potentiometry 85
Section seven. Kinetics of chemical reactions
Chapter 20. Laws of chemical kinetics 93
§ 20.1. General concepts and definitions 93
§ 20.2. Chemical reaction rate 95
§ 20.3. The law of mass action and the principle of independence of reactions 101
Chapter 21. Kinetics of chemical reactions in closed systems. 105
§ 21.1. Unilateral first order reactions 105
§ 21.2. Unilateral Second Order Reactions 109
§ 21.3. One-way reactions of the nth order 111
§ 21.4. Methods for determining the order of the reaction 112
§ 21.5. Bilateral reactions of the first order 113
§ 21.6. Bilateral reactions of the second order 116
§ 21.T. Parallel one-way reactions 117
§ 21.8. Unilateral sequential reactions 119
§ 21.9. Method of quasi-stationary concentrations 125
Chapter 22. Kinetics of reactions in open systems 127
§ 22.1. Reaction kinetics in a perfectly mixed reactor 127
§ 22.2. Reaction kinetics in a plug flow reactor 129
Chapter 23. The theory of the elementary act of chemical interaction 133
§ 23.1. Elementary chemical act 133
§ 23.2. Theory of active collisions 137
§ 23.3. Theory of the activated complex 141
§ 23.4. Preexponential factor in the Arrhenius equation according to the transition state theory 154
§ 23.5. MO symmetry and activation energy of chemical reactions 159
Chapter 24. Kinetics of reactions in solutions, chain and photochemical reactions 166
§ 24.1. Features of the kinetics of reactions in solutions 166
§ 24.2. Influence of medium on the reaction rate constant 170
§ 24.3. Kinetics of ionic reactions in solutions 178
§ 24.4. Chain reactions 181
§ 24.5. Photochemical reactions 189
Chapter 25. Kinetics of electrode processes 196
§ 25.1. The rate of an electrochemical reaction. exchange current 196
§ 25.2. Electrode polarization 197
§ 25.3. Diffusion overvoltage 199
§ 25.4. Electrochemical overvoltage 205
§ 25.5. Other types of overvoltage 210
5 25.6. Temperature-kinetic method for determining the nature of polarization in electrochemical processes 211
§ 25.7. Overvoltage during electrolytic hydrogen evolution 213
§ 25.8. Electrolysis. Decomposition voltage 217
§ 25.9. Polarization phenomena in chemical sources of electric current 220
§ 25.10. Electrochemical corrosion of metals. passivity of metals. Corrosion protection methods 222
Section eight. Catalysis
Chapter 26. Principles of catalytic action 228
§ 26.1. Basic concepts and definitions 228
§ 26.2. Features of the kinetics of catalytic reactions 232
§ 26.3. Activation energy of catalytic reactions 237
§ 26.4. Interaction of reagents with a catalyst and principles of catalytic action 241
Chapter 27. Homogeneous catalysis 245
§ 27.1. Acid-base catalysis 246
§ 27.2. Redox Catalysis 255
§ 27.3. Enzymatic catalysis 260
§ 27.4. Autocatalysis, inhibition and periodic catalytic reactions 266
§ 27.5. Application in industry and prospects for the development of homogeneous catalysis 271
Chapter 28. Heterogeneous catalysis. 273
§ 28.1. Surface structure of heterogeneous catalysts 273
§ 28.2. Adsorption as a stage of heterogeneous catalytic reactions 277
§ 28.3. Mechanism of heterogeneous catalytic reactions 282
§ 28.4. Kinetics of heterogeneous catalytic reactions on an equally accessible surface 285
§ 28.5. Macrokinetics of heterogeneous catalytic processes 292
§ 28.6. Application of heterogeneous catalysis in industry 300
Literature 303
Appendix 305
Index 312
Contents 316

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