• Domestic science and medicine in the 19th - early 20th centuries Chemistry of the 19th century in medicine
• What are peptides? Types and types of peptides. All about peptides and anti-aging cosmetics with peptides Additional peptides are not formed
• Toxic effect on humans of hazardous chemicals
• Radioautography Method of autoradiography in cytology
• Identification of bacteria by antigenic structure
• Organic matter in wastewater What are organic compounds in water
• Hydrogen index (pH factor)
• What is the pH of water and why is it important to know it?
• Redox potential Redox systems
• Reversible redox system Reversible redox systems

1. The rate of chemical reactions. Concept definition. Factors affecting the rate of a chemical reaction: reagent concentration, pressure, temperature, presence of a catalyst. The law of mass action (LMA) as the basic law of chemical kinetics. The rate constant, its physical meaning. Influence on the reaction rate constant of the nature of the reactants, temperature and the presence of a catalyst.

The rate of a homogeneous reaction is a quantity numerically equal to the change in the molar concentration of any participant in the reaction per unit time.

The average reaction rate v cf in the time interval from t 1 to t 2 is determined by the ratio:

The main factors affecting the rate of a homogeneous chemical reaction are:

• - the nature of the reactants;
• - molar concentrations of reagents;
• - pressure (if gases are involved in the reaction);
• - temperature;
• - the presence of a catalyst.

The rate of a heterogeneous reaction is a value numerically equal to the change in the chemical amount of any participant in the reaction per unit time per unit area of ​​the interface: .

By stages, chemical reactions are divided into simple (elementary) and complex. Most chemical reactions are complex processes that occur in several stages, i.e. consisting of several elementary processes.

For elementary reactions, the law of mass action is valid: the rate of an elementary chemical reaction is directly proportional to the product of the concentrations of the reactants in powers equal to the stoichiometric coefficients in the reaction equation.

For an elementary reaction aA + bB > ... the reaction rate, according to the law of mass action, is expressed by the relation:

where c(A) and c(B) are the molar concentrations of reactants A and B; a and b are the corresponding stoichiometric coefficients; k is the rate constant of this reaction.

For heterogeneous reactions, the equation of the law of mass action does not include the concentrations of all reagents, but only gaseous or dissolved ones. So, for the combustion reaction of carbon:

C (c) + O 2 (g) > CO 2 (g)

the velocity equation has the form: .

The physical meaning of the rate constant is that it is numerically equal to the rate of a chemical reaction at concentrations of reactants equal to 1 mol/dm 3 .

The value of the rate constant of a homogeneous reaction depends on the nature of the reactants, temperature and catalyst.

2. Effect of temperature on the rate of a chemical reaction. Temperature coefficient of the rate of a chemical reaction. active molecules. Distribution curve of molecules according to their kinetic energy. Activation energy. Ratio of activation energy and chemical bond energy in initial molecules. Transition state, or activated complex. Activation energy and thermal effect of the reaction (energy scheme). Dependence of the temperature coefficient of the reaction rate on the value of the activation energy.

As the temperature increases, the rate of a chemical reaction usually increases. The value showing how many times the reaction rate increases with an increase in temperature by 10 degrees (or, what is the same, by 10 K), is called the temperature coefficient of the chemical reaction rate (r):

where - the values ​​of the reaction rate, respectively, at temperatures T 2 and T 1; r is the temperature coefficient of the reaction rate.

The dependence of the reaction rate on temperature is approximately determined by the van't Hoff empirical rule: with an increase in temperature for every 10 degrees, the rate of a chemical reaction increases by 2–4 times.

A more accurate description of the dependence of the reaction rate on temperature is feasible within the framework of the Arrhenius activation theory. According to this theory, a chemical reaction can only occur when active particles collide. Particles are called active if they have a certain energy characteristic of a given reaction, which is necessary to overcome the repulsive forces that arise between the electron shells of the reacting particles. The proportion of active particles increases with increasing temperature.

An activated complex is an intermediate unstable group that is formed during the collision of active particles and is in a state of redistribution of bonds. When the activated complex decomposes, reaction products are formed.

The activation energy E and is equal to the difference between the average energy of the reacting particles and the energy of the activated complex.

For most chemical reactions, the activation energy is less than the dissociation energy of the weakest bonds in the molecules of the reactants.

In activation theory, the effect of temperature on the rate of a chemical reaction is described by the Arrhenius equation for the rate constant of a chemical reaction:

where A is a constant factor, independent of temperature, determined by the nature of the reactants; e is the base of the natural logarithm; E a - activation energy; R is the molar gas constant.

As follows from the Arrhenius equation, the higher the rate constant of the reaction, the lower the activation energy. Even a slight decrease in the activation energy (for example, when a catalyst is introduced) leads to a noticeable increase in the reaction rate.

According to the Arrhenius equation, an increase in temperature leads to an increase in the rate constant of a chemical reaction. The smaller the value of E a, the more noticeable the effect of temperature on the reaction rate and, therefore, the greater the temperature coefficient of the reaction rate.

3. Influence of a catalyst on the rate of a chemical reaction. Homogeneous and heterogeneous catalysis. Elements of the theory of homogeneous catalysis. Theory of intermediate compounds. Elements of the theory of heterogeneous catalysis. Active centers and their role in heterogeneous catalysis. The concept of adsorption. Influence of a catalyst on the activation energy of a chemical reaction. Catalysis in nature, industry, technology. biochemical catalysis. Enzymes.

Catalysis is a change in the rate of a chemical reaction under the action of substances whose quantity and nature after the completion of the reaction remain the same as before the reaction.

A catalyst is a substance that changes the rate of a chemical reaction, but remains chemically unchanged.

A positive catalyst speeds up the reaction; a negative catalyst, or inhibitor, slows down the reaction.

In most cases, the effect of a catalyst is explained by the fact that it reduces the activation energy of the reaction. Each of the intermediate processes involving a catalyst proceeds with a lower activation energy than the non-catalyzed reaction.

In homogeneous catalysis, the catalyst and reactants form one phase (solution). In heterogeneous catalysis, the catalyst (usually a solid) and the reactants are in different phases.

In the course of homogeneous catalysis, the catalyst forms an intermediate compound with the reagent, which reacts with the second reagent at a high rate or rapidly decomposes with the release of the reaction product.

An example of homogeneous catalysis: the oxidation of sulfur oxide (IV) to sulfur oxide (VI) with oxygen in the nitrous method for producing sulfuric acid (here the catalyst is nitrogen oxide (II), which easily reacts with oxygen).

In heterogeneous catalysis, the reaction proceeds on the surface of the catalyst. The initial stages are the diffusion of reactant particles to the catalyst and their adsorption (i.e. absorption) by the catalyst surface. Reagent molecules interact with atoms or groups of atoms located on the surfaces of the catalyst, forming intermediate surface compounds. The redistribution of the electron density that occurs in such intermediate compounds leads to the formation of new substances that are desorbed, i.e., removed from the surface.

The process of formation of intermediate surface compounds occurs at the active centers of the catalyst.

An example of heterogeneous catalysis is an increase in the rate of oxidation of sulfur(IV) oxide to sulfur(VI) oxide with oxygen in the presence of vanadium(V) oxide.

Examples of catalytic processes in industry and technology: the synthesis of ammonia, the synthesis of nitric and sulfuric acids, the cracking and reforming of oil, the afterburning of products of incomplete combustion of gasoline in cars, etc.

Examples of catalytic processes in nature are numerous, since most of the biochemical reactions occurring in living organisms are catalytic reactions. These reactions are catalyzed by proteins called enzymes. There are about 30,000 enzymes in the human body, each of which catalyses only one type of process (for example, saliva ptyalin catalyzes only the conversion of starch to glucose).

4. Chemical balance. Reversible and irreversible chemical reactions. state of chemical equilibrium. Chemical equilibrium constant. Factors that determine the value of the equilibrium constant: the nature of the reactants and temperature. Shift in chemical equilibrium. Influence of changes in concentration, pressure and temperature on the position of chemical equilibrium.

Chemical reactions, as a result of which the starting substances are completely converted into reaction products, are called irreversible. Reactions going simultaneously in two opposite directions (forward and backward) are called reversible.

In reversible reactions, the state of the system in which the rates of the forward and reverse reactions are equal () is called the state of chemical equilibrium. Chemical equilibrium is dynamic, i.e. its establishment does not mean the termination of the reaction. In the general case, for any reversible reaction aA + bB - dD + eE, regardless of its mechanism, the following relation holds:

At equilibrium, the product of the concentrations of the reaction products, referred to the product of the concentrations of the starting materials, for a given reaction at a given temperature is a constant value, called the equilibrium constant (K).

The value of the equilibrium constant depends on the nature of the reactants and temperature, but does not depend on the concentrations of the components of the equilibrium mixture.

Changing the conditions (temperature, pressure, concentration), under which the system is in a state of chemical equilibrium (), causes an imbalance. As a result of unequal changes in the rates of direct and reverse reactions () over time, a new chemical equilibrium () is established in the system, corresponding to new conditions. The transition from one equilibrium state to another is called a shift, or a shift in the equilibrium position.

If, during the transition from one equilibrium state to another, the concentrations of substances recorded on the right side of the reaction equation increase, they say that the equilibrium shifts to the right. If, upon transition from one equilibrium state to another, the concentrations of substances recorded on the left side of the reaction equation increase, they say that the equilibrium shifts to the left.

The direction of the shift in chemical equilibrium as a result of changes in external conditions is determined by the Le Chatelier principle: If an external influence is exerted on a system in a state of chemical equilibrium (change the temperature, pressure or concentration of substances), then it will favor the flow of one of the two opposite processes, which weakens this effect.

According to Le Chatelier's principle:

An increase in the concentration of the component written on the left side of the equation leads to a shift in equilibrium to the right; an increase in the concentration of the component written on the right side of the equation leads to a shift in equilibrium to the left;

With an increase in temperature, the equilibrium shifts towards the occurrence of an endothermic reaction, and with a decrease in temperature, in the direction of an exothermic reaction;

• - With an increase in pressure, the equilibrium shifts towards a reaction that reduces the number of molecules of gaseous substances in the system, and with a decrease in pressure - towards a reaction that increases the number of molecules of gaseous substances.
• 5. Photochemical and chain reactions. Features of the course of photochemical reactions. Photochemical reactions and wildlife. Unbranched and branched chemical reactions (on the example of the reactions of the formation of hydrogen chloride and water from simple substances). Conditions for the initiation and termination of chains.

Photochemical reactions are reactions that take place under the influence of light. A photochemical reaction proceeds if the reagent absorbs radiation quanta, which are characterized by an energy that is quite specific for this reaction.

In the case of some photochemical reactions, by absorbing energy, the reactant molecules pass into an excited state, i.e. become active.

In other cases, a photochemical reaction proceeds if quanta of such high energy are absorbed that chemical bonds are broken and the molecules dissociate into atoms or groups of atoms.

The rate of the photochemical reaction is the greater, the greater the intensity of irradiation.

An example of a photochemical reaction in wildlife is photosynthesis, i.e. the formation of organic substances of cells due to the energy of light. In most organisms, photosynthesis takes place with the participation of chlorophyll; in the case of higher plants, photosynthesis is summarized by the equation:

CO 2 + H 2 O organic matter + O 2

Photochemical processes also underlie the functioning of vision processes.

A chain reaction is a reaction that is a chain of elementary acts of interaction, and the possibility of each act of interaction occurring depends on the success of the previous act.

The stages of a chain reaction are chain initiation, chain development, and chain termination.

The origin of the chain occurs when, due to an external source of energy (quantum of electromagnetic radiation, heating, electric discharge), active particles with unpaired electrons (atoms, free radicals) are formed.

In the course of chain development, the radicals interact with the initial molecules, and new radicals are formed in each act of interaction.

Chain termination occurs if two radicals collide and transfer the energy released in this case to a third body (a molecule resistant to decay, or a vessel wall). The chain can also be terminated if an inactive radical is formed.

There are two types of chain reactions - unbranched and branched.

In unbranched reactions, at the stage of chain development, one new radical is formed from each reacting radical.

In branched reactions at the stage of chain development, 2 or more new radicals are formed from one reacting radical.

6. Factors determining the direction of a chemical reaction. Elements of chemical thermodynamics. Concepts: phase, system, environment, macro- and microstates. Basic thermodynamic characteristics. The internal energy of the system and its change in the course of chemical transformations. Enthalpy. The ratio of enthalpy and internal energy of the system. The standard enthalpy of a substance. Enthalpy change in systems during chemical transformations. Thermal effect (enthalpy) of a chemical reaction. Exo- and endothermic processes. Thermochemistry. Hess' law. thermochemical calculations.

Thermodynamics studies the patterns of energy exchange between the system and the environment, the possibility, direction and limits of the spontaneous flow of chemical processes.

A thermodynamic system (or simply a system) is a body or a group of interacting bodies mentally distinguished in space. The rest of the space outside the system is called the environment (or simply the environment). The system is separated from the environment by a real or imaginary surface.

A homogeneous system consists of one phase, a heterogeneous system consists of two or more phases.

A phase is a part of a system that is homogeneous at all its points in terms of chemical composition and properties and is separated from other parts of the system by an interface.

The state of the system is characterized by the totality of its physical and chemical properties. The macrostate is determined by the averaged parameters of the entire set of particles in the system, and the microstate is determined by the parameters of each individual particle.

Independent variables that determine the macrostate of the system are called thermodynamic variables, or state parameters. Temperature T, pressure p, volume V, chemical quantity n, concentration c, etc. are usually chosen as state parameters.

A physical quantity, the value of which depends only on the state parameters and does not depend on the transition path to a given state, is called a state function. The state functions are, in particular:

U - internal energy;

H - enthalpy;

S - entropy;

G - Gibbs energy (free energy or isobaric-isothermal potential).

The internal energy of the system U is its total energy, consisting of the kinetic and potential energy of all particles of the system (molecules, atoms, nuclei, electrons) without taking into account the kinetic and potential energy of the system as a whole. Since a complete account of all these components is impossible, then in the thermodynamic study of a system, the change in its internal energy during the transition from one state (U 1) to another (U 2) is considered:

U 1 U 2 U = U 2 -U1

The change in the internal energy of the system can be determined experimentally.

The system can exchange energy (heat Q) with the environment and do work A, or, conversely, work can be done on the system. According to the first law of thermodynamics, which is a consequence of the law of conservation of energy, the heat received by the system can only be used to increase the internal energy of the system and to perform work by the system:

Q= U+A

In the future, we will consider the properties of such systems, which are not affected by any other forces, except for the forces of external pressure.

If the process in the system proceeds at a constant volume (i.e., there is no work against the forces of external pressure), then A \u003d 0. Then the thermal effect of the process proceeding at a constant volume, Q v is equal to the change in the internal energy of the system:

Most chemical reactions encountered in everyday life take place at constant pressure (isobaric processes). If no other forces act on the system, except for constant external pressure, then:

A = p(V2 - V 1 ) = pV

Therefore, in our case (p = const):

Qp=U + pV

Q p \u003d U 2 - U 1 + p(V 2 - V 1 ), where

Qp = (U 2 +pV 2 )-(U 1 +pV 1 ).

The function U + pV is called the enthalpy; it is denoted by the letter N. Enthalpy is a function of state and has the dimension of energy (J).

Qp= H 2 - H 1 =H,

i.e., the thermal effect of the reaction at constant pressure and temperature T is equal to the change in the enthalpy of the system during the reaction. It depends on the nature of the reactants and products, their physical state, the conditions (T, p) of the reaction, and also on the amount of substances involved in the reaction.

The enthalpy of a reaction is the change in the enthalpy of a system in which the reactants interact in amounts equal to the stoichiometric coefficients in the reaction equation.

The reaction enthalpy is called standard if the reactants and reaction products are in standard states.

The standard state of a substance is the aggregate state or crystalline form of a substance in which it is thermodynamically most stable under standard conditions (T \u003d 25 o C or 298 K; p \u003d 101.325 kPa).

The standard state of a substance that exists at 298 K in solid form is considered to be its pure crystal under a pressure of 101.325 kPa; in liquid form - pure liquid under pressure of 101.325 kPa; in gaseous form - a gas with its own pressure of 101.325 kPa.

For a solute, its state in solution at a molality of 1 mol/kg is considered standard, and it is assumed that the solution has the properties of an infinitely dilute solution.

The standard enthalpy of the reaction of formation of 1 mol of a given substance from simple substances in their standard states is called the standard enthalpy of formation of this substance.

Recording example: (CO 2) \u003d - 393.5 kJ / mol.

The standard enthalpy of formation of a simple substance that is in the most stable (for given p and T) state of aggregation is taken equal to 0. If an element forms several allotropic modifications, then only the most stable (for given p and T) modification has zero standard enthalpy of formation.

Usually, thermodynamic quantities are determined under standard conditions:

p \u003d 101.32 kPa and T \u003d 298 K (25 ° C).

Chemical equations that indicate changes in enthalpy (heat effects of reactions) are called thermochemical equations. There are two forms of writing thermochemical equations in the literature.

Thermodynamic form of the thermochemical equation:

C (graphite) + O 2 (g) CO 2 (g); = - 393.5 kJ.

The thermochemical form of the thermochemical equation for the same process:

C (graphite) + O 2 (g) CO 2 (g) + 393.5 kJ.

In thermodynamics, the thermal effects of processes are considered from the point of view of the system. Therefore, if the system releases heat, then Q< 0, а энтальпия системы уменьшается (ДH < 0).

In classical thermochemistry, thermal effects are considered from the standpoint of the environment. Therefore, if the system releases heat, then it is assumed that Q > 0.

An exothermic process is a process that proceeds with the release of heat (DH< 0).

Endothermic is a process that proceeds with the absorption of heat (DH> 0).

The basic law of thermochemistry is Hess's law: "The thermal effect of a reaction is determined only by the initial and final state of the system and does not depend on the path of the system's transition from one state to another."

Consequence from Hess' law: The standard thermal effect of a reaction is equal to the sum of the standard heats of formation of the reaction products minus the sum of the standard heats of formation of the starting materials, taking into account the stoichiometric coefficients:

• (reactions) = (cont.) -(out.)
• 7. The concept of entropy. Entropy change during phase transformations and chemical processes. The concept of the isobaric-isothermal potential of the system (Gibbs energy, free energy). The ratio between the magnitude of the change in the Gibbs energy and the magnitude of the change in the enthalpy and entropy of the reaction (basic thermodynamic relationship). Thermodynamic analysis of the possibility and conditions for the occurrence of chemical reactions. Features of the course of chemical processes in living organisms.

Entropy S is a value proportional to the logarithm of the number of equiprobable microstates (W) through which a given macrostate can be realized:

S=k ln W

The unit of entropy is J/mol?K.

Entropy is a quantitative measure of the degree of disorder in a system.

Entropy increases during the transition of a substance from a crystalline state to a liquid and from a liquid to a gaseous state, during the dissolution of crystals, during the expansion of gases, during chemical interactions leading to an increase in the number of particles, and above all particles in the gaseous state. On the contrary, all processes, as a result of which the ordering of the system increases (condensation, polymerization, compression, reduction in the number of particles), are accompanied by a decrease in entropy.

There are methods for calculating the absolute value of the entropy of a substance, therefore, in the tables of thermodynamic characteristics of individual substances, data are given for S 0, and not for DS 0.

The standard entropy of a simple substance, as opposed to the enthalpy of formation simple matter, is not equal to zero.

For entropy, a statement similar to that considered above for H is true: the change in the entropy of the system as a result of a chemical reaction (S) is equal to the sum of the entropies of the reaction products minus the sum of the entropies of the starting substances. As in the calculation of enthalpy, summation is carried out taking into account stoichiometric coefficients.

The direction in which a chemical reaction spontaneously proceeds in an isolated system is determined by the combined action of two factors: 1) the tendency for the system to transition to a state with the lowest internal energy (in the case of isobaric processes, with the lowest enthalpy); 2) the tendency to achieve the most probable state, i.e., the state that can be realized in the largest number of equiprobable ways (microstates), i.e.:

DH > min, DS > max.

The state function that simultaneously reflects the influence of both tendencies mentioned above on the direction of chemical processes is the Gibbs energy (free energy, or isobaric-isothermal potential), related to enthalpy and entropy by the relation

where T is the absolute temperature.

As you can see, the Gibbs energy has the same dimension as the enthalpy, and therefore is usually expressed in J or kJ.

For isobaric-isothermal processes (i.e., processes occurring at constant temperature and pressure), the change in Gibbs energy is:

G= H-TS

As in the case of H and S, the change in the Gibbs energy G as a result of a chemical reaction (the Gibbs energy of the reaction) is equal to the sum of the Gibbs energies of the formation of the reaction products minus the sum of the Gibbs energies of the formation of the initial substances; summation is carried out taking into account the number of moles of the substances involved in the reaction.

The Gibbs energy of formation of a substance is related to 1 mole of this substance and is usually expressed in kJ/mol; in this case, G 0 of the formation of the most stable modification of a simple substance is taken equal to zero.

At a constant temperature and pressure, chemical reactions can spontaneously proceed only in such a direction, in which the Gibbs energy of the system decreases (G0). This is a condition for the fundamental possibility of implementing this process.

The table below shows the possibility and conditions for the reaction to proceed with various combinations of H and S signs:

By the sign of G, one can judge the possibility (impossibility) of a spontaneous flow of a single process. If the system is affected, then it is possible to carry out a transition from one substance to another, characterized by an increase in free energy (G>0). For example, in the cells of living organisms reactions of formation of complex organic compounds proceed; the driving force of such processes are solar radiation and oxidation reactions in the cell.

Lecture 1 Chemical thermodynamics. Chemical kinetics and catalysis PLAN 1. Basic concepts of thermodynamics. 2. Thermochemistry. 3. Chemical balance. 4. The rate of chemical reactions. 5. Effect of temperature on the rate of reactions. 6. The phenomenon of catalysis. Prepared by: Ph.D., Assoc. Ivanets L.M., ass. Kozachok S.S. Lecturer Assistant of the Department of Pharmaceutical Chemistry Kozachok Solomeya Stepanovna

Thermodynamics - Thermodynamics is a branch of physics that studies the mutual transformations of various types of energy associated with the transfer of energy in the form of heat and work. The great practical significance of thermodynamics is that it makes it possible to calculate the thermal effects of a reaction, to indicate in advance the possibility or impossibility of a reaction, as well as the conditions for its passage.

Internal energy Internal energy is the kinetic energy of all particles of the system (molecules, atoms, electrons) and the potential energy of their interactions, except for the kinetic and potential energy of the system as a whole. Internal energy is a state function, i.e. its change is determined by the given initial and final states of the system and does not depend on the path of the process: U = U 2 - U 1

The first law of thermodynamics Energy does not disappear without a trace and does not arise from nothing, but only passes from one form to another in an equivalent amount. A perpetual motion machine of the first kind, that is, a periodically operating machine that gives work without wasting energy, is impossible. Q \u003d U + W In any isolated system, the total energy supply remains unchanged. Q=U+W

The thermal effect of a chemical reaction at constant V or p does not depend on the reaction path, but is determined by the nature and state of the starting materials and reaction products Hess's law H 1 H 2 H 3 H 4 4 H 1 \u003d H 2 + H 3 + H 4

The second law of thermodynamics, like the first, is the result of centuries of human experience. There are various formulations of the second law, but they all determine the direction of spontaneous processes: 1. Heat cannot spontaneously transfer from a cold body to a hot one (Clausius' postulate). 2. A process whose only result is the conversion of heat into work is impossible (Thomson's postulate). 3. It is impossible to build a batch machine that only cools the heat reservoir and performs work (Planck's first postulate). 4. Any form of energy can be completely converted into heat, but heat is only partially converted into other types of energy (Planck's second postulate).

Entropy is a thermodynamic function of state, therefore its change does not depend on the path of the process, but is determined only by the initial and final states of the system. then S 2 - S 1 = ΔS = S 2 - S 1 = ΔS = The physical meaning of entropy is the amount of bound energy, which is related to one degree: in isolated systems, the direction of the flow of spontaneous processes is determined by the change in entropy.

Characteristic functions U is a function of an isochoric-isoentropic process: dU = TdS – pdV. For an arbitrary process: U 0 H is a function of an isobaric isoentropic process: dH = TdS + Vdp For an arbitrary process: H 0 S is a function of an isolated system For an arbitrary process: S 0 For an arbitrary process: S 0 F is a function of an isochoric isothermal process dF = dU – TdS. For an arbitrary process: F 0 G is a function of an isobaric-isothermal process: dG = dH- TdS For an arbitrary process: G 0

Classification of chemical reactions according to the number of stages Simple ones proceed in one elementary chemical act Complex ones proceed in several stages Reverse reaction A B

The effect of temperature on the rate of reactions The effect of temperature on the rate of enzymatic reactions t t

Comparison of van't Hoff: Calculation of the shelf life of medicines according to the "accelerated aging" method of van't Hoff: at t 2 t 1 Temperature coefficient of speed:

Basic concepts and laws of chemistry. Chemical bond. The structure and properties of matter

1. What substances are called simple? Complex? From the given substances, select simple ones: CO, O 3, CaO, K, H 2, H 2 O.

2. What substances are called oxides? Acids? Reasons? Salts?

3. From the given oxides - SO 2, CaO, ZnO, Cr 2 O 3, CrO, P 2 O 5, CO 2, Cl 2 O 3, Al 2 O 3 - select basic, acidic and amphoteric.

4. What salts are classified as acidic, basic, medium, double, mixed, complex?

5. Name the following compounds: ZnOHCl, KHSO 3 , NaAl(SO 4) 2 . What class of compounds do they belong to?

6. What is called the basicity of an acid?

7. From the given hydroxides, select amphoteric ones: Fe (OH) 2, KOH, Al (OH) 3, Ca (OH) 2, Fe (OH) 3, Pb (OH) 2.

8. What is called a reaction scheme? reaction equation?

9. What is the name of the numbers in the reaction equation? What do they show?

10. How to go from the reaction scheme to the equation?

11. What substances do basic oxides interact with? Amphoteric oxides? Acid oxides?

12. What substances do bases interact with?

13. What substances do acids interact with?

14. What substances do salts interact with?

15. Determine the mass fractions of elements in nitric acid HNO 3.

16. What metals interact with alkalis?

17. What metals interact with solutions of sulfuric and hydrochloric acids?

18. What products are formed during the interaction of metals with nitric acid of various concentrations?

19. What reactions are called decomposition reactions? Connections? Substitutions? Redox?

20. Write the reaction equations: CrCl 3 + NaOH→; CrCl 3 + 2NaOH→; CrCl 3 + 3NaOH→; CrCl 3 + NaOH (excess) →.

21. Write the reaction equations: Al + KOH →; Al + KOH + H 2 O →.

22. What is called an atom? chemical element? Molecule?

23. What elements are classified as metals? Nonmetals? Why?

24. What is called the chemical formula of a substance? What does she show?

25. What is called the structural formula of a substance? What does she show?

26. What is called the amount of substance?

27. What is called a mole? What does it show? How many structural units are there in a mole of a substance?

28. What masses of elements are indicated in the Periodic system?

29. What is called the relative atomic, molecular masses? How are they defined? What are their units of measurement?

30. What is called the molar mass of a substance? How is it defined? What is its unit of measurement?

31. What conditions are called normal conditions?

32. What is the volume of 1 mol of gas at N.C.? 5 moles of gas at n.o.?

33. What does an atom consist of?

34. What does the nucleus of an atom consist of? What is the charge on the nucleus of an atom? What determines the charge of the nucleus of an atom? What determines the mass of the nucleus of an atom?

35. What is called a mass number?

36. What is called the energy level? How many electrons are in a single energy level?

37. What is called an atomic orbital? How is she portrayed?

38. What characterizes the main quantum number? Orbital quantum number? Magnetic quantum number? Spin quantum number?

39. What is the relationship between the principal and orbital quantum numbers? Between orbital and magnetic quantum numbers?

40. What is the name of electrons with \u003d 0? = 1? = 2? = 3? How many orbitals correspond to each of the given states of the electron?

41. What state of an atom is called the ground state? Excited?

42. How many electrons can be located in one atomic orbital? What is the difference?

44. How many and what sublevels can be located on the first energy level? On the second? On the third? On the fourth?

45. Formulate the principle of least energy, Klechkovsky's rules, Pauli's principle, Hund's rule, periodic law.

46. ​​What changes periodically for the atoms of elements?

47. What do the elements of one subgroup have in common? One period?

48. How do the elements of the main subgroups differ from the elements of the secondary subgroups?

49. Compose the electronic formulas of the ions Cr +3, Ca +2, N -3. How many unpaired electrons do these ions have?

50. What energy is called ionization energy? Affinity for an electron? Electronegativity?

51. How do the radii of atoms and ions change in a group and in a period of D.I. Mendeleev?

52. How do the electronegativity of atoms in a group and in a period of the Periodic system of D.I. Mendeleev?

53. How do the metallic properties of the elements and the properties of their compounds change in the group and in the period of the Periodic system of D.I. Mendeleev?

54. Make formulas for higher oxides of aluminum, phosphorus, bromine, manganese.

55. How is the number of protons, neutrons and electrons in an atom determined?

56. How many protons, neutrons and electrons are there in a zinc atom?

57. How many electrons and protons are contained in Cr +3, Ca +2, N -3 ions?

58. Formulate the law of conservation of mass? What remains constant during any chemical reaction?

59. What parameter remains constant in isobaric chemical reactions?

60. Formulate the law of composition constancy. For substances of what structure is it valid?

61. Formulate Avogadro's law and its consequences.

62. If the nitrogen density of a gas is 0.8, then what is the molar mass of the gas?

63. In the event of a change in what external parameters does the molar volume of a gas change?

64. Formulate the combined gas law.

65. For equal volumes of different gases under the same conditions, will the masses of gases be equal?

66. Formulate Dalton's law. If the total pressure of a mixture of nitrogen and hydrogen is 6 atm, and the volume content of hydrogen is 20%, then what are the partial pressures of the components?

67. Write down the Mendeleev-Clapeyron equation (state of an ideal gas).

68. What is the mass of a gas mixture consisting of 11.2 liters of nitrogen and 11.2 liters of fluorine (n.o.)?

69. What is called a chemical equivalent? Molar mass equivalent?

70. How is the molar mass of equivalents of simple and complex substances determined?

71. Determine the molar masses of the equivalents of the following substances: O 2, H 2 O, CaCl 2, Ca (OH) 2, H 2 S.

72. Determine the equivalent of Bi(OH) 3 in the reaction Bi(OH) 3 + HNO 3 = Bi(OH) 2 (NO 3) + H 2 O.

73. Formulate the law of equivalents.

74. What is called the molar volume of the equivalent of a substance? How is it defined?

75. Formulate the law of volumetric relations.

76. What volume of oxygen will be required for the oxidation of 8 m 3 of hydrogen (n.o.) according to the reaction 2H 2 + O 2 ↔ 2H 2 O?

77. What volume of hydrogen chloride is formed by the interaction of 15 liters of chlorine and 20 liters of hydrogen?

78. What is meant by a chemical bond? Specify the characteristics of a chemical bond.

79. What is a measure of the strength of a chemical bond?

80. What affects the distribution of electron density?

81. What determines the shape of a molecule?

82. What is called valency?

83. Determine the nitrogen valencies in the following compounds: N 2, NH 3, N 2 H 4, NH 4 Cl, NaNO 3.

84. What is called the degree of oxidation?

85. What bond is called covalent?

86. Indicate the properties of a covalent bond.

87. How does the polarity of a bond change in the series KI, KBr, KCl, KF?

88. Molecules of what substance are non-polar: oxygen, hydrogen chloride, ammonia, acetic acid.

89. What is meant by hybridization of valence orbitals?

90. Determine the types of hybridization of central atoms in the following substances: beryllium fluoride, aluminum chloride, methane.

91. How does the type of hybridization affect the spatial structure of molecules?

92. What bond is called ionic? Under the influence of what forces does it arise?

93. What bond is called metallic?

94. What properties do substances with a metallic type of chemical bond have?

95. What is the maximum number of -bonds that can be formed between two atoms in a molecule?

96. How is the absolute electronegativity of an atom of an element determined?

97. Arrange the elements in ascending order of their electronegativity: Fe, C, Ag, H, Cl.

98. What is called the dipole moment of communication? How is it calculated?

99. What features do substances with an atomic crystal lattice have? With a molecular crystal lattice?

100. What bond is called hydrogen? What determines its strength? Between the molecules of which inorganic substances does it occur?

Thermodynamics and kinetics of chemical reactions

1. What does thermodynamics study?

2. What is called a thermodynamic system? What types of systems exist?

3. What are called state parameters? What parameters are called intensive, extensive? Name the main parameters of a chemical system.

4. What is called a process? Spontaneous process? Cycle? Equilibrium process? An unbalanced process? Reversible process?

5. What is called a phase? Homogeneous, heterogeneous system?

6. What is called the state function?

7. What characterizes the internal energy U? What does internal energy depend on?

8. What is called heat Q? Which reactions are exothermic or endothermic? How do heat and enthalpy change during their flow?

9. What is called work p∆V?

10. Formulate the first law of thermodynamics. Write it down mathematically.

11. Formulate the first law of thermodynamics for isothermal, isochoric and isobaric processes.

12. What is called enthalpy?

13. What is called the thermal effect of the reaction? What determines the thermal effect of a reaction?

14. What equation is called thermodynamic? Thermochemical?

15. What conditions are called standard?

16. What is called the enthalpy of reaction? Standard enthalpy of reaction?

17. What is called the enthalpy of formation of a substance? The standard enthalpy of formation of a substance?

18. What is the standard state of matter? What is the enthalpy of formation of a simple substance in the standard state?

19. The enthalpy of formation of H 2 SO 3 is equal in magnitude to the heat effect of the reaction: H 2 (g) + S (tv) + 1.5O 2 (g) H 2 SO 3 (g); H 2 (g) + SO 2 (g) + 0.5O 2 (g) H 2 SO 3 (g); H 2 O (g) + SO 2 (g) H 2 SO 3 (g); 2H (g) + S (tv) + 3O (g) H 2 SO 3 (g).

20. The interaction of 1 mole of hydrogen and 1 mole of bromine released 500 kJ of heat. What is ∆Н arr, HBr?

21. In the formation of 5 moles of substance A x B y, 500 kJ of heat was absorbed. What is ∆Н arr of this substance?

22. What is called the enthalpy of combustion? Standard enthalpy of combustion? Heat capacity?

23. Formulate the law of Hess, the first and second consequences of it.

24. Which expression is applicable to calculate ∆Н р of the reaction 2A + 3B 2C according to Hess' law:

∆Н r = 2∆Н arr, С + 2∆Н arr, A + 3∆Н arr, B; ∆Н r = 2∆Н arr, С – (2∆Н arr, A + 3∆Н arr, B);

∆Н r = 2∆Н arr, A + 3∆Н arr, B –2∆Н arr, C; ∆Н r = – 2∆Н arr, C – (2∆Н arr, A + 3∆Н arr, B)?

25. The standard enthalpy of combustion (∆H 0 combust) of methanol CH 4 O (l) (M = 32 g / mol) is -726.6 kJ / mol. How much heat will be released when 2.5 kg of a substance is burned?

26. In what case is the standard enthalpy of combustion of one substance equal to the standard enthalpy of formation of another substance?

27. For what substances is the standard enthalpy of combustion equal to zero: CO, CO 2, H 2, O 2?

28. For the reaction 2Cl 2 (g) + 2H 2 O (g) 4HCl (g) + O 2 (g), calculate the standard enthalpy (kJ) if the standard enthalpies of formation of substances are known:

29. ∆H = -1410.97 kJ/mol; ∆H = -2877.13 kJ/mol. How much heat will be released during the joint combustion of 2 moles of ethylene and 4 moles of butane?

30. ∆H = -1410.97 kJ/mol; ∆H = -2877.13 kJ/mol. What amount of heat will be released when burning 0.7 kg of a gas mixture consisting of 20% ethylene and 80% butane?

31. The standard enthalpy of the reaction MgCO 3 (tv) → MgO (tv) + CO 2 (g) is 101.6 kJ; standard enthalpies of formation of MgO (tv) and CO 2 (g): -601.0 and -393.5 kJ / mol, respectively. What is the standard enthalpy of formation of magnesium carbonate MgCO 3 ?

32. What is called the thermodynamic probability of a system? What is called entropy? How is entropy expressed in terms of thermodynamic probability?

33. Formulate the second law of thermodynamics.

34. What is called the standard entropy of a substance?

35. Formulate the third law of thermodynamics (Planck's postulate).

36. What is called the entropy of a reaction? The standard entropy of the reaction?

37. Which expression is applicable to calculate ∆S p of the reaction CH 4 + CO 2 2CO + 2H 2:

∆S p \u003d S + S + S + S; ∆S p \u003d S + S + 2S + 2S;

∆S p \u003d 2S + 2S - S + S; ∆S p \u003d 2S + 2S - S - S?

38. For the reaction 2Cl 2 (u) + 2H 2 O (l) 4HCl (g) + O 2 (g), calculate the standard entropy (J / K) if the standard entropies of the formation of substances are known:

39. What is called Gibbs free energy? What is its relationship with other thermodynamic functions?

40. How is the direction of the reaction determined by the sign of the Gibbs energy of a reaction?

41. At what temperatures is a reaction possible if ∆H<0, ∆S>0; ∆H<0, ∆S<0; ∆H>0, ∆S>0; ∆H>0, ∆S<0.

42. How is the equilibrium temperature of the process determined?

43. What is called the Gibbs energy of the reaction ∆G p? The standard Gibbs energy of the reaction?

44. What expression is applicable to calculate ∆G p of the reaction 4NH 3 (g) + 5O 2 (g) 4NO (g) + 6H 2 O (l)

∆G p \u003d ∆G 4 + ∆G 5 + ∆G 4 + ∆G 6; ∆G p = ∆G + ∆G + ∆G + ∆G ;

∆G p \u003d 4∆G + 5∆G - 4∆G - 6∆G; ∆G p \u003d 4∆G + 6∆G - 4∆G - 5∆G ?

45. For the reaction HNO 3 (l) + HNO 2 (l) 2NO 2 (g) + H 2 O (l), calculate the standard Gibbs energy (kJ) if the standard Gibbs energies of the formation of substances are known:

46. ​​For the reaction Fe (tv) + Al 2 O 3 (tv) → Al (tv) + Fe 2 O 3 (tv), determine the equilibrium temperature and the possibility of the process occurring at 125 0 С, if ∆Н = 853.8 kJ / mole; ∆S = 37.68 J/mol K.

47. What is meant by the rate of a chemical reaction?

48. Formulate the law of mass action.

49. In 40 s, as a result of two reactions Zn + 2HCl \u003d ZnCl 2 + H 2 (1) and Zn + 2HBr \u003d ZnBr 2 + H 2 (2), 8 g of zinc chloride and zinc bromide were formed. Compare reaction rates.

50. If in the reaction 3Fe (NO 3) 2 (solution) + 4HNO 3 \u003d 3Fe (NO 3) 3 (solution) + NO (g) + 2H 2 O (g) the concentration of Fe (NO 3) 2 increase by 7 times, and the concentration of HNO 3 by 4 times, how will the reaction rate change?

51. Make a kinetic equation for the reaction Sb 2 S 3 (tv) + 3H 2 (g) 2Sb (tv) + 3H 2 S (g).

52. How is the rate of a multistage reaction determined?

53. How will the rate of the direct reaction CO (g) + 3H 2 (g) CH 4 (g) + H 2 O (g) change with an increase in system pressure by 3 times?

54. What is called a speed constant? What does it depend on?

55. What is called activation energy? What does it depend on?

56. The rate constant of a certain reaction at a temperature of 310 K is 4.6 ∙ 10 -5 l mol -1 s -1, and at a temperature of 330 K 6.8 ∙ 10 -5 l mol -1 s -1. What is the activation energy equal to?

57. The activation energy of a certain reaction is 250 kJ / mol. How will the rate constant change when the reaction temperature changes from 320 K to 340 K?

58. Write down the Arrhenius equation and the van't Hoff rule.

59. The activation energy of reaction (1) is 150 kJ/mol, the activation energy of reaction (2) is 176 kJ/mol. Compare the rate constants k 1 and k 2 .

60. How to explain the increase in the reaction rate with increasing temperature?

61. What is called the temperature coefficient of the reaction?

62. What is the temperature coefficient of the reaction if the rate constant of a certain reaction at 283 and 308 K is 1.77 and 7.56 l mol -1 s -1, respectively?

63. At a temperature of 350 K, the reaction ended in 3 s, and at a temperature of 330 K, in 28 s. How long will it take to finish at a temperature of 310 K?

64. How does the activation energy affect the temperature coefficient of the reaction?

65. What is called a catalyst? Inhibitor? A promoter? Catalytic poison?

66. What is called chemical equilibrium? How long does the system remain in equilibrium?

67. How are the rates of direct and reverse reactions related at the moment of equilibrium?

68. What is called the equilibrium constant? What does it depend on?

69. Express the equilibrium constant of the reactions 2NO + O 2 ↔ 2NO 2; Sb 2 S 3 (tv) + 3H 2 ↔ 2Sb (tv) + 3H 2 S (g).

70. At a certain temperature, the equilibrium constant of the reaction N 2 O 4 ↔ 2NO 2 is 0.16. There was no NO 2 in the initial state, and the equilibrium concentration of NO 2 was 0.08 mol/L. What will be the equilibrium and initial concentration of N 2 O 4 ?

71. Formulate Le Chatelier's principle. How do changes in temperature, concentration, and total pressure affect the mixing of equilibrium?

72. Chemical dynamic equilibrium in the system was established at 1000 K and a pressure of 1 atm. atm. What is the equilibrium constant K p of this reaction?

73. Equilibrium concentrations (mol / l) of the components of the gas-phase system in which the reaction took place

3N 2 H 4 ↔ 4NH 3 + N 2 are equal to: \u003d 0.2; =0.4; =0.25. What is the equilibrium constant of the reversible

74. Equilibrium concentrations (mol / l) of the components of the gas-phase system in which the reaction occurs

N 2 + 3H 2 ↔ 2NH 3 are equal to: = 0.12; =0.14; =0.1. Determine the initial concentrations of N 2 and H 2 .

75. Equilibrium concentrations of the components of the gas phase of the system in which the reaction occurs

C (tv) + CO 2 ↔ 2CO at 1000 K and P total \u003d 1 atm., Equal to CO 2 - 17% vol. and CO - 83% vol. What is the constant

reaction equilibrium?

76. The equilibrium constant K with a reversible gas-phase reaction CH 4 + H 2 O ↔ CO + 3H 2 at a certain temperature is 9.54 mol 2 l -2. The equilibrium concentrations of methane and water are 0.2 mol/l and 0.4 mol/l, respectively. Determine the equilibrium concentrations of CO and H 2.

77. Write down the relationship between the equilibrium constant K p and the Gibbs energy ∆G of a reversible reaction occurring under isothermal conditions.

78. Determine the equilibrium constant K p of the gas-phase reversible reaction COCl 2 ↔ CO + Cl 2; ∆H 0 = 109.78 kJ,

∆S 0 = 136.62 J/K at 900 K.

79. Equilibrium constant K p gas-phase reaction PCl 3 + Cl 2 ↔ PCl 5; ∆H 0 \u003d -87.87 kJ at 450 K is 40.29 atm -1. Determine the Gibbs energy of this process (J/K).

80. Write down the relationship between K p and K with a reversible gas-phase reaction 2CO + 2H 2 ↔ CH 4 + CO 2.

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1 4. Chemical process. Why and how do chemical reactions take place? Thermodynamics and kinetics In the first half of the 19th century, there was a need to improve heat engines that produce mechanical work due to chemical combustion reactions. Such heat engines at that time were firearms and steam engines. As a result, in the middle of the 19th century, thermodynamics or the mechanical theory of heat was created. The term thermodynamics "thermodynamics" was proposed in 1851 by the English scientist William Thomson (Lord Kelvin from 1892) (). The German researcher Rudolf Julius Emanuel Clausius () called the new science Mechanische Warmeteorie "the mechanical theory of heat". Modern definition: Chemical thermodynamics is the science of the dependence of the direction and limits of transformations of substances on the conditions in which these substances are located. Unlike other sections of physical chemistry (structure of matter and chemical kinetics), chemical thermodynamics can be applied without knowing anything about the structure of matter. Such a description requires much less initial data. A specific object of thermodynamic research is called a thermodynamic system or simply a system isolated from the surrounding world by real or imaginary surfaces. A system can be a gas in a vessel, a solution of reagents in a flask, a crystal of a substance, or even a mentally selected part of these objects. According to the levels of interaction with the environment, thermodynamic systems are usually divided into: open systems exchange matter and energy with the environment (for example, living objects); closed ones exchange only energy (for example, a reaction in a closed flask or a flask with a reflux condenser), the most frequent object of chemical thermodynamics; isolated ones do not exchange either matter or energy and retain a constant volume (approximation of the reaction in a thermostat). A rigorous thermodynamic consideration is possible only for isolated systems that do not exist in the real world. At the same time, thermodynamics can quite accurately describe closed and even open systems. In order for a system to be described thermodynamically, it must consist of a large number of particles, comparable to the Avogadro number and thus comply with the laws of statistics. The properties of the system are divided into extensive (summing), for example, total volume, mass, and intensive (equalizing) pressure, temperature, concentration, etc. The most important for state function calculations are thermodynamic functions whose values ​​depend only on the state of the system and do not depend on the path of transition between states. The process in thermodynamics is not the development of an event in time, but a sequence of equilibrium states of the system leading from the initial set of thermodynamic variables to the final one. Thermodynamics makes it possible to completely solve the problem, if the process under study as a whole is described by a set of equilibrium stages. eleven

2 In thermodynamic calculations, numerical data (tabular) on the thermodynamic properties of substances are used. Even small datasets of such data make it possible to calculate many different processes. To calculate the equilibrium composition of the system, it is not necessary to write down the equations of possible chemical reactions, it is enough to take into account all substances that can, in principle, constitute an equilibrium mixture. Thus, chemical thermodynamics does not give a purely calculated (non-empirical) answer to the question why? and even more so how? ; it solves problems according to the principle if ..., then .... For thermal calculations, the first law of thermodynamics is the most important one of the forms of the law of conservation of energy. His formulations: Energy is neither created nor destroyed. A perpetual motion machine (perpetuum mobile) of the first kind is impossible. In any isolated system, the total amount of energy is constant. Yu.R. Mayer (1842) [1] was the first to discover the relationship between chemical reactions and mechanical energy; the mechanical equivalent of heat was measured by J.P. Joule (). For thermochemical calculations, the law of conservation of energy is used in the formulation of G.I. Hess: “When a chemical compound is formed, the same amount of heat is always released, regardless of whether the formation of this compound occurs directly or indirectly and in several steps." Hess announced this law of “constancy of heat sums” in a report at a conference of the Russian Academy of Sciences on March 27, 1840. In this case, the work done by a chemical reaction at constant pressure consists of a change in internal energy and the work of expansion of the resulting gas: ΔQ p = ΔU + pδv isobaric (i.e. running at constant pressure) process. This function is called enthalpy (from the Greek enthalpo I heat) [ 3 ]: ΔQ p = ΔH = ΔU + pδv Another definition: the difference in enthalpies in two states of the system is equal to the thermal effect of the isobaric process. 1. In 1840, the German doctor Julius Robert Mayer () worked as a ship's doctor on a flight from Europe to Java. He noticed that the venous blood in the tropics is lighter than in Germany, and concluded that in the tropics less oxygen is needed to maintain the same body temperature. Therefore, heat and work can be mutually transformed. In 1842 Mayer theoretically estimated the mechanical equivalent of heat at 365 kgm (modern 427 kgm) 2 Trifonov D.N. "Direct and noble character" (On the 200th anniversary of Herman Ivanovich Hess) 3. The name enthalpy was proposed by the Dutch physicist Geike Kamerling-Onnes (). 12

3 It is enthalpy that has proved to be convenient for describing the operation of both steam engines and firearms, since in both cases the expansion of hot gases or water vapor is used. There are extensive tables containing data on the standard enthalpies of formation of substances ΔH o 298. The indices mean that for chemical compounds the enthalpies of formation of 1 mol of them from simple substances are given, taken in the most stable modification at 1 atm (1, Pa or 760 mm Hg. st) and 298.15 K (25 ° C). When it comes to ions in solution, the standard concentration is 1 mol/l. For the simple substances themselves, the enthalpy of formation equal to 0 is adopted (except for white phosphorus, not the most stable, but the most reproducible form of phosphorus). The sign of enthalpy is determined from the point of view of the system itself: when heat is released, the change in enthalpy is negative, when heat is absorbed, the change in enthalpy is positive. An example of a thermochemical calculation of an extremely complex reaction: The enthalpy of formation of glucose from carbon dioxide and water cannot be determined by direct experiment, it is impossible to obtain glucose from simple substances. But we can calculate the enthalpies of these processes. 6 C + 6 H O 2 = C 6 H 12 O 6 (ΔH x -?) Such a reaction is impossible 6 CO H 2 O = C 6 H 12 O O 2 (ΔH y -?) The reaction takes place in green leaves, but together with others processes Find ΔH x algebraically. Using Hess' law, it is enough to combine three combustion equations: 1) C + O 2 = CO 2 ΔH 1 = -394 kJ 2) H 2 + 1/2 O 2 = H 2 O (steam) ΔH 2 = -242 kJ 3) C 6 H 12 O O 2 \u003d 6 CO H 2 O ΔH 3 \u003d kj 6 CO 2 ΔH 1 \u003d 6 (-394) kJ 2) 6 H O 2 \u003d 6 H 2 O (steam) ΔH 2 \u003d 6 (-242) kJ 3) 6 CO H 2 O \u003d C 6 H 12 O O 2 ΔH 3 = kJ When calculating the enthalpy, we take into account that when the “reversal” of equation 3, it changed sign: that ΔH y corresponds to the reverse process of photosynthesis, i.e. burning glucose. Then ΔH y \u003d -ΔH 3 \u003d kJ When solving, no data on the structure of glucose were used; the mechanism of its combustion was also not considered. Problem Determine the enthalpy of obtaining 1 mol of ozone O 3 from oxygen, if it is known that the combustion of 1 mol of oxygen in an excess of hydrogen releases 484 kJ, and the combustion of 1 mol of ozone in an excess of hydrogen releases 870 kJ The second law of thermodynamics. Entropy The second law of thermodynamics according to W. Thomson (1851): a process is impossible in nature, the only result of which would be mechanical work done by cooling the heat reservoir. 13

4 The formulation of R. Clausius (1850): heat itself cannot transfer from a colder body to a warmer one, or: it is impossible to design a machine that, acting through a circular process, will only transfer heat from a colder body to a warmer one. The earliest formulation of the second law of thermodynamics appeared before the first law, based on the work done in France by S. Carnot (1824) and its mathematical interpretation by E. Clapeyron (1834) as the efficiency of an ideal heat engine: efficiency = (T 1 - T 2) / T 1 Carnot and Clapeyron formulated the law of conservation of caloric for a weightless indestructible liquid, the content of which determines the body temperature. The theory of caloric dominated thermodynamics until the middle of the 19th century, while the laws and relations derived from the concept of caloric proved to be valid in the framework of the molecular-kinetic theory of heat. To find out the reasons for the occurrence of spontaneous processes that occur without heat release, it became necessary to describe heat by the method of generalized forces, similarly to any mechanical work (A), through a generalized force (F) and a generalized coordinate (in this case, thermal) [4]: ​​da = Fdx For thermal reversible processes, we get: dq = TdS I.e. Initially, the entropy S is the thermal coordinate of the state, which was introduced (Rudolf Clausius, 1865) to standardize the mathematical apparatus of thermodynamics. Then for an isolated system, where dq = 0, we get: In a spontaneous process ΔS > 0 In an equilibrium process ΔS = 0 In a non-spontaneous process ΔS< 0 В общем случае энтропия изолированной системы или увеличивается, или остается постоянной: ΔS 0 Энтропия свойство системы в целом, а не отдельной частицы. В 1872 г. Л.Больцман [ 5 ] предложил статистическую формулировку второго закона термодинамики: изолированная система эволюционирует преимущественно в направлении большей термодинамическоой вероятности. В 1900 г. М.Планк вывел уравнение для статистического расчета энтропии: S = k b lnw W число различных состояний системы, доступное ей при данных условиях, или термодинамическая вероятность макросостояния системы. k b = R/N A = 1, эрг/град постоянная Больцмана 4. Полторак О.М., Термодинамика в физической химии. Учеб. для хим. и хим-технол. спец. вузов, М.: Высш. шк., с., стр Больцман Людвиг (Boltzmann, Ludwig) (), австрийский физик. Установил фундаментальное соотношение между энтропией физической системы и вероятностью ее состояния, доказал статистический характер II начала термодинамики Современный биограф Людвига Больцмана физик Карло Черчиньяни пишет: Только хорошо поняв второе начало термодинамики, можно ответить на вопрос, почему вообще возможна жизнь. В 1906 г. Больцман покончил с собой, поскольку обманулся в любви; он посвятил свою жизнь атомной теории, но любовь его осталась без взаимности, потому что современники не могли понять масштаб его картины мира 14

5 It should always be remembered that the second law of thermodynamics is not absolute; it loses its meaning for systems containing a small number of particles and for systems on a cosmic scale. The second law, especially in a statistical formulation, is not applicable to living objects, which are open systems and constantly reduce entropy, creating perfectly ordered molecules, for example, due to the energy of sunlight. Living systems are characterized by self-organization, which the Chilean neuroscientist Humberto Maturana called autopoiesis (self-creation) in 1970. Living systems not only constantly move away from the classical thermodynamic equilibrium, but also make the environment non-equilibrium. As early as 1965, the American atmospheric chemist James Lovelock suggested that the criterion for the existence of life on Mars be the equilibrium composition of the atmosphere. The Earth's atmosphere contains simultaneously oxygen (21% by volume), methane (0.00018%), hydrogen (0.00005%), carbon monoxide (0.00001%) is a clearly non-equilibrium mixture at temperatures C. The Earth's atmosphere is an open system, in the formation of which living organisms constantly participate. The atmosphere of Mars is dominated by carbon dioxide (95% - compare with 0.035% on Earth), oxygen in it is less than 1%, and reducing gases (methane) have not yet been detected. Consequently, the atmosphere of Mars is practically in equilibrium, all reactions between the gases contained in it have already taken place. From these data, Lovelock concluded that there is currently no life on Mars Gibbs energy The introduction of entropy made it possible to establish criteria for determining the direction and depth of any chemical process (for a large number of particles in equilibrium). Macroscopic systems reach equilibrium when the energy change is compensated by the entropy component: At constant pressure and temperature: ΔH p = TΔS p or Δ(H-TS) ΔG = 0 Gibbs energy[6] or Gibbs free energy or isobaric-isothermal potential Gibbs energy change as a criterion for the possibility of a chemical reaction For a given temperature ΔG = ΔH - TΔS At ΔG< 0 реакция возможна; при ΔG >0 reaction is not possible; at ΔG = 0 the system is in equilibrium. 6 Gibbs Josiah Willard (), American physicist and mathematician, one of the founders of chemical thermodynamics and statistical physics. Gibbs published a fundamental treatise On the Equilibrium of Heterogeneous Substances, which became the basis of chemical thermodynamics. fifteen

6 The possibility of a spontaneous reaction in an isolated system is determined by the combination of the signs of the energy (enthalpy) and entropy factors: Sign of ΔH Sign of ΔS Possibility of spontaneous reaction + No + Yes Depends on the ratio of ΔH and TΔS + + Depends on the ratio of ΔH and TΔS There are extensive tabular data on standard values ΔG 0 and S 0, allowing to calculate ΔG 0 of the reaction. 5. Chemical kinetics Predictions of chemical thermodynamics are most correct in their forbidding part. If, for example, for the reaction of nitrogen with oxygen, the Gibbs energy is positive: N 2 + O 2 \u003d 2 NO ΔG 0 \u003d +176 kJ, then this reaction will not go spontaneously, and no catalyst will help it. The well-known factory process for obtaining NO from the air requires huge expenditures of energy and a non-equilibrium process (quenching of products by rapid cooling after passing a mixture of gases through an electric arc). On the other hand, far from all reactions for which ΔG< 0, спешат осуществиться на практике. Куски каменного угля могут веками лежать на воздухе, хотя для реакции C + O 2 = CO 2 ΔG 0 = -395 кдж Предсказание скорости химической реакции, а также выяснение зависимости этой скорости от условий проведения реакции осуществляет химическая кинетика наука о химическом процессе, его механизме и закономерностях протекания во времени. Скорость химической реакции определяется как изменение концентрации одного из участвующих в реакции веществ (исходное вещество или продукт реакции) в единицу времени. Для реакции в общем виде aa + bb xx + yy скорость описывается кинетическим уравнением: v = -ΔC (A) /Δt = ΔC (X) /Δt = k C m n (A) C (B) k называется константой скорости реакции. Строго говоря, скорость определяется не как конечная разность концентраций, а как их производная v = -dc (A) /dt; степенные показатели m и n обычно не совпадают со стехиометрическими коэффициентами в уравнении реакции. Порядком реакции называется сумма всех показателей степеней m и n. Порядок реакции по реагенту A равен m. Большинство реакций являются многостадийными, даже если они описываются простыми стехиометрическими уравнениями. В этом случае обычно получается сложное кинетическое уравнение реакции. Например, для реакции H 2 + Br 2 = 2 HBr dc (HBr) /dt = kc (H2) C (Br2) 0,5 / (1 + k C (HBr) / C (Br2)) 16

7 Such a complex dependence of the rate on concentrations indicates a multistage reaction mechanism. A chain mechanism has been proposed for this reaction: Br 2 Br. +Br. initiation of the Br chain. + H 2 HBr + H. chain continuation H. + Br 2 HBr + Br. chain continuation H. + HBr H 2 + Br. Br inhibition. +Br. Br 2 chain termination The number of reactant molecules participating in a simple one-step reaction consisting of one elementary act is called the molecularity of the reaction. Monomolecular reaction: C 2 H 6 \u003d 2 CH 3. Bimolecular reaction: CH 3. + CH 3. \u003d C 2 H 6 Examples of relatively rare trimolecular reactions: 2 NO + O 2 \u003d 2 NO 2 2 NO + Cl 2 \u003d 2 NOCl H. + H. + Ar = H 2 + Ar A feature of the 1st order reactions proceeding according to the scheme: And the products is the constancy of the half-life t 0.5 of the time during which half of the starting substance will turn into products. This time is inversely proportional to the reaction rate constant k. t 0.5 = 0.693/k i.e. the half-life for a first order reaction is a constant and characteristic of the reaction. In nuclear physics, the half-life of a radioactive isotope is its important property Temperature dependence of the rate of reactions Most reactions of practical importance are accelerated by heating. The dependence of the reaction rate constant on temperature is expressed by the Arrhenius equation [ 7 ] (1889): k = Aexp(-E a /RT) The factor A is related to the frequency of collisions of particles and their orientation during collisions; E a is the activation energy of a given chemical reaction. To determine the activation energy of a given reaction, it is sufficient to measure its rate at two temperatures. The Arrhenius equation describes the temperature dependence not only for simple chemical processes. Psychological studies of people with different body temperatures (from 36.4 to 39 ° C) have shown that the subjective sense of time (clock counting speed) and 7 Svante August Arrhenius (Arrhenius) () Swedish physical chemist, creator of the theory of electrolytic dissociation, academician of the Swedish Royal Academy of Sciences. On the basis of ideas about the formation of active particles in electrolyte solutions, Arrhenius put forward a general theory of the formation of "active" molecules in chemical reactions. In 1889, studying the inversion of cane sugar, he showed that the rate of this reaction is determined by the collision of only "active" molecules. A sharp increase in this rate with increasing temperature is determined by a significant increase in the number of "active" molecules in the system. To enter into a reaction, the molecules must have some additional energy compared to the average energy of the entire mass of the molecules of the substance at a certain temperature (this additional energy will later be called the activation energy). Arrhenius outlined ways to study the nature and form of the temperature dependence of the reaction rate constants. 17

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Topic 3. General laws of chemical processes.

Chemical thermodynamics and kinetics

Introduction

Central to chemistry is the doctrine of the transformation of substances, including energy and the kinetics of chemical reactions. The assimilation of this doctrine makes it possible to predict the possibility and direction of chemical processes, calculate the energy effects and energy costs, the rate of production and yield of products in the reaction, influence the rate of chemical processes, and also prevent undesirable reactions in certain devices, installations and devices.

3.1. Chemical thermodynamics and kinetics

The exchange of energy between the system under study and the externalenvironment describe the laws that it studiesthermodynamics. The application of the laws of thermodynamics in chemistry allows us to solve the problem of the fundamental possibility of various processes, the conditions for their implementation,divide the degree of conversion of reactants into chimic reactions and evaluate their energetics.

Chemical thermodynamics , examines the relationship between work and energy in relation to chemical transformations.

Thermal and mechanical energy - algebraicquantities. Signs of quantitiesQ and BUT in thermodynamicsviewed in relation to the system. Energy, receivedreceived by the system is indicated by the sign “+”, given to the systemstem - sign "-".

Variables that determine the state of the SIstems are calledstate parameters. Among them in chemistry, the most commonly used pressure, temperature, volume, composition of the system. System status and prooutgoing changes in it are also characterized with the help ofstate functions, depending on the state parameters and not depending on the path of the system transition fromone state to another. These include internalenergy, enthalpy, entropy, isobaric-isothermal potential, etc.

Processes occurring at constant pressure -isobaric, at constant volumeisochoric, at constant temperature -isothermal. Majority chemical reactions take place in open vessels,i.e. at a constant pressure equal to atmospheric.

Chemical kineticsstudies the characteristics of a chemical process, such as the rate of a reaction and its dependence on external conditions.

3.2. Energy of chemical processes

Breakdown occurs during a chemical reactionsome chemical bonds and the formation of new ones. This process is accompanied by the release or absorption of heat.you, light or some other kind of energy. Energy effReaction effects are studied by the science of thermochemistry. In thermochemistryuse thermochemical reaction equations, whichwhich take into account:

aggregate state of matter;

thermal effect of the reaction (Q).

These equations often use fractional coefficients. So, the reaction equations for the formation of 1 mol of gasfigurative water is written as follows:

H 2 (g) + 1 / 2O 2 (g) \u003d H 2 O (g) + 242 kJ (*)

The symbol (d) indicates that hydrogen, oxygen andwater is in the gas phase. "+242 kJ" - means thatAs a result of this reaction, so much heat is released atformation of 1 mole of water.

The importance of taking into account the state of aggregation is related to the fact thatthat the heat of formation of liquid water is greater byheat released during the condensation of steam:

H 2 (g) + 1 / 2O 2 (g) \u003d H 2 O (g) + 286 kJ (**)

Condensation process:

H 2 O (g) \u003d H 2 O (g) + 44 kJ (***)

In addition to the thermal effect, thermodynamics usesyut the concept of "change in heat content" - enthalpy(reserve of internal energy) during the reaction ( H)

Exothermic reactions: heat is released Q > 0

internal energy is decreasing H<0

Endothermic reactions: heat is absorbed Q< 0

internal energy increases H>0.

Thus, the reaction (*) of water formation is exothermic.Thermal effect of the reaction:Q = 242 kJ,  H = -242 kJ.

Enthalpy of formation of chemical compounds

Standard enthalpy (heat) of formation chemical compound  H 0 f,V,298 is the change in enthalpy in the process of formation of one mole of this compound, which is in the standard state (p = 1 atm; T = 25 0 C), from simple substances, also in standard states and phases and modifications that are thermodynamically stable at a given temperature .

The standard enthalpies of formation of simple substances are taken equal to zero if their states of aggregation and modifications are stable under standard conditions.

The standard enthalpies of formation of substances are collected and summarized in reference books.

3.2. 1. Thermochemical calculations

The independence of the heat of a chemical reaction from the process path at p=const was established in the first half of the 19th century. Russian scientist G.I. Hess: the thermal effect of a chemical reaction does not depend on the path of its occurrence, but depends only on the nature and physical state of the initial substances and reaction products.

For most reactions, the change in the thermal effect within the temperature limits of practical importance is small. Therefore, in the future we will use  H 0 f,B,298 and are assumed to be constant in calculations.

Consequence from Hess' law– the heat effect of a chemical reaction is equal to the sum of the heats (enthalpies) of the formation of the reaction products minus the sum of the heats (enthalpies) of the formation of the starting substances.

When using the corollary from the Hess law in thermochemical calculations, it should be borne in mind that stoichiometric coefficients in the reaction equation should be taken into account in algebraic summation.

So, for the reaction equation aA + bB = cC + dD, the thermal effect  H is equal to

H=(s  N ex.C +d N ex.D) – (а N ex.A +v N ex.B) (*)

Equation (*) makes it possible to determine both the thermal effect of the reaction from the known enthalpies of formation of the substances participating in the reaction, and one of the enthalpies of formation, if the thermal effect of the reaction and all other enthalpies of formation are known.

Fuel combustion heat

The thermal effect of the oxidation reaction with oxygen of the elements that make up the substance to the formation of higher oxides is called the heat of combustion of this substance
.

Example: determine the heat of combustion of ethanol C 2 H 5 OH (g)

If a calculation conducted for
with the formation of liquid water, then the heat of combustion is called higher, if with the formation of gaseous water, then lower. By default, they usually mean the higher calorific value.

In technical calculations, the specific heat of combustion Q T is used, which is equal to the amount of heat released during the combustion of 1 kg of a liquid or solid substance or 1 m 3 of a gaseous substance, then

Q T = -  N ST  1000/M (for w, tv.)

Q T = -  N ST  1000/22.4 (for city),

where M is the mass of a mole of a substance, 22.4 l is the volume of a mole of gas.

3.3. Chemical and phase equilibrium

3.3.1. Chemical equilibrium

Reversible reactions - chemical reactions occurring simultaneously in two opposite directions.

Chemical equilibrium - the state of the system in which the rate of the direct reaction (V 1 ) is equal to the rate of the reverse reaction (V 2 ). In chemical equilibrium, the concentrations of substances remain unchanged. Chemical equilibrium has a dynamic character: forward and reverse reactions do not stop at equilibrium.

The state of chemical equilibrium is quantitatively characterized by the equilibrium constant, which is the ratio of the constants of the straight line (K 1 ) and reverse (K 2 ) reactions.

For the reaction mA + nB « pC + dD the equilibrium constant is

K=K 1 /k 2 = ([C] p[D] d) / ([A] m[B] n)

The equilibrium constant depends on the temperature and the nature of the reactants. The larger the equilibrium constant, the more the equilibrium is shifted towards the formation of direct reaction products.

Ways to shift the balance

Le Chatelier's principle. If an external influence is made on a system in equilibrium (concentration, temperature, pressure change), then it favors the flow of one of the two opposite reactions that weakens this effect.

V 1

A+B

V 2

Pressure. An increase in pressure (for gases) shifts the equilibrium towards a reaction leading to a decrease in volume (i.e., to the formation of a smaller number of molecules).

V 1

A+B

; an increase in P leads toV 1 >V 2

V 2

An increase in temperature shifts the equilibrium position towards an endothermic reaction (i.e. towards a reaction proceeding with the absorption of heat)

V 1

A+B

B + Q, then increase t° C leads to V 2 > V 1

V 2

V 1

A+B

B - Q, then increase t° C leads to V 1 > V 2

V 2

Increasing the concentration of starting materials and removing products from the reaction sphere shifts the equilibrium towards the direct reaction. Increasing concentrations of starting materials [A] or [B] or [A] and [B]: V 1 > V 2 .

Catalysts do not affect the equilibrium position.

3.3.2. Phase equilibria

The equilibrium of the process of transition of a substance from one phase to another without changing the chemical composition is called phase equilibrium.

Examples of phase equilibrium:

Solid.............Liquid

Liquid .......... Steam

3.3.3. Reaction rate and methods of its regulation

Speed ​​reaction is determined by the change in the molar concentration of one of the reactants:

V = ± (С 2 - С 1) / (t 2 - t 1) \u003d ± D FROM / D t

where C 1 and C 2 - molar concentrations of substances at time t 1 and t2 respectively (sign (+) - if the rate is determined by the reaction product, sign (-) - by the starting material).

Reactions occur when molecules of reactants collide. Its speed is determined by the number of collisions and the likelihood that they will lead to a transformation. The number of collisions is determined by the concentrations of the reacting substances, and the probability of a reaction is determined by the energy of the colliding molecules.

Factors affecting the rate of chemical reactions

The nature of the reactants. An important role is played by the nature of chemical bonds and the structure of the molecules of the reagents. Reactions proceed in the direction of the destruction of less strong bonds and the formation of substances with stronger bonds. For example, to break bonds in H 2 and N 2 high energies are required; such molecules are not very reactive. To break bonds in highly polar molecules (HCl, H 2 O) less energy is required and the reaction rate is much faster. Reactions between ions in electrolyte solutions proceed almost instantaneously.

Examples: Fluorine reacts explosively with hydrogen at room temperature; bromine reacts with hydrogen slowly even when heated.

Calcium oxide reacts vigorously with water, releasing heat; copper oxide - does not react.

Concentration. With an increase in concentration (the number of particles per unit volume), collisions of reactant molecules occur more often - the reaction rate increases.

The law of active masses (K. Guldberg, P. Waage, 1867)

The rate of a chemical reaction is directly proportional to the product of the concentrations of the reactants.

aA + bB + . . .® . . .

V=k[A] a[B] b . . .

The reaction rate constant k depends on the nature of the reactants, temperature, and catalyst, but does not depend on the concentrations of the reactants.

The physical meaning of the rate constant is that it is equal to the reaction rate at unit concentrations of the reactants.

For heterogeneous reactions, the concentration of the solid phase is not included in the reaction rate expression.

Temperature. With an increase in temperature for every 10° C, the reaction rate increases by 2-4 times (Van't Hoff's Rule). As the temperature increases from t 1 to t 2 the change in reaction rate can be calculated by the formula:

(t 2 - t 1 ) / 10

Vt 2 / Vt 1

= g

(where Vt 2 and Vt 1 - reaction rates at temperatures t 2 and t1 respectively;gis the temperature coefficient of this reaction).

Van't Hoff's rule is applicable only in a narrow temperature range. More accurate is the Arrhenius equation:

k \u003d Ae -Ea / RT

where

A is a constant depending on the nature of the reactants;

R is the universal gas constant;

Ea is the activation energy, i.e. the energy that colliding molecules must have in order for the collision to result in a chemical transformation.

Energy diagram of a chemical reaction.

exothermic reaction

Endothermic reaction

A - reagents, B - activated complex (transition state), C - products.

The higher the activation energy Ea, the more the reaction rate increases with increasing temperature.

1. The contact surface of the reactants. For heterogeneous systems (when substances are in different states of aggregation), the larger the contact surface, the faster the reaction proceeds. The surface of solids can be increased by grinding them, and for soluble substances by dissolving them.

3.3.4. Mechanisms of chemical reactions, oscillatory reactions

Classification of chemical reactions

I . According to the number and composition of the starting materials and reaction products:

1) Reactions connections are reactions in which two or more substances form one substance of a more complex composition. Reactions of the combination of simple substances are always redox reactions. Complex substances can also participate in compound reactions.

2) Reactions decomposition Reactions in the course of which two or more simpler substances are formed from one complex substance.
Decomposition products of the initial substance can be both simple and complex substances.

Decomposition reactions usually proceed when substances are heated and are endothermic reactions. Like compound reactions, decomposition reactions can proceed with or without changing the oxidation states of the elements;

3) Reactions substitution - these are reactions between simple and complex substances, during which the atoms of a simple substance replace the atoms of one of the elements in the molecule of a complex substance, as a result of the substitution reaction, a new simple and a new complex substance are formed.
These reactions are almost always redox reactions.

4) Reactions exchange - these are reactions between two complex substances, the molecules of which exchange their constituent parts.
Exchange reactions always proceed without electron transfer, that is, they are not redox reactions.

II . On the basis of changes in the degree of oxidation

1) Reactions that go without changing the oxidation state - neutralization reactions

2) With a change in the degree of oxidation

III . Depending on the presence of a catalyst

1) Non-catalytic (go without the presence of a catalyst);

2) catalytic (comes with a catalyst)

IV . According to the thermal effect

1) exothermic (with heat release):

2) Endothermic (with heat absorption):

V . On the basis of reversibility

1) irreversible (flow in one direction only):

2) reversible (flowing simultaneously in the forward and reverse directions):

VI . On the basis of homogeneity

1) homogeneous (flowing in a homogeneous system):

2) Heterogeneous (flowing in an inhomogeneous system):

According to the flow mechanism All reactions can be divided into simple and complex. Simple reactions proceed in one stage and are called one-stage.

Complex reactions proceed either sequentially (multi-stage reactions), or in parallel, or in series-parallel.

Each reaction step can involve one molecule (monomolecular reactions), two molecules (bimolecular reactions), and three molecules (trimolecular reactions).

Vibrational reactions - a class of chemical reactions occurring in an oscillatory mode, in which some reaction parameters (color, concentration of components, temperature, etc.) change periodically, forming a complex spatio-temporal structure of the reaction medium.

(System bromate-malonic acid-cerium Belousov-Zhabotinsky reaction)

3.4. Catalysis

Substances that participate in reactions and increase its rate, remaining unchanged at the end of the reaction, are calledcatalysts .

The mechanism of action of catalysts is associated with a decrease in the activation energy of the reaction due to the formation of intermediate compounds.

At homogeneous catalysis the reactants and the catalyst constitute one phase (they are in the same state of aggregation).

At heterogeneous catalysis - different phases (they are in different states of aggregation).

In some cases, the course of undesirable chemical processes can be drastically slowed down by adding to the reaction mediuminhibitors(phenomenon " negative catalysis ").

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