PHYSICAL CHEMISTRY

§ 1. The subject of physical chemistry. Its meaning

The relationship of chemical and physical phenomena studies physical chemistry. This branch of chemistry is the boundary between chemistry and physics. Using the theoretical and experimental methods of both sciences, as well as its own methods, physical chemistry is engaged in a multifaceted study of chemical reactions and the physical processes accompanying them. Since, however, even a multifaceted study is never complete and does not cover the phenomenon in an exhaustive way, the laws and laws of physical chemistry, like those of other natural sciences, always simplify the phenomenon and do not fully reflect it.

The rapid development and growing importance of physical chemistry are associated with its boundary position between physics and chemistry. The main general task of physical chemistry is the prediction of the time course of the process and the final result (equilibrium state) under various conditions based on data on the structure and properties of the substances that make up the system under study.

§ 2. Brief outline of the history of the development of physical chemistry

The term "physical chemistry" and the definition of this science were first given by M.V. Lomonosov, who in 1752-1754. read a course in physical chemistry to the students of the Academy of Sciences and left the manuscript of this course "Introduction to True Physical Chemistry" (1752). Lomonosov carried out many studies, the topics of which correspond to the "Plan for the course of physical chemistry" compiled by him (1752) and the program of experimental work "Experience in Physical Chemistry" (1754). Under his leadership, a student workshop in physical chemistry was also held.

Lomonosov gave the following definition of physical chemistry: "Physical chemistry is a science that explains, on the basis of the provisions and experiments of physics, what happens in mixed bodies during chemical operations." This definition is close to modern.

For the development of physical chemistry, the discovery of two laws of thermodynamics in the middle of the 19th century (S. Carnot, Yu.R. Mayer, G. Helmholtz, D.P. Joule, R. Clausius, W. Thomson) was of great importance.

The number and variety of research, lying in the field that borders between physics and chemistry, constantly increased in the 19th century. The thermodynamic theory of chemical equilibrium was developed (K.M. Guldberg, P. Waage, D.W. Gibbs). The studies of L.F. Wilhelmi laid the foundation for the study of the rates of chemical reactions (chemical kinetics). The transfer of electricity in solutions was studied (I.V. Gittorf, F.V.G. Kolrausch), the laws of equilibrium of solutions with steam were studied (D.P. Konovalov) and the theory of solutions was developed (D.I. Mendeleev).

The recognition of physical chemistry as an independent science and academic discipline was expressed in the establishment at the University of Leipzig (Germany) in 1887 of the first department of physical chemistry headed by W. Ostwald and in the foundation of the first scientific journal on physical chemistry there. At the end of the 19th century, the University of Leipzig was the center for the development of physical chemistry, and the leading physical chemists were W. Ostwald, J. H. Van't Hoff, S. Arrhenius and W. Nernst. By this time, three main sections of physical chemistry were defined - chemical thermodynamics, chemical kinetics and electrochemistry.

The most important areas of science, the development of which is a necessary condition for technical progress, include the study of chemical processes; physical chemistry plays a leading role in the development of this problem.

§ 3. Sections of physical chemistry. Research methods

Chemical thermodynamics. In this section, on the basis of the laws of general thermodynamics, the laws of chemical equilibrium and the doctrine of phase equilibria are expounded.

The doctrine of solutions aims to explain and predict the properties of solutions (homogeneous mixtures of several substances) on the basis of the properties of the substances that make up the solution.

The doctrine of surface phenomena. Various properties of surface layers of solids and liquids (interfaces between phases) are studied; one of the main studied phenomena in the surface layers is adsorption(accumulation of matter in the surface layer).

In systems where the interfaces between liquid, solid, and gaseous phases are highly developed (emulsions, mists, smokes, etc.), the properties of the surface layers become of primary importance and determine many of the unique properties of the entire system as a whole. Such dispersed (microheterogeneous) systems are being studied colloid chemistry, which is a major independent branch of physical chemistry.

The above list of the main sections of physical chemistry does not cover some areas and smaller sections of this science, which can be considered as parts of larger sections or as independent sections of physical chemistry. It should be emphasized once again the close interrelationship between the various branches of physical chemistry. In the study of any phenomenon, one has to use an arsenal of ideas, theories and methods for studying many branches of chemistry (and often other sciences). Only with an initial acquaintance with physical chemistry is it possible for educational purposes to distribute the material into the indicated sections.

Methods of physical and chemical research. The basic methods of physical chemistry are naturally the methods of physics and chemistry. This is, first of all, an experimental method - the study of the dependence of the properties of substances on external conditions, the experimental study of the laws of the flow of various processes and the laws of chemical equilibrium.

The theoretical understanding of experimental data and the creation of a coherent system of knowledge is based on the methods of theoretical physics.

The thermodynamic method, which is one of them, makes it possible to quantitatively relate various properties of a substance (“macroscopic” properties) and calculate some of these properties based on the experimental values ​​of other properties.

CHAPTER I
THE FIRST LAW OF THERMODYNAMICS

§ 1. Energy. The law of conservation and transformation of energy

An integral property (attribute) of matter is movement; it is indestructible, like matter itself. The motion of matter manifests itself in different forms, which can pass one into another. The measure of motion of matter is energy. Quantitatively, energy is expressed in a certain way through the parameters characteristic of each specific form of movement, and in units specific to this form.

In the SI system of units, the unit of energy (heat and work) is the joule ( J), equal to the work of force in 1 H on the way to 1 m. 1 J = 1 Nm.

The widely used unit of energy (heat), the calorie, is currently an off-system unit that is allowed for use. The currently used calorie, by definition, equates to a certain number of joules: 1 feces equals 4.1868 joules. This unit is used in heat engineering and can be called thermal calorie. In chemical thermodynamics, a slightly different unit is used, equated to 4.1840 joules and called thermochemical calorie. The expediency of its application is connected with the convenience of using the extensive experimental thermochemical material collected in reference books and expressed in these units.

When one form of motion is transformed into another, the energies of the disappeared and appeared motion, expressed in different units, are equivalent to each other, i.e., the energy of the disappeared motion is in a constant quantitative relation to the energy of the motion that has arisen (the law of equivalent transformations of energy). This ratio does not depend on the energies of the two forms of motion and on the specific conditions under which the transition from one form of motion to another took place. So, when the energy of an electric current is converted into the energy of chaotic molecular motion, one joule of electrical energy always turns into 0.239 feces energy of molecular motion.

Thus, energy as a measure of the motion of matter always manifests itself in a qualitatively original form, corresponding to a given form of motion, and is expressed in the appropriate units of measurement. On the other hand, it quantitatively reflects the unity of all forms of movement, their mutual convertibility and the indestructibility of movement.

The above law of equivalent transformations of energy is a physical experimental law. The law of equivalent energy transformations can be expressed differently, namely in the form the law of conservation and transformation of energy: energy is neither created nor destroyed; in all processes and phenomena, the total energy of all parts of an isolated material system participating in this process does not increase or decrease, remaining constant.

The law of conservation and transformation of energy is universal in the sense that it is applicable to phenomena occurring in arbitrarily large bodies, representing an aggregate of a huge number of molecules, and to phenomena occurring with the participation of one or a few molecules.

For various forms of mechanical motion, the law of conservation of energy has long been expressed in a qualitative form (Descartes - 1640) and a quantitative form (Leibniz - 1697).

For the mutual transformations of heat and work (see below), the law of conservation of energy was proved as a natural science law by the studies of Yu. R. Mayer, G. Helmholtz and D.P. Joule, carried out in the forties of the XIX century.

Using the law of equivalent transformations, it is possible to express the energies of various forms of motion in units characteristic of one type of energy (one form of motion), and then perform operations of addition, subtraction, etc.

§ 2. Subject, method and limits of thermodynamics

Thermodynamics is one of the main branches of theoretical physics. Thermodynamics studies the laws of mutual transformations of various types of energy associated with the transfer of energy between bodies in the form of heat and work. Focusing its attention on heat and work as forms of energy transfer in a variety of processes, thermodynamics involves numerous energy connections and dependencies between various properties of a substance in its circle of consideration and gives very widely applicable generalizations called the laws of thermodynamics.

When establishing the basic thermodynamic laws, energy transformations (often very complex) occurring inside the body are usually not detailed. The types of energy inherent in the body in its given state are also not differentiated; the totality of all these types of energy is considered as a single internal energy of the system .

The subject matter of thermodynamics outlined above defines the method and boundaries of this science. The distinction between heat and work, taken as a starting point by thermodynamics, and the opposition of heat to work makes sense only for bodies consisting of many molecules, since for one molecule or for a set of a small number of molecules, the concepts of heat and work lose their meaning. Therefore, thermodynamics considers only bodies consisting of a large number of molecules, the so-called macroscopic systems moreover, thermodynamics in its classical form does not take into account the behavior and properties of individual molecules.

The thermodynamic method is also characterized by the fact that the object of study is a body or a group of bodies isolated from the material world into thermodynamic system (hereinafter referred to simply system).

The system has certain boundaries separating it from the outside world (environment).

The system is homogeneous , if each of its parameters has the same value in all parts of the system or continuously changes from point to point.

The system is heterogeneous , if it consists of several macroscopic (consisting in turn of many molecules) parts, separated from one another by visible interfaces. On these surfaces, some parameters change abruptly. Such, for example, is the system "solid salt - saturated aqueous salt solution - saturated water vapor." Here, at the boundaries of salt - solution and solution - vapor, the composition and density change abruptly.

Homogeneous parts of the system, separated from other parts by visible interfaces, are called phases . In this case, the set of individual homogeneous parts of the system that have the same physical and thermodynamic properties is considered to be one phase (for example, a set of crystals of one substance or a set of liquid droplets suspended in a gas and forming fog). Each phase of the system is characterized by its own equation of state.

A system that cannot exchange matter and energy with the environment (in the form of heat or work) is called isolated .

A system that can exchange matter and energy with the environment (in the form of heat or work) is called open.

A system that cannot exchange matter with the environment, but can exchange energy (in the form of heat or work) is called closed .

Thermodynamics studies the relationship between such measurable properties of a material system as a whole and its macroscopic parts (phases), such as temperature, pressure, mass, density and chemical composition of the phases included in the system, and some other properties, as well as the relationship between changes in these properties.

The set of properties studied by thermodynamics (the so-called thermodynamic parameters of the system) defines thermodynamic state of the system. A change in any thermodynamic properties (even if only one) leads to a change in the thermodynamic state of the system.

All processes occurring in nature can be divided into spontaneous (natural) and non-spontaneous.

Spontaneous processes These are processes that do not require external energy input. For example, the transfer of heat from a body with a higher temperature to a body with a lower temperature, the dissolution of salt in water, etc., proceed by themselves.

Non-spontaneous processes require energy from the outside for their flow, for example, the separation of air into nitrogen and oxygen.

Thermodynamics mainly considers such states of a system in which its parameters (temperature, pressure, electrostatic potential, etc.) do not change spontaneously with time and have the same value at all points in the volume of individual phases. Such states are called balanced.

One of the basic postulates of thermodynamics is the statement that the course of any spontaneous process ultimately brings the isolated system to an equilibrium state, when its properties will no longer change, i.e., equilibrium will be established in the system.

States characterized by uneven and time-varying distributions of temperature, pressure, and composition within phases are nonequilibrium. They are considered by the thermodynamics of non-equilibrium (irreversible) processes, in which, in addition to the basic thermodynamic laws, additional assumptions are used.

Thermodynamics, built on the basis of the basic laws of thermodynamics, which are considered as a generalization of experience, is often called classical or phenomenological thermodynamics. Thermodynamics provides the theoretical foundations for the theory of heat engines; this section is called technical thermodynamics. The study of chemical processes from a thermodynamic point of view is engaged in chemical thermodynamics, which is one of the main branches of physical chemistry.

§ 3. Heat and work

Changes in the forms of motion during its transition from one body to another and the corresponding transformations of energy are very diverse. The forms of the transition of motion itself and the transitions of energy connected with it can be divided into two groups.

The first group includes only one form of motion transition by chaotic collisions of molecules of two adjoining bodies, i.e. by conduction (and at the same time by radiation). The measure of the movement transmitted in this way is heat .

The second group includes various forms of movement transition, a common feature of which is the movement of macroscopic masses under the action of any external forces that have a directed character. Such are the rise of bodies in a gravitational field, the transition of a certain amount of electricity from a larger electrostatic potential to a smaller one, the expansion of a gas under pressure, etc. The general measure of the movement transmitted by such means is Work .

Heat and work characterize qualitatively and quantitatively two different forms of transmission of motion from one part of the material world to another.

The transmission of motion is a kind of complex motion of matter, the two main forms of which we distinguish. Heat and work are measures of these two complex forms of motion of matter, and they should be considered as types of energy.

The common property of heat and work is that they matter only during the time intervals in which these processes take place. In the course of such processes, in some bodies, movement in one form or another decreases and the corresponding energy decreases, while in other bodies movement in the same or other forms increases and the corresponding types of energy increase.

We are not talking about the stock of heat or work in any body, but only about the heat and work of a known process. After its completion, there is no need to talk about the presence of heat or work in the bodies.

§ 4. Equivalence of heat and work

A constant equivalent ratio between heat and work during their mutual transitions was established in the classical experiments of D.P. Joule (1842-1867). A typical Joule experiment is as follows.

Joule device for determining the mechanical equivalent of heat.

Weights falling from a known height rotate a stirrer immersed in water in a calorimeter (a weight and a calorimeter with water constitute a thermodynamic system.) The rotation of the stirrer blades in water causes the water in the calorimeter to heat up; the corresponding rise in temperature is quantified.

After the specified process is completed, the system must be brought to its original state. This can be done through mental experience. The weights rise to their original height, while external work is expended, which increases the energy of the system. In addition, heat is removed from the calorimeter (transferred to the environment) by cooling it to the initial temperature. These operations return the system to its original state, i.e., all measurable properties of the system acquire the same values ​​that they had in the initial state. The process during which the properties of the system changed, and at the end of which it returned to its original state, is called circular (cyclic) process or cycle .

The only result of the described cycle is the removal of work from the environment surrounding the system, and the transfer to this environment of the heat taken from the calorimeter.

Comparison of these two quantities, measured in the corresponding units, shows a constant relationship between them, independent of the size of the load, the size of the calorimeter, and the specific amounts of heat and work in different experiments.

It is advisable to write the heat and work in a cyclic process as the sum (integral) of infinitely small (elementary) heats  Q and infinitesimal (elementary) jobs  W, and the initial and final limits of integration coincide (cycle).

Then the equivalence of heat and work in a cyclic process can be written as follows:

(I, 1)

In equation (I, 1), the sign denotes integration over a cycle. Coefficient constancy k reflects the equivalence of heat and work ( k is the mechanical equivalent of heat). Equation (I, 1) expresses the law of conservation of energy for a particular, very important case of the transformation of work into heat.

In the studies of Joule, Rowland (1880), Miculescu (1892), and others, the methods of friction in metals, impact, direct conversion of the work of an electric current into heat, stretching of solids, etc. were used. k always constant within the experimental error.

In what follows, it is always assumed that work and heat, with the help of the coefficient k expressed in the same units (no matter what) and the coefficient k goes down.

§ 5. Internal energy

For a non-circular process, the equality (I, 1) is not observed, since the system does not return to its original state. Instead, the equalities for a non-circular process can be written (omitting the coefficient k):

≠

Since the limits of integration are generally arbitrary, then for elementary quantities  W and  Q:

Q   W,

Consequently:

Q – W  0

Denote the difference  Q –  W for any elementary thermodynamic process through dU:

dU   Q – W (I, 2)

or for the final process:

–
(I, 2a)

Returning to the circular process, we obtain (from Equation I, 1):

=
–
= 0 (I, 3)

Thus, the value dU is the total differential of some system state function. When the system returns to its original state (after a cyclic change), the value of this function acquires its original value.

System state functionU , defined by the equalities (I, 2) or (I, 2a) is calledinternal energy systems .

Obviously, expression (I, 2a) can be written as follows:

= U 2 – U 1 = ∆ U = – (I, 2b)

U 2 – U 1 = ∆U = Q – W

This reasoning substantiates empirically the presence of a certain function of the state of the system, which has the meaning of the total measure of all movements that the system possesses.

In other words, internal energy includes the translational and rotational energy of molecules, the vibrational energy of atoms and groups of atoms in a molecule, the energy of electron motion, intranuclear and other types of energy, i.e. the totality of all types of particle energy in the system, with the exception of the potential and kinetic energy of the system itself .

Let us assume that the cyclic process was carried out in such a way that after the system returned to its initial state, the internal energy of the system did not take the initial value, but increased. In this case, the repetition of circular processes would cause the accumulation of energy in the system. It would be possible to convert this energy into work and obtain work in this way not at the expense of heat, but “out of nothing”, since in a circular process work and heat are equivalent to each other, which is shown by direct experiments.

Inability to complete the specified build cycle perpetuum mobile (perpetuum mobile) of the first kind, that gives work without spending an equivalent amount of another type of energy, is proved by the negative result of thousands of years of human experience. This result leads to the same conclusion that we obtained in a particular but more rigorous form by analyzing Joule's experiments.

Let us formulate the result obtained once again. The total energy supply of the system (its internal energy) as a result of a cyclic process returns to its original value, i.e., the internal energy of the system located in given state, has one specific value and does not depend on what changes the system was subjected to before reaching this state.

In other words, the internal energy of the system is a single-valued, continuous and finite function of the state of the system.

The change in the internal energy of the system is determined by expression (I, 2b); the expression (I, 3) is valid for a circular process. With an infinitesimal change in some properties (parameters) of the system, the internal energy of the system also changes infinitesimally. This is a property of a continuous function.

Within thermodynamics, there is no need to use general definition concept of internal energy. A formal quantitative definition through expressions (I, 2) or (I, 2a) is sufficient for all further thermodynamic reasoning and conclusions.

Since the internal energy of the system is a function of its state, then, as already mentioned, the increase in internal energy with infinitesimal changes in the parameters of the system states is the total differential of the state function. Breaking the integral in equation (I, 3) into two integrals over the sections of the path from the state 1 up to the state 2 (path "a") (see Fig. I) and vice versa - from the state 2

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Basic lecture notes

on academic discipline: "EN.03 Chemistry"

specialty: "260502 Technology of catering products"

"Physical and colloid chemistry"

annotation

Compiled by: Ivanova L.V.

Reviewers:

Polyanskaya T.V., teacher of natural disciplines, FGOU SPO "OKTES";

Chudnovskaya VG, lecturer, chairman of the PUK of chemical disciplines.

The reference abstract of lectures was compiled in accordance with the requirements of the Federal State Educational Standard (FSES) of secondary vocational education to a minimum of content in the discipline "EN.03 Chemistry" for the training of mid-level specialists: "260807 Technology of public catering products."

Working with the basic lecture notes contributes to the transition from the usual descriptive perception of physical and chemical data to quantitative representations, i.e. leads to a deep and correct understanding of them and, as a result, to the predictability of the processes occurring in colloidal and other systems. This helps to professionally develop, using the scientific foundations of physical and colloidal chemistry, approaches to the technology of obtaining, storing and processing food.

The manual is intended for organizing classroom and extracurricular work of students in the discipline "EN.03 Chemistry" (section 1 "Physical chemistry", section 3 "Colloid chemistry").

Introduction

Section 1. Physical chemistry

1.1 Basic concepts and laws of thermodynamics. Thermochemistry

1.1.1 Basic concepts of thermodynamics

1.1.2 First law of thermodynamics

1.1.3 Thermochemistry

1.1.4 The second law of thermodynamics

1.2 Aggregate states of substances, their characteristics

1.2.1 Characteristics of the gaseous state of matter

1.2.2 Characteristics of the liquid state of matter

1.2.3 Characterization of the solid state of matter

1.3 Chemical kinetics and catalysis. Chemical equilibrium

1.3.1 The rate of a chemical reaction

1.3.2 Catalysis and catalysts

1.3.3 Chemical equilibrium

1.4 Properties of solutions

1.4.1 general characteristics solutions

1.4.2 Solutions of gases in liquids

1.4.3 Mutual solubility of liquids

1.4.4 Solutions of solids in liquids

1.4.5 Diffusion and osmosis in solutions

1.4.6 Saturated vapor pressure of solution

1.4.7 Freezing and boiling solutions

1.4.8 Properties of electrolyte solutions

1.5 Surface phenomena. Adsorption

1.5.1 Adsorption, its types

1.5.2 Adsorption at the solution-gas interface

1.5.3 Ion exchange adsorption

Section 2. Colloid chemistry

2.1 Subject of colloidal chemistry. Disperse systems

2.1.1 General characteristics dispersed systems

2.1.2 Classification of disperse systems

2.2 Colloidal solutions

2.2.1 Acquisition methods

2.2.2 The structure of a colloidal particle

2.2.3 Properties of colloidal solutions

2.3 Coarse systems

2.3.2 Suspensions

2.3.3 Emulsions

2.3.4 Aerosols

2.4 Physical and chemical changes organic matter food products

2.4.1 Proteins, their chemical structure and amino acid composition

2.4.2 Carbohydrates - high molecular weight polysaccharides

2.4.4 Jelly

Bibliographic list

Introduction

Physical chemistry is a science that studies the relationship between the chemical and physical properties of substances, chemical and physical phenomena and processes.

Only on the basis of the laws of physical chemistry can such common in various industries be understood and implemented. Food Industry processes such as evaporation, crystallization, drying, sublimation, separation, distillation, extraction and dissolution. Without knowledge of the methods of physical chemistry, technological control of food production is impossible: determination of moisture, acidity, content of sugars, proteins, fats, vitamins, etc.

The founder of physical chemistry is M.V. Lomonosov. He in 1752-1754. He was the first scientist to give students a course in physical chemistry. The reading of the course was accompanied by a demonstration of experiments and laboratory work. Lomonosov was the first to propose the term "physical chemistry" and gave this scientific discipline the following definition: "Physical chemistry is a science that explains, on the basis of the provisions and experiments of physics, what happens in mixed bodies during chemical operations." Thus, M.V. Lomonosov considered physical chemistry as a science designed to give a physical explanation of the essence of chemical processes.

M.V. Lomonosov wrote the world's first textbook on physical chemistry. The discovery by the great scientist of the law of conservation of matter and energy, the doctrine of the existence of absolute zero, the kinetic theory of gases, a number of works on the study of solutions formed the basis of the emerging physical chemistry, contributed to its formation into an independent science. The period of separation into a separate science lasted more than 100 years. The course of physical chemistry during this time was not read by any of the scientists.

One of the branches of physical chemistry, which has become an independent science, is colloidal chemistry.

colloid chemistry is a science that studies the properties of heterogeneous highly dispersed systems and polymer solutions.

Culinary processes: coagulation of proteins (during the heat treatment of meat, fish, eggs, etc.), obtaining stable emulsions (many sauces), foams (whipping cream, proteins, mousses), aging of jellies (hardening of bread, separation of liquid from jelly, jelly, etc.), adsorption (clarification of broths) - refer to colloidal processes. They are at the heart of all food production.

The laws of physical and colloidal chemistry underlie environmental protection measures. Usually, wastewater, factory chimney smoke - also colloidal systems. Methods for the destruction of these colloidal systems are based on the laws of physical colloidal chemistry.

Section 1. Physical chemistry

1. 1 Mainconcepts and laws of thermodynamics. Termaboutchemistry

1.1.1 Basic concepts of thermodynamics

Thermodynamics- a science that studies the general laws of the mutual transformation of energy from one form to another.

Chemical thermodynamics quantifies the thermal effects of various processes, clarifies the fundamental possibility of a spontaneous flow of chemical reactions and the conditions under which chemical reactions can be in a state of equilibrium.

The object of study in thermodynamics is system- a body or a group of bodies, actually or mentally separated from the environment. A system can be called a mineral crystal, a solution of any substance in a container, a gas in a cylinder, etc.

The system is called thermodynamic, if between the bodies that make up it, there can be an exchange of heat, matter, and if the system is described completely by thermodynamic parameters.

Types of systems (depending on the nature of interaction with the environment)

open	Closed	Isolated
It exchanges energy and matter with the environment.	It cannot exchange matter with the environment, but it can exchange energy and work with it.	It does not exchange matter and energy with the environment. Heat transfer, mutual transformations of energy, and concentration equalization can occur inside the system, but the internal energy of the system remains constant.
An open flask containing a solution from which the solvent can evaporate and which can be heated and cooled.	A tightly closed flask with a substance.	The reaction taking place in the thermostat.

The system can be homogeneous- consists of one phase (air, crystal, salt) and heterogeneous- consists of several phases (ice-water, water-benzene).

Phase- a part of a heterogeneous system separated by interfaces and characterized by the same physical properties at all its points.

Environment is everything that is in direct or indirect contact with the system. It is generally accepted that the environment has such a large size that the transfer or acquisition of heat by it does not change its temperature.

The state of a thermodynamic system is determined by mass, volume, pressure, composition, heat capacity, and other characteristics, which are called status parametersInia.

If the parameters of the state of the system do not change over time, then such a state is considered equilibrium. In an equilibrium thermodynamic system, the state parameters are interconnected by certain mathematical equations - the equations of state (for example, the Claiperon-Mendeleev equation for the state of an ideal gas).

Parameters that can be directly measured are called main parameters of the state. State parameters that cannot be directly measured (internal energy, enthalpy, entropy, thermodynamic potentials) are considered as functions of the main parameterstditch state.

thermodynamicallyeprocesss-changes in system state parameters:

isothermal (T=const);

· isobaric (Р=const);

Isochoric (V=const).

All bodies in nature, regardless of the state of aggregation, have a certain reserve internal energy.

Energy is made up of the kinetic energy of molecules, including the energy of translational and rotational motion, the energy of motion of atoms into molecules, electrons in atoms, intranuclear energy, the energy of interaction of particles with each other, etc. The kinetic and potential energy of the body itself is not included in the internal energy. Internal energy is a state function. The absolute value of the internal energy cannot be determined, only the change in internal energy (U) can be measured. The change in internal energy does not depend on the transition path, but depends only on the initial and final states of the system.

Heat (Q)(or the thermal effect of the process) is a quantitative characteristic of the energy that the system receives (gives off) from the environment during this process. Heat is a form of energy transfer realized by changing the kinetic energy of the thermal (chaotic) motion of particles (atoms, molecules). If the process is accompanied by the transfer of energy from the environment to the system, it is called endothermic, otherwise - exothermic. Any exothermic reaction in the forward direction becomes endothermic if it goes in the opposite direction, and vice versa.

Work (A), performed by the system, is due to the interaction of the system with the external environment, as a result of which external forces are overcome, i.e. work is one of the forms of energy exchange with the environment and serves as a quantitative characteristic of the transferred energy, and the energy transfer is realized through the ordered (organized) movement of molecules under the action of a certain force.

1.1. 2 First law of thermodynamics

This is a universal law of nature, the law of conservation and transformation of energy, corresponding to the basic position of dialectical materialism about the eternity and indestructibility of motion. This law was first formulated in 1842 by the outstanding German physicist J. Meyer.

Energy does not disappear and does not arise from nothing, but only transforms from one form to another in strictly equivalent ratios.

Depending on the type of system, the first law of thermodynamics has different formulations.

For a closed system, this law of thermodynamics establishes a relationship between the heat received or released by the system in some process, the change in the internal energy of the system, and the work produced in this case.

In an isolated system, internal energyRgia is constant, i.e. U=0.

If heat Q is supplied to a closed system, then this energy is dissipatedaboutpouts to increase the internal energy of the system U and on committing siWiththeme of work A versus outsidewthem environmental forces:

Under isobaric-isothermal conditions in which living organisms function:

where: p - external pressure,

V - change in the volume of the system.

We substitute (1.2) into (1.1).

Qр = U+рV = (U end - U start) + (рV end - рV start) = (U end + рV end) - (U end + рV start) (1.3)

The sum of the internal energy of the system and the product of volume and pressure (U + pV) is called enthalpy (N) - thermodynamic function characterizing the energy state of the system under isobaric-isothermal conditions. In this way:

Enthalpy is the sum of all types of energy concentrated in a given system, including the mechanical energy of particles, which can manifest itself in the form of work during expansion. Chemical reactions and physico-chemical processes can proceed with the release and absorption of energy. They are divided into exothermic and endothermic.

Processes in which heat is released are called exothermicand, processes occurring with the absorption of heat, - endothermiceskim.

In exothermic processes, the enthalpy decreases (H con H start), therefore:

ДH = (H end - H start);

In endothermic processes, the enthalpy increases (H con H start), therefore:

ДH = (H end - H start) 0,

The enthalpy of a system depends on pressure, temperature, and the amount of substance.

Under isobaric-isothermal conditions, the amount of heat that is released or absorbed during a chemical reaction is characterized by a change in enthalpy and is called reaction enthalpy H. The change in the enthalpy of reaction, determined at standard conditions, is called the standard enthalpy of reaction and is denoted H 0.

Enthalpy of reaction, i.e. the thermal effect of the reaction depends only on the nature and state of the initial substances and final products and does not depend onandsieve from the way, by toaboutto which the reaction proceeds.

Standard conditions:

The amount of substance is 1 mol;

pressure 760 mm. rt. Art. or 101.325 kPa;

temperature 298 0 K or 25 0 C.

1.1. 3 Thermochemistry

Chemical the equation, which indicates the value of the enthalpy (or thermal effect) of the reaction, is called thermochemical.

Thermochemical equations are used in thermochemistry. Thermochemistry determines the thermal effects of a chemical reaction and transitions from one state to another. The thermochemical equation differs from the chemical one in that the thermochemical equations indicate the absolute value and the sign of the thermal effect of the reaction, which is related to one mole of the starting or obtained substance, therefore, the stoichiometric coefficients in thermochemical equations can be fractional. In thermochemical equations, the state of aggregation and the crystalline form are also noted.

The enthalpy of reaction can be determined both experimentally and by calculation using the enthalpies of formation of substances involved in a chemical reaction based on Hess' law(1840):

In thermochemical calculations great importance have consequences from Hess's law:

1 consequence. The enthalpy of the reaction is equal to the difference between the algebraic sum of the enthalpies of formation of products and initial substances, taking into account the stoichiometric coefficients in the reaction equation.

2 consequence. The enthalpy of the direct realization is numerically equal to the enthalpy of the reverse reaction, but with the opposite sign.

1.1. 4 Second law of thermodynamics

This has the following formulations:

The transfer of heat from a cold body to a hot one is associated with compensation, i.e. with the need for additional work, which ultimately turns into heat absorbed by a hot body (for example, in a home refrigerator, heat is transferred from objects to parts of the device, and then to air. This transfer requires the expenditure of electricity). The processes, the implementation of which is associated with compensation, are called irreversibleandmy.

Spontaneous (natural, spontaneous) transition of energy (in the form of heat) from a less heated body to a more heated one is impossible.eto that.

The heat of the ocean, for example, can in principle be converted into work (according to the first law of thermodynamics), but only in the presence of an appropriate refrigerator (according to the second law of thermodynamics).

It is impossible to create a perpetual motion machine of the 2nd kind.

With regard to chemical reactions (at P, T=const), this position is expressed by the following mathematical equation:

H = G + TS or G = H - TS, (1.5)

where H is the thermal effect of the reaction observed during its irreversible flow;

G - change Gibbs free energy(free energy at constant pressure), or a change in the isobaric-isothermal potential, that is, this is the maximum part of the energy of the system that, under given conditions, can turn into useful work. At G 0 the reaction proceeds spontaneously.

Even with a reversible flow of the reaction, only part of the heat of the process can go into work. The other part, not converted into pabot, is transmitted at the same time from more heated to colder parts of the systemewe.

The function S introduced into equation (1.5) is called enteraboutFDI.

Entropy is a function of each specific, stationary state and does not depend on the path to reaching a new state (for example, on what intermediate stages the system goes through when moving from state 1 to state 2).

The product TS is the transferred heat (Q), which cannot be converted into work even with a reversible course of the reaction (the value of "bound energy"). This product shows the amount of internal energy lost in the form of heat:

TS = Q, or S = Q/T, (1.6)

The change in the entropy of the system during the reaction, equal to the heat imparted to the system, divided by the absolute temperature at which the system receives (gives off) this heat.

In addition to the thermodynamic potential - the Gibbs free energy G, in thermodynamics, as an auxiliary function for describing processes, another introduced thermodynamic potential is also of great importance - free energy Helmholtz F(free energy at constant volume), or isochoric-isothermal potential:

F = U - TS (for V, T=const) (1.7)

Spontaneous processes can produce work. Equilibrium occurs when this possibility is exhausted. Since negative changes in F and G correspond to spontaneous processes, the sign of the change in function G (at P, T=const) or function F (at V, T=const) will show the possibility or impossibility of a spontaneous reaction. If the changes in these functions for system states 1 and 2 are zero, then the system is in equilibrium.

Entropy differs from other system state parameters (P, T, V) in that its numerical value and the value of its change cannot be directly measured and can only be obtained indirectly, by calculation. To calculate the entropy S of the reaction aA + bB = cC = dD, it is necessary to subtract the sum of the entropies of the substances on the left side of the equation from the sum of the entropies of the substances on the right side of the equation (taking into account the stoichiometric coefficients). So, for standard conditions:

S 0 298K = S 0 298K (products) - S 0 298K (reagents), (1.8)

Only those processes that are associated with an increase in entropy can occur spontaneously in an isolated system, i.e. the system passes from a less probable state to a more probable one and reaches such a macroscopic state, which corresponds to a small number of microscopic states. In other words, processes are spontaneous when the final state can be realized a large number microstates and entropy is a measure of the system's striving for equilibrium. Such processes must be accompanied by an increase in entropy.

Questions for self-control:

1. What fundamental questions does chemical thermodynamics solve?

2. What is called a system, a thermodynamic system?

3. What are called state parameters? What are the state options?

4. What is called a thermodynamic process?

5. How is the first law of thermodynamics formulated?

6. What is the ratio of enthalpy to the internal energy of the system?

7. What is the standard enthalpy of formation?

8. What is different chemical equations from thermochemical?

9. What determines the second law of thermodynamics?

10. What do you need to know in order to determine the fundamental possibility of a particular reaction under given conditions?

11. What thermodynamic factors determine the direction of chemical reactions?

12. How do isobaric-isothermal and isochoric-isothermal potentials change in a spontaneously ongoing process?

1. 2 Aggregate states of substances, their characteristics

Depending on the external conditions(temperature and pressure), each substance can be in one of three states of aggregation: hard, livingdcom or gaseous.These states are called aggregate states.For some substances, only two or even one state of aggregation is characteristic. For example, naphthalene, iodine, when heated under normal conditions, from a solid state into a gaseous state, bypassing the liquid state. Substances such as proteins, starch, rubbers, which have huge macromolecules, cannot exist in a gaseous state.

Gases do not have a constant shape and constant volume. Liquids have a constant volume but do not have a constant shape. Solids are characterized by constancy of shape and volume.

1.2. 1 Character of the gaseous state of matter

For gases, the following properties are:

Uniform filling of the entire provided volume;

Low density compared to liquid and solids and high diffusion rate;

Relatively easy compressibility.

These properties are determined by the forces of intermolecular attraction and the distance between molecules.

In a gas, the molecules are at a very large distance from each other, the forces of attraction between them are negligible. At low pressures, the distances between gas molecules are so large that, compared with them, the size of molecules, and, consequently, the volume of molecules in the total volume of gas, can be neglected. At large distances between molecules, there are practically no forces of attraction between them. A gas in this state is called perfect.Under normal conditions T \u003d 273 0 K (0 0 C) and p \u003d 101.325 kPa, real gases, regardless of nature, can be considered ideal and applied to them the equation isIideal gafor (Claiperon equation-Mendeleev):

where P is the gas pressure,

V is the volume of gas,

The amount of substance

R - universal gas constant (in SI units R \u003d 8.314 J / molK),

T is the absolute temperature.

Real gases at high pressures and low temperatures do not obey the equation of state of an ideal gas, since under these conditions interaction forces between molecules begin to manifest themselves and it is no longer possible to neglect the intrinsic volume of molecules compared to the volume of the body. To mathematically describe the behavior of real gases, the equation is used van der Waals:

(p + n 2 a/V 2) (V - nb) = vRT, (2.2)

where a and b are constants,

a / V 2 - correction for mutual attraction,

b is the correction for the intrinsic volume of the molecules,

n is the number of moles of gas.

With an increase in pressure and a decrease in temperature, the distances between molecules decrease, and the interaction forces increase so that a substance can go from a gaseous state to a liquid one. For every gas there is a limit critical temperature, above which a gas cannot be liquefied at any pressure. The pressure required to liquefy a gas at a critical temperature is called critical pressure, and the volume of one mole of gas under these conditions critical volumeemom.

Rice. 1. Real gas isotherms

The state of the gas at critical parameters is called critical withaboutstanding.In the critical state, the difference between liquid and gas disappears, they have the same physical properties.

The transition of a gas into a liquid can be shown graphically. Figure 1 shows a graphical relationship between volume and pressure at constant temperatures. Such curves are called fromaboutterms. The isotherms can be divided into three sections: AB, BC, CD at low temperatures. AB - corresponds to the gaseous state, BC - corresponds to the transition of gas into liquid, CD - characterizes the liquid state. As the temperature rises, the section BC decreases and turns into an inflection point K, called critical point.

Liquefied gases find a great industrial application. Liquid CO 2 is used for carbonating fruit and mineral waters, making sparkling wines. Liquid SO 2 is used as a disinfectant to kill molds in cellars, cellars, wine barrels, fermentation tanks. Liquid nitrogen is widely used in medicine and biology to obtain low temperatures during canning and freezing of blood and biological tissues. Liquid gases are more convenient to transport.

1.2. 2 Characteristics of the liquid state of matter

Unlike gases, rather large forces of mutual attraction act between liquid molecules, which determines the peculiar nature of molecular motion. The thermal motion of a liquid molecule includes oscillatory and translational motions. Each molecule oscillates around a certain equilibrium point for some time, then moves and again occupies a new equilibrium position. This determines its fluidity. The forces of intermolecular attraction do not allow molecules to move far from each other during their movement. The total effect of the attraction of molecules can be represented as the internal pressure of liquids, which reaches very high values. This explains the constancy of volume and the practical incompressibility of liquids, although they easily take any form.

The properties of liquids also depend on the volume of molecules, their shape and polarity. If the liquid molecules are polar, then two or more molecules combine (associate) into a complex complex. Such liquids are called associateaboutbathrooms liquids. Associated liquids (water, acetone, alcohols) have higher boiling points, lower volatility, and higher dielectric constant. For example, ethyl alcohol and dimethyl ether have the same molecular formula (C 2 H 6 O). Alcohol is an associated liquid and boils at a higher temperature than dimethyl ether, which is a non-associated liquid.

The liquid state is characterized by such physical properties as plotness, viscosity, surface tension.

Surface tension.

The state of the molecules in the surface layer differs significantly from the state of the molecules in the depth of the liquid. Consider a simple case - liquid - vapor (Fig. 2).

Rice. 2. Action of intermolecular forces on the interface and inside the liquid

On fig. 2, the molecule (a) is inside the liquid, the molecule (b) is in the surface layer. The spheres around them are the distances over which the forces of intermolecular attraction of the surrounding molecules extend.

The molecule (a) is uniformly affected by intermolecular forces from the surrounding molecules, so the forces of intermolecular interaction are compensated, the resultant of these forces is equal to zero (f=0).

The density of a vapor is much less than the density of a liquid, since the molecules are far apart from each other. Therefore, the molecules in the surface layer almost do not experience the force of attraction from these molecules. The resultant of all these forces will be directed inside the liquid perpendicular to its surface. Thus, the surface molecules of a liquid are always under the influence of a force that tends to draw them in and, thereby, reduce the surface of the liquid.

To increase the liquid interface, it is necessary to expend work A (J). The work required to increase the interface S by 1 m 2 is a measure of the surface energy or surface tension.

Thus, the surface tension d (J / m 2 \u003d Nm / m 2 \u003d N / m) is the result of uncompensated intermolecular forces in the surface layer:

q = F/S (F - surface energy) (2.3)

There are many methods for determining surface tension. The most common are the stalagmometric method (the method of counting drops) and the method of the highest pressure of gas bubbles.

Using the methods of X-ray diffraction analysis, it was found that in liquids there is some orderliness in the spatial arrangement of molecules in individual microvolumes. Near each molecule, the so-called short-range order is observed. At some distance from it, this regularity is violated. And in the entire volume of the liquid there is no order in the arrangement of particles.

Rice. 3. Stalagmometer 4. Viscometer

Viscosity h (Pa s) - the property to resist the movement of one part of the liquid relative to the other. AT practical life a person is faced with a large variety of liquid systems, the viscosity of which is different - water, milk, vegetable oils, sour cream, honey, juices, molasses, etc.

The viscosity of liquids is due to intermolecular effects that limit the mobility of molecules. It depends on the nature of the liquid, temperature, pressure.

Viscosity is measured by devices called viscometers. The choice of viscometer and method for determining the viscosity depends on the state of the system under study and its concentration.

For liquids with a low viscosity or low concentration, capillary-type viscometers are widely used.

1.2. 3 Characteristics of the solid state of matter

Solids, unlike liquids and gases, retain their shape. The attractive forces between the particles that make up a solid body are so great that they cannot move freely relative to each other, but only oscillate around some middle position.

All solids are divided into crystalline and amorphous.In crystalline bodies, the particles are arranged in a certain order characteristic of each substance, and this order extends to the entire volume. Throughout the volume of an amorphous body, there is no order in the arrangement of particles. In this respect, amorphous bodies can be considered as liquids with an abnormally high viscosity.

Very often, amorphous and crystalline forms are different states of the same substance. So, silicon dioxide is found in nature both in the form of quartz crystals (rock crystal), and in an amorphous form - the mineral flint. Known crystalline and amorphous carbon.

The crystalline form is the most stable, substances gradually pass from the amorphous state to the crystalline state. Under normal conditions, this process is very slow, an increase in temperature can speed it up. For example, sugar can be in crystalline (granulated sugar, lump sugar) and in amorphous (caramelized) states. Over time, caramel can crystallize, which is undesirable in the confectionery industry. kinetics adsorption dispersed colloidal

The order in the spatial arrangement of particles and crystalline bodies - crystal cell- determines the external signs of the crystalline state. These include: 1) a definite and pronounced melting point; 2) a certain geometric shape of single crystals; 3) anisotropy.

Questions for self-control:

Under what conditions do the properties of a real gas approach those of an ideal gas?

Is it possible to compress a real gas indefinitely?

What physical meaning constants in the equation of state of a real gas?

Is it possible, knowing the temperature and pressure, to determine the number of molecules per unit volume?

What causes low compressibility of liquids?

How does the formation of a hydrogen bond between molecules affect the properties of a liquid?

How can one explain the decrease in surface tension and viscosity with increasing temperature?

How can a crystalline solid be distinguished from an amorphous one?

What is the main difference in the structure of crystalline and amorphous bodies?

1. 3 Chemical kinetics and catalysis.Chemical equilibrium

1.3.1 The rate of a chemical reaction

Kinetics- the study of the rate and mechanism of chemical reactions.

The question of the rate of a chemical reaction is of great practical and theoretical importance. The course of biochemical processes in the body, physicochemical changes in food products during heat treatment, and the performance of factory equipment depend on the reaction rate.

The speed of chemical processes can be controlled by changing the conditions of their occurrence. In some cases, it is desirable to intensify the process in order to obtain more product per unit of time. Sometimes it is necessary to reduce the rate of a chemical reaction, for example, to slow down the oxidation of fats in foods. All these problems can be solved by applying the laws of chemical kinetics.

Speed ​​reaction- change in the concentration of reacting substances per unit of time.

where c is the change in the concentration of reactants,

t - time interval.

The dependence of high-speed chemical reactions on concentration is determined the law of mass action, open empirical way K.M. Guldberg and P. Waage in 1867.

For the reaction aA + bB = C

where: A and B are the concentrations of reactants,

a and b are the coefficients in the equation,

k - coefficient of proportionality, called the rate constant, depending on the nature of the reactants and temperature.

The rate of a chemical reaction is proportional to the product of the endnfractions of reactants taken in powers equal toaboutcoefficients in the rea equationtotions.

Reaction rate constant numerically equal to the reaction rate at concentrations of reactants equal to unity.

Factors affecting the rate of a chemical reaction:

the nature of the reactants;

the concentration of reactants;

· temperature;

pressure (for gases);

the area of ​​contact of the reactants;

the presence of a catalyst.

As the temperature rises, the speed of movement of molecules increases, and, consequently, the number of collisions between them per unit time.

The influence of temperature on the rate of a chemical reaction obeys the van't Hoff rule.

For every 10 degrees increase in temperature, thebMost of the reactions increase by 2-4 times.

The number showing how many times the rate of a given reaction increases when the temperature rises by 10 degrees is called temperature toaboutreaction factor. Mathematically, this dependence is expressed by the relation:

where - temperature coefficient reactions,

and 0 - reaction rates at initial (t 1) and final (t 2) temperatures;

t - temperature change t 2 - t 1.

Van't Hoff's rule is approximate and can be applied to reactions occurring at temperatures from 0 to 300 degrees and in a small temperature range. As the temperature rises, the temperature coefficient of the reaction rate decreases, approaching unity.

A more accurate dependence of the rate of a chemical reaction on temperature was experimentally established by Arrhenius:

where k is the reaction rate constant,

B and A are constants for this reaction.

1.3. 2 Catalysis and catalysts

Catalyst A substance that changes the rate of a chemical reaction but is not consumed. Catalysts are either accelerating or decelerating.

Catalysis- the phenomenon of changing the reaction rate in the presence of catalysts.

catalytic reactions- reactions proceeding with the participation of catalysts.

If the catalyst is one of the products of the reaction, then the reaction is called autocatalytic, and the phenomenon itself autocatalysis.

Inhibitor a catalyst that slows down the reaction.

An example of positive catalysts is water in the interaction of aluminum powder with iodine.

Enzymes-biological catalysts of protein nature.

Enzymes are present in all living cells. It is customary to divide enzymes into simple and complex, or one-component or two-component. Simple enzymes consist only of protein, complex enzymes consist of protein and a non-protein part, which is called coenzyme.

Enzymes are characterized by high catalytic activity and selectivity. In terms of catalytic activity, they are significantly superior to inorganic catalysts. For example, 1 mole of catalase at 0 degrees decomposes 200,000 moles of H 2 O 2 in one second, and 1 mole of platinum at 20 degrees decomposes from 10 to 80 moles of hydrogen peroxide in one second.

Such accelerations of the reaction are due to the fact that enzymes sharply reduce energy barriers in the reaction path. For example, the activation energy for the decomposition reaction of H 2 O 2 under the action of an iron (II) ion and catalase molecules, respectively, is 42 and 7.1 kJ / mol; for the hydrolysis of urea with acid and urease, respectively, 103 and 28 kJ/mol.

Enzymes are very specific compared to inorganic catalysts. For example, amylase, contained in saliva, easily and quickly breaks down starch, but does not catalyze the process of sugar breakdown. Urease is extremely effective in catalyzing the hydrolysis of urea, but has no effect on its derivatives. This feature of enzymes allows living organisms, having an appropriate set of enzymes, to actively respond to external influences. For example, it has been observed that in stressful situations our body is amazing. The fact is described when weak woman she lifted a car by the bumper and held it while the people who came to the rescue freed the child who fell under it; a person pursued by an angry animal easily overcomes obstacles that are insurmountable for him in his usual state; in important competitions, athletes lose several kilograms in weight during the period of performance.

All that has been said about the remarkable properties of enzymes is explained by the fact that the selectivity of action (selectivity) and activity are interrelated: the higher the selectivity, the higher its activity. Enzymes have a unique selectivity, and therefore their activity is the highest.

1.3. 3 Chemical equilibrium

Reversible reactions can go in two opposite directions. They do not reach the end, but end with the establishment of chemical equilibrium.

Chemical equilibrium The state of a system when the rates of the forward and reverse reactions become equal.

The state of chemical equilibrium is maintained until conditions change. When external conditions change, the equilibrium is disturbed, and after a while the system will come to a new state of equilibrium.

Balance shift the transition of a system from one state of equilibrium to another.

The direction of equilibrium shift is determined Le Chat principleeleagues.

If the equilibrium system is affected, then it is equal toeyou are shifting thistin the direction that weakens this effect.

For example, an increase in temperature shifts the equilibrium towards an endothermic reaction, an increase in the concentration of the starting substances shifts the equilibrium towards the products of the reaction. Pressure only changes the equilibrium of reactions involving gases. An increase in pressure shifts the equilibrium in the direction of a reaction proceeding with a change in volume.

Questions for Samokontroll:

1. What does kinetics study?

2. What is called the rate of chemical reactions?

3. Why is there a minus sign in the mathematical equation for the rate of a chemical reaction?

4. List the factors that affect the rate of a chemical reaction.

5. Describe the effect of concentration, temperature, nature of reactants on the rate of a chemical reaction.

6. What is called catalysis and catalyst?

7. How are catalytic reactions classified?

8. What are inhibitors?

9. What is called chemical equilibrium?

10. What is called a shift in chemical equilibrium?

11. Formulate Le Chatelier's principle.

12. In which direction will the equilibrium of the equilibrium reaction shift with increasing temperature? Pressure (if gases are involved in the reactions)? The concentration of one of the reactants?

1. 4 Solution properties

1.4. 1 General characteristics of solutions

Solutions are of great importance in human life and practical activities. So, the processes of assimilation of food by humans and animals are associated with the transfer of nutrients into solution. Solutions are all the most important physiological fluids (blood, lymph, etc.). Industries based on chemical processes are usually associated with the use of solutions.

Solutions- multicomponent homogeneous systems in which one or more substances are distributed in the form of molecules, atoms or ions in the medium of another substance - a solvent.

The solution can have any state of aggregation - solid, liquid or gaseous. Any solution consists of solutes and a solvent. Usually, a solvent is considered to be a component that exists in its pure form in the same state of aggregation as the resulting solution (for example, a solution of salt in water: salt is a solute, water is a solvent). If both components before dissolution were in the same state of aggregation (for example, alcohol and water), then the component in a larger amount is considered the solvent.

In terms of structure, solutions occupy an intermediate position between mechanical mixtures and chemical compounds. With mechanical mixtures, they have in common the variability of composition, and with chemical compounds - the uniformity of the composition throughout the phase and the presence of a thermal effect during formation. In accordance with this, at first there were two opposing theories: "physical" and "chemical", each of which defended its own views on the structure of solutions.

Modern ideas about the structure of solutions are based on the solvation theory put forward by Mendeleev and developed by his followers. According to this theory, two processes simultaneously occur in the system during dissolution: diffusion of the solute in the volume of the solvent (a physical process) and the formation of unstable compounds of variable composition - solvates ( chemical process). If the solvent is water, then these compounds are called hydrates.

The formation of solutions is a spontaneous process that proceeds with an increase in the disorder of the system, i.e. with an increase in entropy. For example, when a crystal is dissolved, the system goes from a completely ordered state to a less ordered one. In this case, with an increase in entropy (AS > 0), the free energy of the system decreases (AG<0).

If the solution is formed from 2 liquids, then the driving force of the dissolution process is due to the tendency of the components of the solution to equalize the concentrations, which also leads to an increase in entropy, i.e. AS > 0 and AQ< 0. Растворение вещества - процесс обратимый. И как всякий обратный процесс, растворение заканчивается установлением динамического равновесия: нерастворенное вещество - вещество в растворе. Раствор, находящийся в равновесии с растворяющимся веществом, называют насыщенным раствором, а достигнутую предельную концентрацию насыщенного раствора - растворимостью.

The most important characteristic of a solution is its composition or concentration of components.

Solution concentration- the amount of a solute contained in a certain amount of a solution or solvent.

The concentration of solutions can be expressed in different ways. In chemical practice, the following methods of expressing concentrations are most commonly used:

1. Mass fraction of a dissolved substance (percentage concentration)- shows how many grams of a substance are dissolved in 100 g of a solution. It is determined by the formula:

where W is the mass fraction of the solute,

m in-va - the mass of the dissolved substance,

m solution - the mass of the solution.

2. Molar concentration- shows how many moles of a solute are contained in 1 liter of solution.

3. Molar concentration- shows how many moles of a substance are contained in 1 kg of solvent.

1.4. 2 Solutions of gases in liquids

The solubility of gases in liquids depends on their nature, the nature of the solvent, temperature and pressure. As a rule, the solubility of a gas is greater if the dissolution is accompanied by chemical interaction with the solvent, and less if there is no chemical interaction. For example, in 1 liter of water at n.o. dissolves 0.0002 g of hydrogen, which does not interact with water, and 875 g of ammonia, which reacts with water to form ammonium hydroxide.

The dependence of the solubility of gases on the nature of the solvent can be shown in the following examples. Under the same conditions, 87.5 g of NH 3 is dissolved in 1000 g of water, and only 25 g is dissolved in 100 g of ethyl alcohol. The solubility of gases largely depends on temperature. As the temperature rises, their solubility decreases, and as the temperature decreases, it increases. So at 0 0 C, 171 cm 3 CO 2 dissolves in 100 ml of water, at 20 0 C - only 87.8 cm 3. Therefore, prolonged boiling can almost completely remove dissolved gases from a liquid, and it is advisable to saturate liquids with gas at low temperatures.

The solubility of a gas also depends on pressure. The dependence of gas solubility on pressure is determined ge lawnri.

C = k p, (4.2)

where C is the gas concentration in the solution,

k - coefficient of proportionality, depending on the nature of the liquid and gas,

p is the pressure of the gas above the solution.

The mass of dissolved gas at constant temperature is directly praboutis proportional to the gas pressure over the solutionaboutrum.

Henry's law is valid only for dilute solutions at low pressures. Gases that interact with the solvent NH 3 , SO 2 , HC1 with water do not obey Henry's law. Their solubility also increases with increasing pressure, but according to a more complex law.

The manifestation of Henry's law is illustrated by the formation of copious foam when uncorking a bottle of soda water or a bottle of champagne; here there is a sharp decrease in the solubility of the gas with a decrease in its partial pressure. The same law explains the occurrence of decompression sickness. At a depth of 40 m below sea level, the total pressure is 600 kPa and the solubility of nitrogen in blood plasma is 9 times greater than on the sea surface. When a diver rises rapidly from a depth, dissolved nitrogen is released into the blood in bubbles that clog blood vessels, which can lead to serious consequences.

The solubility of a gas decreases when a third component is present in the solution. Thus, gases dissolve much worse in electrolyte solutions than in pure water. For example, 3 10 3 m 3 of chlorine dissolves in 1 g of water at 0 0 C, and 10 times less dissolves in 1 g of a saturated NaCl solution, therefore, when chlorine is stored above a liquid, water is replaced with a sodium chloride solution.

1.4. 3 Mutual solubility of liquids

Unlike the solubility of gases in liquids, the dissolution of a liquid is a more complex process. When two liquids are mixed, they can:

Dissolve in each other in any ratio;

Practically insoluble;

Dissolve limited.

The mutual solubility of liquids depends primarily on their chemical structure. Even the alchemists noticed that "like dissolves in like", i.e. the polar is usually soluble in the polar, and the non-polar in the non-polar. For this reason, water (polar liquid) is a good solvent for polar liquids (ethyl alcohol, acetic acid, etc.) and does not dissolve non-polar liquids (benzene, kerosene, etc.) at all. If the liquids differ from each other in polarity, then they are limitedly soluble in each other. With limited solubility, each of the liquids passes into the other up to a certain limit, resulting in a two-layer system. For example, with an increase in temperature, their mutual solubility usually increases, and at a certain temperature both liquids mix in any ratio, and the boundary between them disappears. This temperature is called critical.

The critical temperature reached by heating is called upper critical temperature.

Known mixtures of liquids, where the solubility decreases with increasing temperature. Therefore, the critical temperature is reached as the temperature decreases and is called lower critical temperatureatRoy.

Using the critical dissolution temperature, some analytical determinations are sometimes carried out.

Of particular interest is the solubility of various substances in two-layer systems consisting of two insoluble liquids.

If a third substance capable of solubility in each of them is introduced into a system consisting of two immiscible liquids, then the solute will be distributed between both liquids in proportion to its solubility in each of them.

The ratio of the concentrations of a substance distributed between two immiscible liquids at a constant temperature remains constant, regardless of the total amount of solute.

С 1 /С 2 = k, (4.3)

where C 1 and C 2 are the concentration of the solute in the 1st and 2nd solvents,

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Chemical thermodynamics. Basic concepts of thermodynamics. First law of thermodynamics. Applications of the first law of thermodynamics to chemical processes. The dependence of the thermal effect of the reaction on temperature. Kirchhoff's law. The second law of thermodynamics.

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

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

The first book details the following sections of the course: quantum mechanical foundations of the theory chemical bond, the structure of atoms and molecules, spectral methods for studying the 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 ideas 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.

Format: djvu

The size: 11.2 MB

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Book 2.

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The size: 7 MB

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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 tasks. 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. Saturability 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 heat capacity of a crystalline substance 211
§ 7.7. Quantum statistical theory of heat capacity 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 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 body 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 chart 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 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
Application 305
Index 312
Contents 316

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. Yakobi 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. K. 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 late XIX in. 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 by 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 strong theoretical background. Here it is necessary to note the major studies of N. S. Kurnakov on physical and chemical analysis, work in the field of electrochemistry by A. N. Frumkin, the creation of the theory of chain reactions by N. N. Semenov, 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 basis 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 heterogeneous system is a system consisting of various parts that differ in structure or chemical composition . 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, multiples and submultiple units pressure: 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