1. The rate of chemical reactions. Definition of the concept. 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. Rate constant, its physical meaning. The influence of the nature of the reactants, temperature and the presence of a catalyst on the reaction rate constant.
The rate of a homogeneous reaction is a value numerically equal to the change in the molar concentration of any reaction participant per unit time.
The average reaction speed v avg in the time interval from t 1 to t 2 is determined by the relation:
The main factors influencing the rate of a homogeneous chemical reaction:
- - the nature of the reacting substances;
- - molar concentrations of reagents;
- - pressure (if gases are involved in the reaction);
- - temperature;
- - presence of a catalyst.
The rate of a heterogeneous reaction is a value numerically equal to the change in the chemical amount of any reaction participant per unit time per unit interface area: .
According to the stages, chemical reactions are divided into simple (elementary) and complex. Most chemical reactions are complex processes occurring 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 the 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 reactants, but only gaseous or dissolved ones. So, for the carbon combustion reaction:
C (k) + O 2 (g) > CO 2 (g)
the speed 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 for a homogeneous reaction depends on the nature of the reactants, temperature and catalyst.
2. The influence 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. The ratio of activation energy and chemical bond energy in the original molecules. Transition state, or activated complex. Activation energy and thermal effect of the reaction (energy diagram). Dependence of the temperature coefficient of the reaction rate on the activation energy.
As 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, which is the same, by 10 K) is called the temperature coefficient of the rate of a chemical reaction (g):
where are the reaction rates at temperatures T 2 and T 1, respectively; r is the temperature coefficient of the reaction rate.
The dependence of the reaction rate on temperature is approximately determined by Van't Hoff's empirical rule: with every 10 degree increase in temperature, 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 possible within the framework of the Arrhenius activation theory. According to this theory, a chemical reaction can occur when only active particles collide. Active particles are those that have a certain energy characteristic of a given reaction that 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 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 a 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 reacting substances.
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 reacting substances; 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 lower the activation energy, the greater the reaction rate constant. Even a slight decrease in activation energy (for example, when adding a catalyst) 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. The 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. The 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 influence of substances, the quantity and nature of which 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 a reaction. Each of the intermediate processes involving a catalyst occurs with a lower activation energy than a 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.
During homogeneous catalysis, the catalyst forms an intermediate compound with a reagent, which reacts with a second reagent at a high speed or quickly decomposes to release a reaction product.
An example of homogeneous catalysis: the oxidation of sulfur(IV) oxide to sulfur(VI) oxide with oxygen using the nitrous method for producing sulfuric acid (here the catalyst is nitrogen(II) oxide, which easily reacts with oxygen).
In heterogeneous catalysis, the reaction occurs on the surface of the catalyst. The initial stages are the diffusion of reagent particles to the catalyst and their adsorption (i.e., absorption) by the surface of the catalyst. Reactant molecules interact with atoms or groups of atoms located on the surface of the catalyst, forming intermediate surface compounds. The redistribution of 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 by oxygen in the presence of vanadium(V) oxide.
Examples of catalytic processes in industry and technology: ammonia synthesis, synthesis of nitric and sulfuric acids, cracking and reforming of oil, afterburning of products of incomplete combustion of gasoline in cars, etc.
Examples of catalytic processes in nature are numerous, since most biochemical reactions occurring in living organisms are classified as catalytic reactions. The catalysts for such reactions are protein substances called enzymes. There are about 30,000 enzymes in the human body, each of which catalyzes only one type of process (for example, salivary ptyalin catalyzes only the conversion of starch into glucose).
4. Chemical equilibrium. 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. The 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 that occur simultaneously in two opposite directions (forward and reverse) 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 cessation of the reaction. In general, for any reversible reaction aA + bB - dD + eE, regardless of its mechanism, the following relation holds:
At steady equilibrium, the product of the concentrations of the reaction products divided by the product of the concentrations of the starting substances 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.
A change in 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 forward 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 displacement of the equilibrium position.
If, during the transition from one equilibrium state to another, the concentrations of substances written on the right side of the reaction equation increase, the equilibrium is said to shift to the right. If, during the transition from one equilibrium state to another, the concentrations of substances written on the left side of the reaction equation increase, the equilibrium is said to shift to the left.
The direction of the shift in chemical equilibrium as a result of changes in external conditions is determined by Le Chatelier’s principle: If an external influence is exerted on a system in a state of chemical equilibrium (change temperature, pressure or concentrations of substances), then it will favor the occurrence of whichever of the two opposite processes occurs. 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 of equilibrium to the right; an increase in the concentration of the component written on the right side of the equation leads to a shift of equilibrium to the left;
When the temperature increases, the equilibrium shifts towards the endothermic reaction, and when the temperature decreases, towards the exothermic reaction;
- - When the pressure increases, the equilibrium shifts towards a reaction that reduces the number of molecules of gaseous substances in the system, and when the pressure decreases, 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 living nature. Unbranched and branched chemical reactions (using the example of 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 occur under the influence of light. A photochemical reaction occurs if the reagent absorbs radiation quanta characterized by an energy quite specific for a given reaction.
In the case of some photochemical reactions, absorbing energy, the molecules of the reagent pass into an excited state, i.e. become active.
In other cases, a photochemical reaction occurs if quanta of such high energy are absorbed that chemical bonds are broken and molecules dissociate into atoms or groups of atoms.
The greater the irradiation intensity, the greater the speed of the photochemical reaction.
An example of a photochemical reaction in living nature is photosynthesis, i.e. formation of organic cell substances due to light energy. In most organisms, photosynthesis occurs 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
The functioning of vision processes is also based on photochemical processes.
A chain reaction is a reaction that is a chain of elementary acts of interaction, and the possibility of each act of interaction depends on the success of the previous act.
The stages of a chain reaction are chain initiation, chain development and chain termination.
The initiation of a chain occurs when, due to an external source of energy (quanta of electromagnetic radiation, heating, electrical discharge), active particles with unpaired electrons (atoms, free radicals) are formed.
During the development of the chain, radicals interact with the original molecules, and new radicals are formed in each act of interaction.
Chain termination occurs when two radicals collide and transfer the energy released in the process to a third body (a molecule resistant to decay or the wall of a vessel). The chain can also terminate if a low-active 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 that determine the direction of a chemical reaction. Elements of chemical thermodynamics. Concepts: phase, system, environment, macro- and microstates. Basic thermodynamic characteristics. Internal energy of the system and its change during chemical transformations. Enthalpy. The relationship between enthalpy and internal energy of a system. Standard enthalpy of a substance. Changes in enthalpy in systems during chemical transformations. Thermal effect (enthalpy) of a chemical reaction. Exo- and endothermic processes. Thermochemistry. Hess's law. Thermochemical calculations.
Thermodynamics studies the patterns of energy exchange between the system and the external environment, the possibility, direction and limits of the spontaneous occurrence of chemical processes.
A thermodynamic system (or simply a system) is a body or a group of interacting bodies, mentally isolated 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 chemical composition and properties and is separated from other parts of the system by an interface.
The state of a 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 by the parameters of each individual particle.
The 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 parameters of the state and does not depend on the path of transition to a given state, is called a state function. The functions of the state are, in particular:
U - internal energy;
H - enthalpy;
S - entropy;
G - Gibbs energy (free energy or isobaric-isothermal potential).
The internal energy of a 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 full account of all these components is impossible, when studying a system thermodynamically, we consider the change in its internal energy during the transition from one state (U 1) to another (U 2):
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 that are not affected by any forces other than external pressure forces.
If a process occurs in a system at a constant volume (i.e., there is no work against external pressure forces), then A = 0. Then the thermal effect of a process occurring at a constant volume, Qv, is equal to the change in the internal energy of the system:
Most chemical reactions encountered in everyday life occur at constant pressure (isobaric processes). If no forces other than constant external pressure act on the system, then:
A = p(V2 - V 1 ) = pV
Therefore, in our case (p = const):
Qp=U + pV
Q р = U 2 - U 1 + p(V 2 - V 1 ), where
Q p = (U 2 +pV 2 ) - (U 1 +pV 1 ).
The function U + pV is called 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,
that is, the thermal effect of a reaction at constant pressure and temperature T is equal to the change in enthalpy of the system during the reaction. It depends on the nature of the reagents and products, their physical state, conditions (T, p) for the reaction, as well as on the amount of substances participating in the reaction.
The enthalpy of a reaction is the change in enthalpy of a system in which the reactants interact in quantities equal to the stoichiometric coefficients in the reaction equation.
The enthalpy of a reaction is called standard if the reactants and products of the reaction 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 = 25 o C or 298 K; p = 101.325 kPa).
The standard state of a substance existing 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 a pressure of 101.325 kPa; in gaseous form - gas with its own pressure of 101.325 kPa.
For a solute, the standard state is considered to be its state in solution at a molality of 1 mol/kg, 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 mole of a given substance from simple substances in their standard states is called the standard enthalpy of formation of this substance.
Example entry: (CO 2) = - 393.5 kJ/mol.
The standard enthalpy of formation of a simple substance in the most stable (given p and T) state of aggregation is assumed to be 0. If an element forms several allotropic modifications, then only the most stable (given p and T) modification has a zero standard enthalpy of formation.
Typically, thermodynamic quantities are determined under standard conditions:
p = 101.32 kPa and T = 298 K (25 o C).
Chemical equations that specify enthalpy changes (heat effects of reactions) are called thermochemical equations. In the literature you can find two forms of writing thermochemical equations.
Thermodynamic form of writing the thermochemical equation:
C (graphite) + O 2 (g) CO 2 (g); = - 393.5 kJ.
Thermochemical form of writing the thermochemical equation of 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 perspective 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 an environmental perspective. Therefore, if the system releases heat, then it is assumed that Q > 0.
Exothermic is a process that occurs with the release of heat (DH< 0).
Endothermic is a process that occurs 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 states of the system and does not depend on the path of transition of the system from one state to another.”
Corollary from Hess's 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 substances, taking into account stoichiometric coefficients:
- (reactions) = (cont.) -(out.)
- 7. The concept of entropy. Changes in entropy during phase transformations and chemical processes. The concept of the isobaric-isothermal potential of a system (Gibbs energy, free energy). The relationship between the magnitude of the change in the Gibbs energy and the magnitude of the change in enthalpy and entropy of the reaction (basic thermodynamic relationship). Thermodynamic analysis of the possibility and conditions of chemical reactions. Features of the flow of chemical processes in living organisms.
Entropy S is a value proportional to the logarithm of the number of equally probable microstates (W) through which a given macrostate can be realized:
S = k lnW
The unit of entropy is J/mol?K.
Entropy is a quantitative measure of the degree of disorder of 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 especially particles in the gaseous state. On the contrary, all processes as a result of which the order 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.
Standard entropy of a simple substance, as opposed to enthalpy of formation simple substance is not equal to zero.
For entropy, a statement similar to that discussed above for H is valid: the change in the entropy of a 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 with the calculation of enthalpy, the summation is carried out taking into account the stoichiometric coefficients.
The direction in which a chemical reaction spontaneously occurs 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) a tendency to achieve the most probable state, i.e. a state that can be realized in the largest number of equally probable ways (microstates), i.e.:
DH > min, DS > max.
The state function, which simultaneously reflects the influence of both of the above-mentioned trends on the direction of the flow 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 can be seen, the Gibbs energy has the same dimension as enthalpy and is therefore 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 equal to:
G= H-TS
As in the case of H and S, the change in Gibbs energy G as a result of a chemical reaction (Gibbs energy of the reaction) is equal to the sum of the Gibbs energies of formation of the reaction products minus the sum of the Gibbs energies of formation of the starting materials; the summation is carried out taking into account the number of moles of substances participating in the reaction.
The Gibbs energy of formation of a substance is referred 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 constant temperature and pressure, chemical reactions can spontaneously proceed only in a direction in which the Gibbs energy of the system decreases (G0). This is a condition for the fundamental possibility of carrying out this process.
The table below shows the possibility and conditions for the reaction to occur with various combinations of H and S signs:
By the sign of G one can judge the possibility (impossibility) of the spontaneous occurrence of a particular process. If the system is influenced, 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 occur to form complex organic compounds; The driving force behind such processes is solar radiation and oxidation reactions in the cell.
“FUNDAMENTALS OF CHEMICAL THERMODYNAMICS, CHEMICAL KINETICS AND EQUILIBRIUM”
Fundamentals of chemical thermodynamics
1 . What chemical thermodynamics studies:
1) the rate of chemical transformations and the mechanisms of these transformations;
2) energy characteristics of physical and chemical processes and the ability of chemical systems to perform useful work;
3) conditions for shifting chemical equilibrium;
4) the influence of catalysts on the rate of biochemical processes.
2. An open system is a system that:
3. A closed system is a system that:
1) does not exchange either matter or energy with the environment;
2) exchanges both matter and energy with the environment;
3) exchanges energy with the environment, but does not exchange matter;
4) exchanges matter with the environment, but does not exchange energy.
4. An isolated system is a system that:
1) does not exchange either matter or energy with the environment;
2) exchanges both matter and energy with the environment;
3) exchanges energy with the environment, but does not exchange matter;
4) exchanges matter with the environment, but does not exchange energy.
5. To what type of thermodynamic systems does the solution located in a sealed ampoule placed in a thermostat belong?
1) isolated;
2) open;
3) closed;
4) stationary.
6. What type of thermodynamic systems does the solution in the sealed ampoule belong to?
1) isolated;
2) open;
3) closed;
4) stationary.
7. What type of thermodynamic systems does a living cell belong to?
1) open;
2) closed;
3) isolated;
4) equilibrium.
8 . What parameters of a thermodynamic system are called extensive?
1) the magnitude of which does not depend on the number of particles in the system;
3) the value of which depends on the state of aggregation of the system;
9. What parameters of a thermodynamic system are called intensive?
!) whose magnitude does not depend on the number of particles in the system;
2) the magnitude of which depends on the number of particles in the system;
3) the value of which depends on the state of aggregation;
4) the magnitude of which depends on time.
10 . Functions of the state of a thermodynamic system are quantities that:
1) depend only on the initial and final state of the system;
2) depend on the process path;
3) depend only on the initial state of the system;
4) depend only on the final state of the system.
11 . What quantities are functions of the state of the system: a) internal energy; b) work; c) warmth; d) enthalpy; d) entropy.
3) all quantities;
4) a, b, c, d.
12 . Which of the following properties are intensive: a) density; b) pressure; c) mass; d) temperature; e) enthalpy; e) volume?
3) b, c, d, f;
13. Which of the following properties are extensive: a) density; b) pressure; c) mass; d) temperature; e) enthalpy; e) volume?
3) b, c, d, f;
14 . What forms of energy exchange between the system and the environment are considered by thermodynamics: a) heat; b) work; c) chemical; d) electric; e) mechanical; e) nuclear and solar?
2) c, d, e, f;
3) a, c, d, e, f;
4) a, c, d, e.
15. Processes occurring at a constant temperature are called:
1) isobaric;
2) isothermal;
3) isochoric;
4) adiabatic.
16 . Processes occurring at constant volume are called:
1) isobaric;
2) isothermal;
3) isochoric;
4) adiabatic.
17 . Processes occurring at constant pressure are called:
1) isobaric;
2) isothermal;
3) isochoric;
4) adiabatic.
18 . The internal energy of a system is: 1) the entire energy reserve of the system, except for the potential energy of its position and the kinetic energy of the system as a whole;
2) the entire energy reserve of the system;
3) the entire energy reserve of the system, except for the potential energy of its position;
4) a quantity characterizing the degree of disorder in the arrangement of particles of the system.
19 . What law reflects the relationship between work, heat and internal energy of a system?
1) the second law of thermodynamics;
2) Hess's law;
3) the first law of thermodynamics;
4) van't Hoff's law.
20 . The first law of thermodynamics reflects the relationship between:
1) work, heat and internal energy;
2) Gibbs free energy, enthalpy and entropy of the system;
3) work and heat of the system;
4) work and internal energy.
21 . Which equation is the mathematical expression of the first law of thermodynamics for isolated systems?
l)AU=0 2)AU=Q-p-AV 3)AG = AH-TAS
22 . Which equation is the mathematical expression of the first law of thermodynamics for closed systems?
1)AU=0; 2)AU=Q-p-AV;
3) AG = AH - T*AS;
23 . Is the internal energy of an isolated system a constant or variable quantity?
1) constant;
2) variable.
24 . In an isolated system, the reaction of hydrogen combustion occurs with the formation of liquid water. Does the internal energy and enthalpy of the system change?
1) internal energy will not change, enthalpy will change;
2) internal energy will change, enthalpy will not change;
3) internal energy will not change, enthalpy will not change;
4) internal energy will change, enthalpy will change.
25 . Under what conditions is the change in internal energy equal to the heat received by the system from the environment?
1) at constant volume;
3) at constant pressure;
4) under no circumstances.
26 . The thermal effect of a reaction occurring at constant volume is called a change:
1) enthalpy;
2) internal energy;
3) entropy;
4) Gibbs free energy.
27 . The enthalpy of a reaction is:
28. Chemical processes during which the enthalpy of the system decreases and heat is released into the external environment are called:
1) endothermic;
2) exothermic;
3) exergonic;
4) endergonic.
29 . Under what conditions is the change in enthalpy equal to the heat received by the system from the environment?
1) at constant volume;
2) at constant temperature;
3) at constant pressure;
4) under no circumstances.
30 . The thermal effect of a reaction occurring at constant pressure is called a change:
1) internal energy;
2) none of the previous definitions are correct;
3) enthalpy;
4) entropy.
31. What processes are called endothermic?
32 . What processes are called exothermic?
1) for which AN is negative;
2) for which AG is negative;
3) for which AN is positive;
4) for which AG is positive.
33 . Specify the formulation of Hess's law:
1) the thermal effect of the reaction depends only on the initial and final state of the system and does not depend on the reaction path;
2) the heat absorbed by the system at a constant volume is equal to the change in the internal energy of the system;
3) the heat absorbed by the system at constant pressure is equal to the change in enthalpy of the system;
4) the thermal effect of the reaction does not depend on the initial and final state of the system, but depends on the reaction path.
34. What law underlies the calculation of caloric content of food?
1) van't Hoff;
3) Sechenov;
35. When oxidizing which substances under body conditions, more energy is released?
1) proteins;
3) carbohydrates;
4) carbohydrates and proteins.
36 . A spontaneous process is a process that:
1) carried out without the help of a catalyst;
2) accompanied by the release of heat;
3) carried out without external energy consumption;
4) proceeds quickly.
37 . Entropy of a reaction is:
1) the amount of heat that is released or absorbed during a chemical reaction under isobaric-isothermal conditions;
2) the amount of heat that is released or absorbed during a chemical reaction under isochoric-isothermal conditions;
3) a value characterizing the possibility of spontaneous occurrence of the process;
4) a quantity characterizing the degree of disorder in the arrangement and movement of particles in the system.
38 . What state function characterizes the tendency of a system to achieve a probable state that corresponds to the maximum randomness of the distribution of particles?
1) enthalpy;
2) entropy;
3) Gibbs energy;
4) internal energy.
39 . What is the relationship between the entropies of three aggregate states of one substance: gas, liquid, solid:
I) S (g) > S (g) > S (tv); 2) S(solid)>S(g)>S(g); 3)S(g)>S(g)>S(TB); 4) the state of aggregation does not affect the entropy value.
40 . Which of the following processes should exhibit the greatest positive change in entropy:
1) CH3OH (s) --> CH,OH (g);
2) CH4OH (s) --> CH 3 OH (l);
3) CH,OH (g) -> CH4OH (s);
4) CH,OH (l) -> CH3OH (sol).
41 . Choose the correct statement: the entropy of the system increases when:
1) increase in pressure;
2) transition from liquid to solid state of aggregation
3) increase in temperature;
4) transition from gaseous to liquid state.
42. Which thermodynamic function can be used to predict whether a reaction will occur spontaneously in an isolated system?
1) enthalpy;
2) internal energy;
3) entropy;
4) potential energy of the system.
43 . Which equation is the mathematical expression of the 2nd law of thermodynamics for isolated systems?
44 . If the system reversibly receives a quantity of heat Q at temperature T, then about T;
2) increases by the amount Q/T;
3) increases by an amount greater than Q/T;
4) increases by an amount less than Q/T.
45 . In an isolated system, a chemical reaction occurs spontaneously to form a certain amount of product. How does the entropy of such a system change?
1) increases
2) decreases
3) does not change
4) reaches the minimum value
46 . Indicate in which processes and under what conditions the change in entropy can be equal to the work of the process?
1) in isobaric conditions, at constant P and T;
2) in isochoric conditions, at constant V and T;
H) the change in entropy is never equal to work; 4) in isothermal conditions, at constant P and 47 . How will the bound energy of the system TS change when heated and when it condenses?
1) increases with heating, decreases with condensation;
2) decreases with heating, increases with condensation;
3) there is no change in T-S;
4) increases with heating and condensation.
48 . What parameters of the system must be kept constant so that the sign of the change in entropy can be used to judge the direction of the spontaneous course of the process?
1) pressure and temperature;
2) volume and temperature;
3) internal energy and volume;
4) only temperature.
49 . In an isolated system, all spontaneous processes proceed in the direction of increasing disorder. How does entropy change?
1) does not change;
2) increases;
3) decreases;
4) first increases and then decreases.
50 . Entropy increases by the amount Q/T for:
1) reversible process;
2) irreversible process;
3) homogeneous;
4) heterogeneous.
51 How does the entropy of the system change due to forward and reverse reactions during ammonia synthesis?
3) entropy does not change during the reaction;
4) entropy increases for forward and reverse reactions.
52 . What simultaneously acting factors determine the direction of a chemical process?
1) enthalpy and temperature;
2) enthalpy and entropy;
3) entropy and temperature;
4) changes in Gibbs energy and temperature.
53. Under isobaric-isothermal conditions, the maximum work performed by the system is:
1) equal to the decrease in Gibbs energy;
2) greater loss of Gibbs energy;
3) less loss of Gibbs energy;
4) is equal to the loss of enthalpy.
54 . What conditions must be met so that the maximum work in the system is accomplished due to the decrease in Gibbs energy?
1) it is necessary to maintain constant V and t;
2) it is necessary to maintain constant P and t;
3) it is necessary to maintain constant AH and AS;
4) it is necessary to maintain constant P&V
55 . What causes the maximum useful work done in a chemical reaction at constant pressure and temperature?
1) due to the decrease in Gibbs energy;
3) due to an increase in enthalpy;
4) due to a decrease in entropy.
56. How is the maximum useful work performed by a living organism under isobaric-isothermal conditions?
1) due to the loss of enthalpy;
2) due to an increase in entropy;
3) due to the decrease in Gibbs energy;
4) due to an increase in the Gibbs energy.
57 . What processes are called endergonic?
58. What processes are called exergonic?
2) AG 0; 4) AG > 0.
59. The spontaneous nature of the process is best determined by assessing:
1) entropy;
3) enthalpy;
2) Gibbs free energy;
4) temperature.
60 . What thermodynamic function can be used to predict the possibility of spontaneous processes occurring in a living organism?
1) enthalpy;
3) entropy;
2) internal energy;
4) Gibbs free energy.
61 . For reversible processes, the change in Gibbs free energy...
1) always equal to zero;
2) always negative;
3) always positive;
62 . For irreversible processes, the change in free energy:
1) always equal to zero;
2) always negative;
3) always positive;
4) positive or negative depending on the circumstances.
63. Under isobaric-isothermal conditions, only such processes can spontaneously occur in a system, as a result of which the Gibbs energy is:
1) does not change;
2) increases;
3) decreases;
4) reaches its maximum value.
64 . For a certain chemical reaction in the gas phase at constant P and TAG > 0. In what direction does this reaction spontaneously proceed?
D) in the forward direction;
2) cannot occur under these conditions;
3) in the opposite direction;
4) is in a state of equilibrium.
65 . What is the sign AG of the ice melting process at 263 K?
66 . In which of the following cases is the reaction not feasible at any temperature?
1)AH>0;AS>0; 2)AH>0;AH
3)A#4)AH= 0;AS = 0.
67. In which of the following cases is the reaction possible at any temperature?
1)DN 0; 2)AH 0; AS > 0; 4)AH = 0;AS = 0.
68 . If AN
1) [AN] > ;
2) for any ratio of AN and TAS; 3)(AH]
4) [AN] = [T-A S].
69 . At what values of the sign of AH and AS are only exothermic processes possible in the system?
70. At what ratios of AN and T* AS is the chemical process directed towards an endothermic reaction:
71 . At what constant thermodynamic parameters can a change in enthalpy serve as a criterion for the direction of a spontaneous process? What sign of DH under these conditions indicates a spontaneous process?
1) at constant S and P, AN
3) with constant Put, AN
2) at constant 5 and P, AN > 0; 4) at constant Vn t, AH > 0.
72 . Is it possible and in what cases to judge by the sign of the enthalpy change during a chemical reaction about the possibility of its occurrence at constant Ti P1
1) possible, if LA » T-AS;
2) under these conditions it is impossible;
3) possible, if AN « T-AS;
4) possible if AN = T-AS.
73 . The reaction ZN 2 + N 2 -> 2NH 3 is carried out at 110°C, so that all reactants and products are in the gas phase. Which of the following values is conserved during the reaction?
2) entropy;
3) enthalpy;
74 . Which of the following statements are true for reactions occurring under standard conditions?
1) endothermic reactions cannot occur spontaneously;
2) endothermic reactions can occur at sufficiently low temperatures;
3) endothermic reactions can occur at high temperatures if AS > 0;
4) endothermic reactions can occur at high temperatures if AS
75 . What are the features of biochemical processes: a) obey the principle of energy coupling; b) usually reversible; c) complex; d) only exergonic (AG
1) a, b, c, d;
2) b, c, d; 3) a, 6, c; 4) c, d.
76 . Exergonic reactions in the body occur spontaneously, since:
77 . Endergonic reactions in the body require energy supply, since: 1) AG >0;
78 . When any peptide AH 0 is hydrolyzed, will this process occur spontaneously?
1) will be, since AG > 0;
3) will not happen, since AG > 0;
2) will be, since AG
4) will not be, since AG
79 . The calorie content of nutrients is called energy:
1) 1 g of nutrients released during complete oxidation;
2) 1 mole of nutrients released during complete oxidation;
3) necessary for complete oxidation of 1 g of nutrients;
4) 1 mole of nutrients required for complete oxidation.
80 . For the process of thermal denaturation of many enzymes, LA > 0 and AS > 0. Can this process occur spontaneously?
1) it can at high temperatures, since \T-AS\ > |BP];
2) can at low temperatures, since \T-AS\
3) cannot, since \T-AS\ > |AH];
4) cannot, because \T-AS\
81 . For the process of thermal hydration of many AN proteins
1) can at sufficiently low temperatures, since |AH| > \T-AS\;
2) can at sufficiently low temperatures, since |АА|
3) can at high temperatures, since |AH)
4) cannot at any temperature.
ProgramParameters chemical reactions, chemical equilibrium; - calculate thermal effects and speed chemical reactions... reactions; - basics physical and colloid chemistry, chemical kinetics, electrochemistry, chemical thermodynamics and thermochemistry; ...
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RegulationsModule 2. Basic physics chemical patterns of occurrence chemical processes Basics chemical thermodynamics. Basics chemical kinetics. Chemical equilibrium. Module 3.. Basics chemistry of solutions General...
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DocumentSimple substances. On this basis V chemical thermodynamics a system for calculating thermal effects has been created..., Cr2O3? TOPIC 2. CHEMICAL KINETICS AND CHEMICAL EQUILIBRIUM As was shown earlier, chemical thermodynamics allows you to predict the fundamental...
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Work program4.1.5. Redox processes. Basics electrochemistry Oxidation-reduction processes. ... Methods for quantitatively expressing the composition of solutions. 5 Chemical thermodynamics 6 Kinetics And equilibrium. 7 Dissociation, pH, hydrolysis 8 ...
The rate of chemical reactions. Definition of the concept. 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. Rate constant, its physical meaning. The influence of the nature of the reactants, temperature and the presence of a catalyst on the reaction rate constant.
1. With. 102-105; 2. With. 163-166; 3. With. 196-207, p. 210-213; 4. With. 185-188; 5. With. 48-50; 6. With. 198-201; 8. With. 14-19
Homogeneous reaction rate - this is a quantity numerically equal to the change in the concentration of any reaction participant per unit time.
Average reaction speed v avg in the time interval from t 1 to t 2 is determined by the relation:
Main factors influencing the rate of a homogeneous chemical reaction :
- the nature of the reacting substances;
- reagent concentration;
- pressure (if gases are involved in the reaction);
- temperature;
- presence of a catalyst.
Heterogeneous reaction rate - this is a quantity numerically equal to the change in the concentration of any reaction participant per unit time on a unit surface: .
According to the stages, chemical reactions are divided into elementary And complex. Most chemical reactions are complex processes occurring in several stages, i.e. consisting of several elementary processes.
For elementary reactions it is true law of mass action: the rate of an elementary chemical reaction at a given temperature is directly proportional to the product of the concentrations of the reacting substances in powers equal to the stoichiometric coefficients of the reaction equation.
For an elementary reaction aA + bB → ... the reaction rate, according to the law of mass action, is expressed by the relation:
wheres (A) and With (IN) - molar concentrations of reactants A And IN; a And b- corresponding stoichiometric coefficients; k – rate constant of a given reaction .
For heterogeneous reactions, the equation of the law of mass action does not include the concentrations of all reactants, but only gaseous or dissolved ones. So, for the carbon combustion reaction:
C(k) + O 2 (g) → CO 2 (g)
the velocity equation has the form .
The physical meaning of the rate constant is 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 for a homogeneous reaction depends on the nature of the reactants, temperature and catalyst.
The influence 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. The ratio of activation energy and chemical bond energy in the original molecules. Transition state, or activated complex. Activation energy and thermal effect of the reaction (energy diagram). Dependence of the temperature coefficient of the reaction rate on the activation energy.
1. With. 106-108; 2. With. 166-170; 3. With. 210-217; 4. With. 188-191; 5. With. 50-51; 6. With. 202-207; 8 . With. 19-21.
As temperature increases, the rate of a chemical reaction usually increases.
The value showing how many times the reaction rate increases when the temperature increases by 10 degrees (or, what is the same, by 10 K) is called temperature coefficient of chemical reaction rate (γ):
where are the reaction rates, respectively, at temperatures T 2 and T 1 ; γ - temperature coefficient of reaction rate.
The dependence of the reaction rate on temperature is approximately determined empirically van't Hoff's rule: with every 10 degree increase in temperature, 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 possible within the framework of the Arrhenius activation theory. According to this theory, a chemical reaction can only occur when active particles collide. Active are called particles that have a certain, characteristic of a given reaction, energy 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.
Activated complex - this is an intermediate unstable group formed during the collision of active particles and is in a state of redistribution of bonds. Reaction products are formed during the decomposition of the activated complex.
Activation energy And E A 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 bond in the molecules of the reacting substances.
In activation theory, the influence temperature on the rate of a chemical reaction is described by the Arrhenius equation for the rate constant of a chemical reaction:
Where A– constant factor, independent of temperature, determined by the nature of the reacting substances; e- the base of the natural logarithm; E a – activation energy; R– molar gas constant.
As follows from the Arrhenius equation, the lower the activation energy, the greater the reaction rate constant. Even a slight decrease in activation energy (for example, when adding a catalyst) 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 larger the value E and, the more noticeable the effect of temperature on the reaction rate and, therefore, the greater the temperature coefficient of the reaction rate.
The 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. The influence of a catalyst on the activation energy of a chemical reaction. Catalysis in nature, industry, technology. Biochemical catalysis. Enzymes.
1. With. 108-109; 2. With. 170-173; 3. With. 218-223; 4 . With. 197-199; 6. With. 213-222; 7. With. 197-202.; 8. With. 21-22.
Catalysis called a change in the rate of a chemical reaction under the influence of substances, the quantity and nature of which after completion of the reaction remain the same as before the reaction.
Catalyst - This is a substance that changes the rate of a chemical reaction and remains chemically unchanged after it.
Positive catalyst speeds up the reaction; 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 a reaction. Each of the intermediate processes involving a catalyst occurs with a lower activation energy than a non-catalyzed reaction.
At homogeneous catalysis the catalyst and reactants form one phase (solution). At heterogeneous catalysis the catalyst (usually a solid) and the reactants are in different phases.
During homogeneous catalysis, the catalyst forms an intermediate compound with a reagent, which reacts with a second reagent at a high speed or quickly decomposes to release a reaction product.
An example of homogeneous catalysis: the oxidation of sulfur(IV) oxide to sulfur(VI) oxide with oxygen using the nitrous method for producing sulfuric acid (here the catalyst is nitrogen(II) oxide, which easily reacts with oxygen).
In heterogeneous catalysis, the reaction occurs on the surface of the catalyst. The initial stages are the diffusion of reagent particles to the catalyst and their adsorption(i.e. absorption) by the catalyst surface. The reagent molecules interact with atoms or groups of atoms located on the surface of the catalyst, forming intermediate surface connections. The redistribution of electron density that occurs in such intermediate compounds leads to the formation of new substances that are desorbed, i.e., are removed from the surface.
The process of formation of intermediate surface compounds occurs on active centers catalyst - on surface areas characterized by a special distribution of electron density.
An example of heterogeneous catalysis: oxidation of sulfur(IV) oxide to sulfur(VI) oxide with oxygen using the contact method for producing sulfuric acid (here the catalyst can be vanadium(V) oxide with additives).
Examples of catalytic processes in industry and technology: ammonia synthesis, synthesis of nitric and sulfuric acids, cracking and reforming of oil, afterburning of products of incomplete combustion of gasoline in cars, etc.
Examples of catalytic processes in nature are numerous, since most biochemical reactions- chemical reactions occurring in living organisms are classified as catalytic reactions. The catalysts for such reactions are protein substances called enzymes. There are about 30 thousand enzymes in the human body, each of which catalyzes the passage of only one process or one type of process (for example, salivary ptyalin catalyzes the conversion of starch into sugar).
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. The influence of changes in concentration, pressure and temperature on the position of chemical equilibrium.
1. With. 109-115; 2. With. 176-182; 3 . With. 184-195, p. 207-209; 4. pp.172-176, p. 187-188; 5. With. 51-54; 8 . With. 24-31.
Chemical reactions, as a result of which starting substances are completely converted into reaction products, are called irreversible. Reactions occurring simultaneously in two opposite directions (forward and reverse) are calledreversible.
In reversible reactions, the state of the system in which the rates of the forward and reverse reactions are equal () is called state of chemical equilibrium. Chemical equilibrium is dynamic, i.e. its establishment does not mean the cessation of the reaction. In the general case, for any reversible reaction aA + bB ↔ dD + eE, regardless of its mechanism, the following relation holds:
At established equilibrium, the product of the concentrations of reaction products divided by the product of the concentrations of the starting substances for a given reaction at a given temperature is a constant value called equilibrium constant(TO).
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.
A change in 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 forward 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 displacement, of the equilibrium position.
If, during the transition from one equilibrium state to another, the concentrations of substances written on the right side of the reaction equation increase, they say that balance shifts to the right. If, during the transition from one equilibrium state to another, the concentrations of substances written on the left side of the reaction equation increase, they say that balance shifts to the left.
The direction of the shift in chemical equilibrium as a result of changes in external conditions is determined Le Chatelier's principle: If an external influence is exerted on a system in a state of chemical equilibrium, then it will favor the occurrence of whichever of the two opposite processes weakens this influence.
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 of equilibrium to the right; an increase in the concentration of the component written on the right side of the equation leads to a shift of equilibrium to the left;
When the temperature increases, the equilibrium shifts towards the endothermic reaction, and when the temperature decreases, towards the exothermic reaction;
As the pressure increases, the equilibrium shifts towards a reaction that reduces the number of molecules of gaseous substances in the system, and as the pressure decreases, towards a reaction that increases the number of molecules of gaseous substances.
Photochemical and chain reactions. Features of the course of photochemical reactions. Photochemical reactions and living nature. Unbranched and branched chemical reactions (using the example of reactions of the formation of hydrogen chloride and water from simple substances). Conditions for the initiation and termination of chains.
2. With. 173-176; 3. With. 224-226; 4. 193-196; 6. With. 207-210; 8. With. 49-50.
Photochemical reactions - These are reactions that take place under the influence of light. A photochemical reaction occurs if the reagent absorbs radiation quanta characterized by an energy quite specific for a given reaction.
In the case of some photochemical reactions, absorbing energy, the molecules of the reagent pass into an excited state, i.e. become active.
In other cases, a photochemical reaction occurs if quanta of such high energy are absorbed that chemical bonds are broken and molecules dissociate into atoms or groups of atoms.
The greater the irradiation intensity, the greater the speed of the photochemical reaction.
An example of a photochemical reaction in living nature: photosynthesis, i.e. the formation of organic cell substances by organisms due to light energy. In most organisms, photosynthesis occurs 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
The functioning of vision is also based on photochemical processes.
Chain reaction - reaction, which is a chain of elementary acts of interaction, and the possibility of each act of interaction depends on the success of the previous act.
Stages chain reaction:
The birth of a chain
Chain development,
Circuit break.
The initiation of a chain occurs when, due to an external source of energy (quanta of electromagnetic radiation, heating, electrical discharge), active particles with unpaired electrons (atoms, free radicals) are formed.
During the development of the chain, radicals interact with the original molecules, and new radicals are formed in each act of interaction.
Chain termination occurs when two radicals collide and transfer the energy released in the process to a third body (a molecule resistant to decay or the wall of a vessel). The chain can also terminate if a low-active radical is formed.
Two types chain reactions: unbranched and branched.
IN unbranched In reactions at the stage of chain development, one new radical is formed from one reacting radical.
IN branched In reactions at the stage of chain development, more than one new radical is formed from one reacting radical.
6. Factors that determine the direction of a chemical reaction. Elements of chemical thermodynamics. Concepts: phase, system, environment, macro- and microstates. Basic thermodynamic characteristics. Internal energy of the system and its change during chemical transformations. Enthalpy. The relationship between enthalpy and internal energy of a system. Standard enthalpy of a substance. Changes in enthalpy in systems during chemical transformations. Thermal effect (enthalpy) of a chemical reaction. Exo- and endothermic processes.
1. With. 89-97; 2. With. 158-163, p. 187-194; 3. With. 162-170; 4. With. 156-165; 5. With. 39-41; 6. With. 174-185; 8. With. 32-37.
Thermodynamics studies the patterns of energy exchange between the system and the external environment, the possibility, direction and limits of the spontaneous occurrence of chemical processes.
Thermodynamic system(or just system) – a body or group of interacting bodies mentally identified in space. The rest of the space outside the system is called environment(or just environment). The system is separated from the environment by a real or imaginary surface .
Homogeneous system consists of one phase, heterogeneous system– of two or more phases.
PhaseA –This is a part of the system, homogeneous at all its points in chemical composition and properties and separated from other phases of the system by an interface.
State a system is characterized by the totality of its physical and chemical properties. Macrostate is determined by the averaged parameters of the entire set of particles in the system, and microstate- parameters of each individual particle.
Independent variables that determine the macrostate of the system are called thermodynamic variables, or state parameters. Temperature is usually chosen as the state parameters T, pressure r, volume V, chemical quantity n, concentration With etc.
A physical quantity whose value depends only on the parameters of the state and does not depend on the path of transition to a given state is called state function. The functions of the state are, in particular:
U- internal energy;
N- enthalpy;
S- entropy;
G- Gibbs energy (or free energy, or isobaric-isothermal potential).
Internal energy of the system U – this 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 it is impossible to fully take into account all these components, when studying thermodynamics the system is considered change its internal energy during transition from one state ( U 1) to another ( U 2):
U 1 U 2 DU = U 2 - U 1
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 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:
In the future, we will consider the properties of such systems that are not affected by any forces other than external pressure forces.
If the process in the system occurs at a constant volume (i.e., there is no work against external pressure forces), then A = 0. Then thermal effectprocess occurring at constant volume, Q v is equal to the change in the internal energy of the system:
Q v = ΔU
Most chemical reactions encountered in everyday life occur at constant pressure ( isobaric processes). If no forces other than constant external pressure act on the system, then:
A = p(V 2 -V 1) = pDV
Therefore, in our case ( r= const):
Q р = U 2 – U 1 + p(V 2 - V 1), whence
Q p = (U 2 + pV 2) - (U 1 + pV 1)
Function U+pV, called enthalpy; it is designated by the letter N . Enthalpy is a function of state and has the dimension of energy (J).
Q p = H 2 - H 1 = DH
Thermal effect of reaction at constant pressure and temperature T is equal to the change in enthalpy of the system during the reaction. It depends on the nature of the reagents and products, their physical state, conditions ( T,r) the reaction, as well as the amount of substances involved in the reaction.
Enthalpy of reactioncall the change in enthalpy of a system in which reactants interact in quantities equal to the stoichiometric coefficients of the reaction equation.
The enthalpy of the reaction is called standard, if the reactants and reaction products are in standard states.
The standard states are:
For a solid - an individual crystalline substance at 101.32 kPa,
For a liquid substance - an individual liquid substance at 101.32 kPa,
For a gaseous substance - gas at a partial pressure of 101.32 kPa,
For a solute, a substance in solution with a molality of 1 mol/kg, and the solution is assumed to have the properties of an infinitely dilute solution.
The standard enthalpy of the reaction of formation of 1 mole of a given substance from simple substances is called standard enthalpy of formation of this substance.
Example entry: D f H o 298(CO 2) = -393.5 kJ/mol.
The standard enthalpy of formation of a simple substance located 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 one has a zero standard enthalpy of formation (for given r And T) modification.
Typically thermodynamic quantities are determined at standard conditions:
r= 101.32 kPa and T= 298 K (25 o C).
Chemical equations that specify enthalpy changes (heat effects of reactions) are called thermochemical equations. In the literature you can find two forms of writing thermochemical equations.
Thermodynamic form of writing the thermochemical equation:
C (graphite) + O 2 (g) ® CO 2 (g); DH o 298= -393.5 kJ
Thermochemical form of writing the thermochemical equation of 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 standpoint of the system, therefore, if the system releases heat, then Q<0, а энтальпия системы уменьшается (ΔH< 0).
In classical thermochemistry, thermal effects are considered from the environmental perspective, therefore, if a system releases heat, then it is assumed that Q>0.
Exothermic is a process that occurs with the release of heat (ΔH<0).
Endothermic is a process that occurs with heat absorption (ΔH>0).
The basic law of thermochemistry is Hess's law: the thermal effect of the reaction is determined only by the initial and final states of the system and does not depend on the path of transition of the system from one state to another.
Corollary of Hess's law : the standard thermal effect of a reaction is equal to the sum of the standard heats of formation of reaction products minus the sum of the standard heats of formation of starting substances, taking into account stoichiometric coefficients:
DН about 298 (r-tions) = åD f Н about 298 (cont.) – åD f Н about 298 (original)
7. The concept of entropy. Changes in entropy during phase transformations and chemical processes. The concept of the isobaric-isothermal potential of a system (Gibbs energy, free energy). The relationship between the magnitude of the change in the Gibbs energy and the magnitude of the change in enthalpy and entropy of the reaction (basic thermodynamic relationship). Thermodynamic analysis of the possibility and conditions of chemical reactions. Features of the flow of chemical processes in living organisms.
1. With. 97-102; 2. With. 189-196; 3. With. 170-183; 4. With. 165-171; 5. With. 42-44; 6. With. 186-197; 8. With. 37-46.
Entropy S- this is a quantity proportional to the logarithm of the number of equally probable microstates through which a given macrostate can be realized:
The unit of entropy is J/mol·K.
Entropy is a quantitative measure of the degree of disorder of 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 especially particles in the gaseous state. On the contrary, all processes as a result of which the order 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 the tables of thermodynamic characteristics of individual substances provide data for S 0, and not for Δ S 0.
The standard entropy of a simple substance, in contrast to the enthalpy of formation of a simple substance, is not zero.
For entropy, a statement similar to that discussed above for DH: the change in the entropy of a system as a result of a chemical reaction (DS) is equal to the sum of the entropies of the reaction products minus the sum of the entropies of the starting substances. As with the calculation of enthalpy, the summation is carried out taking into account the stoichiometric coefficients.
The direction in which a chemical reaction spontaneously occurs 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) a tendency to achieve the most probable state, i.e. a state that can be realized in the largest number of equally probable ways (microstates):
Δ H → min,Δ S → max
The state function, which simultaneously reflects the influence of both of the above-mentioned trends on the direction of the flow of chemical processes, is Gibbs energy (free energy , or isobaric-isothermal potential) , related to enthalpy and entropy by the relation
G = H - TS,
Where T- absolute temperature.
As can be seen, the Gibbs energy has the same dimension as enthalpy and is therefore 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 equal to:
As in case D H and D S, Gibbs energy change D G as a result of a chemical reaction(Gibbs energy of reaction) equal to the sum of the Gibbs energies of the formation of reaction products minus the sum of the Gibbs energies of the formation of the starting substances; the summation is carried out taking into account the number of moles of substances participating in the reaction.
The Gibbs energy of formation of a substance is referred to 1 mole of this substance and is usually expressed in kJ/mol; while D G 0 of the formation of the most stable modification of a simple substance is taken equal to zero.
At constant temperature and pressure, chemical reactions can spontaneously proceed only in a direction in which the Gibbs energy of the system decreases ( D G<0).This is a condition for the fundamental possibility of carrying out this process.
The table below shows the possibility and conditions for the reaction to occur with various combinations of signs D N and D S.
By sign D G one can judge the possibility (impossibility) spontaneous flow individual process. If you put pressure on the system impact, then it is possible to carry out a transition from one substance to another, characterized by an increase in free energy (D G>0). For example, in the cells of living organisms reactions occur to form complex organic compounds; The driving force behind such processes is solar radiation and oxidation reactions in the cell.
Thermodynamics – the science of converting one form of energy into another based on the law of conservation of energy. Thermodynamics establishes the direction of the spontaneous flow of chemical reactions under given conditions. During chemical reactions, bonds in the starting substances are broken and new bonds are formed in the final products. The sum of binding energies after the reaction is not equal to the sum of binding energies before the reaction, i.e. the occurrence of a chemical reaction is accompanied by the release or absorption of energy, and its forms are different.
Thermochemistry is a branch of thermodynamics devoted to the study of the thermal effects of reactions. The thermal effect of a reaction measured at constant temperature and pressure is called enthalpy of reaction and are expressed in joules (J) and kilojoules (kJ).
For exothermic reactions, for endothermic reactions -. The enthalpy of formation of 1 mol of a given substance from simple substances, measured at a temperature of 298 K (25 ° C) and a pressure of 101.825 kPa (1 atm), is called standard (kJ/mol). The enthalpies of simple substances are conventionally assumed to be zero.
Thermochemical calculations are based on Hess’s law: t The thermal effect of a reaction depends only on the nature and physical state of the starting materials and final products, but does not depend on the transition path. Often in thermochemical calculations a corollary from Hess’s law is used: thermal effect of a chemical reaction equal to the sum of the heats of formation reaction products minus the sum of the heats of formation of the starting substances, taking into account the coefficients in front of the formulas of these substances in the reaction equation:
Thermochemical equations indicate the enthalpy of a chemical reaction. At the same time, the formula of each substance indicates its physical state: gaseous (g), liquid (l), solid crystalline (c).
In thermochemical equations, the thermal effects of reactions are given per 1 mole of the starting or final substance. Therefore, fractional odds are allowed here. In chemical reactions, the dialectical law of unity and struggle of opposites is manifested. On the one hand, the system strives for ordering (aggregation) - reducing N, and on the other hand – to disorder (disaggregation). The first trend increases with decreasing temperature, and the second - with increasing temperature. The tendency towards disorder is characterized by a quantity called entropy S[J/(mol. K)]. It is a measure of the disorder of the system. Entropy is proportional to the amount of matter and increases with increasing movement of particles during heating, evaporation, melting, gas expansion, weakening or breaking of bonds between atoms, etc. Processes associated with the orderliness of the system: condensation, crystallization, compression, strengthening of bonds, polymerization, etc. – lead to a decrease in entropy. Entropy is a function of state, i.e.
The general driving force of the process consists of two forces: the desire for order and the desire for disorder. With p = const and T = const, the overall driving force of the process can be represented as follows:
Gibbs energy, or isobaric-isothermal potential, also obeys a corollary of Hess's law:
Processes proceed spontaneously in the direction of decreasing any potential and, in particular, in the direction of decreasing . At equilibrium, the temperature at which the equilibrium reaction begins is equal to:
Table 5
Standard enthalpies of formation , entropy and Gibbs energy formation some substances at 298 K (25°C)
| Substance | , kJ/mol | , J/mol | , kJ/mol |
| CaO(k) | -635,5 | 39,7 | -604,2 |
| CaCO 3 (k) | -1207,0 | 88,7 | -1127,7 |
| Ca(OH) 2 (k) | -986,6 | 76,1 | -896,8 |
| H 2 O (l) | -285,8 | 70,1 | -237,3 |
| H2O (g) | -241,8 | 188,7 | -228,6 |
| Na 2 O (k) | -430,6 | 71,1 | -376,6 |
| NaOH (k) | -426,6 | 64,18 | -377,0 |
| H2S (g) | -21,0 | 205,7 | -33,8 |
| SO2 (g) | -296,9 | 248,1 | -300,2 |
| SO 3 (g) | -395,8 | 256,7 | -371,2 |
| C 6 H 12 O 6 (k) | -1273,0 | - | -919,5 |
| C 2 H 5 OH (l) | -277,6 | 160,7 | -174,8 |
| CO 2 (g) | -393,5 | 213,7 | -394,4 |
| CO(g) | -110,5 | 197,5 | -137,1 |
| C 2 H 4 (g) | 52,3 | 219,4 | 68,1 |
| CH 4 (g) | -74,9 | 186,2 | -50,8 |
| Fe 2 O 3 (k) | -822,2 | 87,4 | -740,3 |
| FeO(k) | -264,8 | 60,8 | -244,3 |
| Fe 3 O 4 (k) | -1117,1 | 146,2 | -1014,2 |
| CS 2 (g) | 115,3 | 65,1 | 237,8 |
| P 2 O 5 (k) | -1492 | 114,5 | -1348,8 |
| NH 4 Cl (k) | -315,39 | 94,56 | -343,64 |
| HCl (g) | -92,3 | 186,8 | -95,2 |
| NH 3 (g) | -46,2 | 192,6 | -16,7 |
| N2O (g) | 82,0 | 219,9 | 104,1 |
| NO (g) | 90,3 | 210,6 | 86,6 |
| NO 2 (g) | 33,5 | 240,2 | 51,5 |
| N 2 O 4 (g) | 9,6 | 303,8 | 98,4 |
| CuO(k) | -162,0 | 42,6 | -129,9 |
| H2(g) | 130,5 | ||
| C (graphite) | 5,7 | ||
| O 2 (g) | 205,0 | ||
| N 2 (g) | 181,5 | ||
| Fe(k) | 27,15 | ||
| Cl 2 (g) | 222,9 | ||
| KNO 3 (k) | -429,71 | 132,93 | -393,13 |
| KNO 2 (k) | -370,28 | 117,15 | -281,58 |
| K 2 O (k) | -361,5 | 87,0 | -193,3 |
| ZnO(k) | -350,6 | 43,6 | -320,7 |
| Al 2 O 3 (k) | -1676,0 | 50,9 | -1582,0 |
| PCl 5 (g) | -369,45 | 362,9 | -324,55 |
| PCl 3 (g) | -277,0 | 311,7 | -286,27 |
| H 2 O 2 (l) | -187,36 | 105,86 | -117,57 |
Reaction speed determined by the nature and concentration of the reacting substances and depends on temperature and catalyst.
Law of mass action: At constant temperature, the rate of a chemical reaction is proportional to the product of the concentration of the reactants to the power of their stoichiometric coefficients.
For the reaction aA + bB = cC + dD, the rate of the direct reaction is:
,
reverse reaction speed: 
, where are the concentrations of dissolved or gaseous compounds, mol/l;
a, b, c, d – stoichiometric coefficients in the equation;
K is the rate constant.
The expression for the reaction rate does not include solid concentrations.
The effect of temperature on the reaction rate is described by Van't Hoff's rule: for every 10 degrees heated, the reaction rate increases by 2-4 times.
Reaction rate at temperatures t 1 and t 2;
Temperature coefficient of reaction.
Most chemical reactions are reversible:
aA + bB cC + dD
the ratio of rate constants is a constant quantity called equilibrium constant
K p = const at T = const.
Le-Chatelier's principle: If a system in a state of chemical equilibrium is subject to any impact (change in temperature, pressure or concentration), the system will react in such a way as to reduce the applied impact:
a) when the temperature in equilibrium systems increases, the equilibrium shifts towards the endothermic reaction, and when the temperature decreases, towards the exothermic reaction;
b) when the pressure increases, the equilibrium shifts towards smaller volumes, and when the pressure decreases, towards larger volumes;
c) as the concentration increases, the equilibrium shifts towards its decrease.
Example 1. Determine the standard enthalpy change of the reaction:
Is this reaction exo- or endothermic?
Solution: The standard enthalpy change of a chemical reaction is equal to the sum of the standard enthalpies of formation of the reaction products minus the sum of the standard enthalpies of formation of the starting substances
With each summation, the number of moles of substances participating in the reaction should be taken into account in accordance with the reaction equation. The standard enthalpies of formation of simple substances are zero:
According to the tabular data:
Reactions that are accompanied by the release of heat are called exothermic, and those that are accompanied by the absorption of heat are called endothermic. At constant temperature and pressure, the change in enthalpy of a chemical reaction is equal in magnitude, but opposite in sign to its thermal effect. Since the standard enthalpy change for a given chemical reaction is , we conclude that this reaction is exothermic.
Example 2. The reduction reaction of Fe 2 O 3 with hydrogen proceeds according to the equation:
Fe 2 O 3 (K) + 3H 2 (G) = 2Fe (K) + 3H 2 O (G)
Is this reaction possible under standard conditions?
Solution: To answer this question of the problem, you need to calculate the standard change in the Gibbs energy of the reaction. Under standard conditions:
The summation is carried out taking into account the number of models involved in the reaction of substances; the formation of the most stable modification of a simple substance is taken equal to zero.
Taking into account the above
According to the tabular data:
Spontaneously occurring processes are decreasing. If< 0, процесс принципиально осуществим, если >0, the process cannot proceed spontaneously.
Therefore, this reaction is impossible under standard conditions.
Example 3. Write expressions for the law of mass action for reactions:
a) 2NO (G) + Cl 2 (G) = 2NOCl (G)
b) CaCO 3 (K) = CaO (K) + CO 2 (G)
Solution: According to the law of mass action, the reaction rate is directly proportional to the product of the concentrations of the reacting substances in powers equal to the stoichiometric coefficients:
a) V = k 2.
b) Since calcium carbonate is a solid whose concentration does not change during the reaction, the desired expression will look like:
V = k, i.e. in this case, the reaction rate at a certain temperature is constant.
Example 4. The endothermic reaction of phosphorus pentachloride decomposition proceeds according to the equation:
PCl 5(G) = PCl 3(G) + Cl 2(G); 
How to change: a) temperature; b) pressure; c) concentration to shift the equilibrium towards the direct reaction - decomposition of PCl 5? Write a mathematical expression for the rates of forward and reverse reactions, as well as the equilibrium constant.
Solution: A displacement or shift in chemical equilibrium is a change in the equilibrium concentrations of reactants as a result of a change in one of the reaction conditions.
A shift in chemical equilibrium is subject to Le Chatelier's principle, according to which a change in one of the conditions under which a system is in equilibrium causes a shift in the equilibrium in the direction of the reaction that counteracts the resulting change.
a) Since the decomposition reaction of PCl 5 is endothermic, to shift the equilibrium towards the direct reaction, the temperature must be increased.
b) Since in this system the decomposition of PCl 5 leads to an increase in volume (two gaseous molecules are formed from one gas molecule), then to shift the equilibrium towards the direct reaction it is necessary to reduce the pressure.
c) A shift in equilibrium in the indicated direction can be achieved either by increasing the concentration of PCl 5 or by decreasing the concentration of PCl 3 or Cl 2 .
According to the law of mass action, the rates of direct (V 1) and reverse (V 2) reactions are expressed by the equations:
V 2 = k
The equilibrium constant of this reaction is expressed by the equation:
Test tasks:
81 - 100. a) calculate the standard change in enthalpy of the forward reaction and determine whether the reaction is exo- or endothermic;
b) determine the change in the Gibbs energy of the direct reaction and draw a conclusion about the possibility of its implementation under standard conditions;
c) write a mathematical expression for the rate of forward and reverse reactions, as well as the equilibrium constant;
d) how should the conditions be changed to shift the equilibrium of the process to the right?
81. CH 4 (g) + CO 2 (g) = 2CO (g) + 2H 2 (g)
82. FeO (K) + CO (g) = Fe (K) + CO 2 (g)
83. C 2 H 4 (g) + O 2 (g) = CO 2 (g) + H 2 O (g)
84. N 2(g) + 3H 2(g) = 2NH 3(g)
85. H 2 O (g) + CO (g) = CO 2 (g) + H 2 (g)
86. 4HCl (g) + O 2 (g) = 2H 2 O (g) + 2Cl 2 (g)
87. Fe 2 O 3 (K) + 3H 2 (g) = 2Fe (K) + 3H 2 O (g)
88. 2SO 2 (g) + O 2 (g) = 2SO 3 (g)
89. PCl 5(g) = PCl 3(g) + Cl 2(g)
90. CO 2 (g) + C (graphite) = 2CO (g)
91. 2H 2 S (g) + 3O 2 (g) = 2SO 2 (g) + H 2 O (g)
92. Fe 2 O 3 (K) + CO (g) = 2FeO (K) + CO 2 (g)
93. 4NH 3 (g) + 5O 2 (g) = 4NO (g) + 6H 2 O (g)
94. NH 4 Cl (K) = NH 3 (g) + HCl (g)
95. CH 4 (g) + 2O 2 (g) = CO 2 (g) + 2H 2 O (g)
96. CS 2(g) + 3O 2(g) = CO 2(g) + 2SO 2(g)
97. 4HCl (g) + O 2 (g) = 2Cl 2 (g) + 2H 2 O (g)
98. 2NO (g) + O 2 (g) = N 2 O 4 (g)
99. NH 3 (g) + HCl (g) = NH 4 Cl (K)
100. CS 2(g) + 3O 2(g) = 2Cl 2(g) + 2SO 2(g)
Topic 6: Solutions. Methods of expressing the concentration of solutions
Solutions are homogeneous systems consisting of a solvent, solutes and possible products of their interaction. The concentration of a solution is the content of a dissolved substance in a certain mass or known volume of a solution or solvent.
Ways to express the concentration of solutions:
Mass fraction() shows the number of grams of solute in 100 g of solution:
Where T– mass of solute (g), T 1 – mass of solution (g).
Molar concentration shows the number of moles of solute contained in 1 liter of solution:
where M is the molar mass of the substance (g/mol), V is the volume of solution (l).
Molal concentration shows the number of moles of solute contained in 1000 g of solvent: n 101-120. Find the mass fraction, molar concentration, molal concentration for the following solutions:
| Option | Substance (x) | Mass of substance (x) | Water volume | Solution density |
| CuSO4 | 320 g | 10 l | 1,019 | |
| NaCl | 0.6 g | 50 ml | 1,071 | |
| H2SO4 | 2 g | 100 ml | 1,012 | |
| Na2SO4 | 13 g | 100 ml | 1,111 | |
| HNO3 | 12.6 g | 100 ml | 1,066 | |
| HCl | 3.6 kg | 10 kg | 1,098 | |
| NaOH | 8 g | 200 g | 1,043 | |
| MgCl2 | 190 g | 810 g | 1,037 | |
| KOH | 224 g | 776 g | 1,206 | |
| CuCl2 | 13.5 g | 800 ml | 1,012 | |
| HCl | 10.8 g | 200 g | 1,149 | |
| CuSO4 | 8 g | 200 ml | 1,040 | |
| NaCl | 6.1 g | 600 ml | 1,005 | |
| Na2SO3 | 4.2 g | 500 ml | 1,082 | |
| H2SO4 | 98 g | 1000 ml | 1,066 | |
| ZnCl2 | 13.6 g | 100 ml | 1,052 | |
| H3PO4 | 9.8 g | 1000 ml | 1,012 | |
| Ba(OH)2 | 100 g | 900 g | 1,085 | |
| H3PO4 | 29.4 g | 600 ml | 1,023 | |
| NaOH | 28 g | 72 g | 1,309 |
Lecture 1 Chemical thermodynamics. Chemical kinetics and catalysis PLAN 1. Basic concepts of thermodynamics. 2. Thermochemistry. 3. Chemical equilibrium. 4. Rate of chemical reactions. 5. The influence of temperature on the rate of reactions. 6. The phenomenon of catalysis. Prepared by: Ph.D., Associate Professor. Ivanets L.M., as. 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 transition of energy in the form of heat and work. The great practical importance of thermodynamics is that it allows one to calculate the thermal effects of a reaction, to indicate in advance the possibility or impossibility of carrying out a reaction, as well as the conditions for its occurrence.
Internal energy Internal energy is the kinetic energy of all particles of the system (molecules, atoms, electrons) and the potential energy of their interactions, in addition to the kinetic and potential energy of the system as a whole. Internal energy is a function of state, i.e. its change is determined by the given initial and final states of the system and does not depend on the process path: 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 type to another in equivalent quantities. A perpetual motion machine of the first kind, that is, a periodically operating machine that produces work without wasting energy, is impossible. Q = 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 path of the reaction, 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 Starting substances, reaction products H 1 = H 2 + H 3 + H 4 H 1 = H 2 + H 3 + H 4
The second law of thermodynamics, like the first, is the result of centuries of human experience. There are different 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’s 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 periodic machine that only cools the thermal reservoir and does 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 – function of the isochoric-isentropic process: dU = TdS – pdV. For an arbitrary process: U 0 Н – function of an isobaric-isentropic process: dН = TdS + Vdp For an arbitrary process: Н 0 S – function of an isolated system For an arbitrary process: S 0 For an arbitrary process: S 0 F – function of an isochoric-isothermal process dF = dU – TdS. For an arbitrary process: F 0 G – 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 Reverse reaction: A B Parallel: B A C Sequential: ABC Conjugate: A D Conjugate: A D C B E B E
The influence of temperature on the rate of reactions The influence of temperature on the rate of enzymatic reactions t t
Van't Hoff comparison: Calculation of the shelf life of drugs using the Van't Hoff "accelerated aging" method: at t 2 t 1 Temperature rate coefficient:
