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
O2 (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 the 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

1. The rate of chemical reactions. Definition of the concept. Factors affecting the rate of a chemical reaction: concentration of the reagent, 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 with 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 to 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;

- 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.
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,

i.e., 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 react 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 standpoint 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 ln W

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 made 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.

Solving problems for the section

The topic “Chemical thermodynamics and kinetics,” which involves the study of conditions affecting the rate of a chemical reaction, appears twice in the school chemistry course – in the 9th and 11th grades. However, this particular topic is one of the most difficult and quite complex not only for understanding by the “average” student, but even for presentation by some teachers, especially non-specialists working in rural areas, for whom chemistry is an additional subject, taking into account the hours of which the teacher accumulates rate, and therefore hope for a more or less decent salary.
In conditions of a sharp decrease in the number of students in rural schools, for well-known reasons, the teacher is forced to be a generalist. After attending 2-3 courses, he begins teaching subjects that are often very far from his main specialty.
This development is aimed primarily at beginning teachers and subject specialists who are forced to teach chemistry in a market economy. The material contains tasks on finding the rates of heterogeneous and homogeneous reactions and increasing the reaction rate with increasing temperature. Despite the fact that these problems are based on school material, although difficult for the “average” student to master, it is advisable to solve several of them in a chemistry lesson in
11th grade, and offer the rest at a club or elective lesson to students who plan to connect their future destiny with chemistry.
In addition to problems analyzed in detail and provided with answers, this development contains theoretical material that will help a chemistry teacher, primarily a non-specialist, understand the essence of this complex topic in a general chemistry course.
Based on the proposed material, you can create your own version of a lesson-lecture, depending on the abilities of the students in the class, and you can use the proposed theoretical part when studying this topic in both the 9th and 11th grades.
Finally, the material contained in this development will be useful for a graduate preparing to enter a university, including one in which chemistry is a major subject, to analyze independently.

Theoretical part on the topic
"Chemical thermodynamics and kinetics"

Conditions affecting the rate of a chemical reaction

1. The rate of a chemical reaction depends on the nature of the reacting substances.

EXAMPLES.

Metallic sodium, which is alkaline in nature, reacts violently with water, releasing a large amount of heat, in contrast to zinc, which is amphoteric in nature, which reacts with water slowly and when heated:

Powdered iron reacts more vigorously with strong mineral hydrochloric acid than with weak organic acetic acid:

2. The rate of a chemical reaction depends on the concentration of the reactants, whether in a dissolved or gaseous state.

EXAMPLES.

In pure oxygen, sulfur burns more energetically than in air:

Powdered magnesium reacts more vigorously with a 30% solution of hydrochloric acid than with a 1% solution:

3. The rate of a chemical reaction is directly proportional to the surface area of the reacting substances in the solid state of aggregation.

EXAMPLES.

A piece of charcoal (carbon) is very difficult to light with a match, but charcoal dust burns explosively:

C + O 2 = CO 2.

Aluminum in the form of granules does not react quantitatively with the iodine crystal, but crushed iodine combines vigorously with aluminum in the form of powder:

4. The rate of a chemical reaction depends on the temperature at which the process occurs.

EXAMPLE

For every 10 °C increase in temperature, the rate of most chemical reactions increases by 2–4 times. A specific increase in the rate of a chemical reaction is determined by a specific temperature coefficient (gamma).

Let's calculate how many times the reaction rate will increase:

2NO + O 2 = 2NO 2,

if the temperature coefficient is 3 and the process temperature has increased from 10 °C to 50 °C.

The temperature change is:

t= 50 °C – 10 °C = 40 °C.

We use the formula:

where is the rate of a chemical reaction at elevated temperature, is the rate of a chemical reaction at the initial temperature.

Consequently, the rate of a chemical reaction when the temperature increases from 10 °C to 50 °C will increase by 81 times.

5. The rate of a chemical reaction depends on the presence of certain substances.

Catalyst is a substance that accelerates the course of a chemical reaction, but is not consumed during the reaction. A catalyst lowers the activation barrier of a chemical reaction.

Inhibitor is a substance that slows down the progress of a chemical reaction, but is not consumed during the reaction process.

EXAMPLES.

The catalyst that accelerates this chemical reaction is manganese(IV) oxide.

The catalyst that accelerates this chemical reaction is red phosphorus.

An inhibitor that slows down the progress of this chemical reaction is an organic substance - methenamine (hexamethylenetetramine).

The rate of a homogeneous chemical reaction is measured by the number of moles of the substance that reacted or formed as a result of the reaction per unit time per unit volume:

where homog is the rate of a chemical reaction in a homogeneous system, is the number of moles of one of the substances that entered into the reaction or one of the substances formed as a result of the reaction, V- volume,
t– time, – change in the number of moles of a substance during the reaction t.

Since the ratio of the number of moles of a substance to the volume of the system represents the concentration With, That

Hence:

The rate of a homogeneous chemical reaction is measured in mol/(l s).

Taking this into account, the following definition can be given:

the rate of a homogeneous chemical reaction is equal to the change in the concentration of one of the substances that entered into the reaction or one of the substances formed as a result of the reaction per unit time.

If a reaction occurs between substances in a heterogeneous system, then the reacting substances do not come into contact with each other throughout the entire volume, but only on the surface of the solid. For example, when a piece of crystalline sulfur burns, oxygen molecules react only with those sulfur atoms that are on the surface of the piece. When a piece of sulfur is crushed, the reacting surface area increases and the rate of sulfur combustion increases.

In this regard, the definition of the rate of a heterogeneous chemical reaction is as follows:

the rate of a heterogeneous chemical reaction is measured by the number of moles of the substance that reacted or formed as a result of the reaction per unit time on a unit surface:

Where S– surface area.

The rate of a heterogeneous chemical reaction is measured in mol/(cm 2 s).

Tasks on the topic
"Chemical thermodynamics and kinetics"

1. 4 moles of nitrogen(II) oxide and excess oxygen were introduced into the vessel for chemical reactions. After 10 s, the amount of nitrogen oxide(II) substance turned out to be 1.5 mol. Find the rate of this chemical reaction if it is known that the volume of the vessel is 50 liters.

2. The amount of methane substance in the vessel for carrying out chemical reactions is 7 mol. Excess oxygen was introduced into the vessel and the mixture was exploded. It was experimentally established that after 5 s the amount of methane substance decreased by 2 times. Find the rate of this chemical reaction if it is known that the volume of the vessel is 20 liters.

3. The initial concentration of hydrogen sulfide in the gas combustion vessel was 3.5 mol/l. Excess oxygen was introduced into the vessel and the mixture was exploded. After 15 s, the hydrogen sulfide concentration was 1.5 mol/l. Find the rate of this chemical reaction.

4. The initial concentration of ethane in the gas combustion vessel was 5 mol/L. Excess oxygen was introduced into the vessel and the mixture was exploded. After 12 s, the ethane concentration was 1.4 mol/L. Find the rate of this chemical reaction.

5. The initial concentration of ammonia in the gas combustion vessel was 4 mol/l. Excess oxygen was introduced into the vessel and the mixture was exploded. After 3 s, the ammonia concentration was 1 mol/l. Find the rate of this chemical reaction.

6. The initial concentration of carbon monoxide (II) in the gas combustion vessel was 6 mol/l. Excess oxygen was introduced into the vessel and the mixture was exploded. After 5 s, the concentration of carbon(II) monoxide was halved. Find the rate of this chemical reaction.

7. A piece of sulfur with a reacting surface area of 7 cm2 was burned in oxygen to form sulfur(IV) oxide. In 10 s, the amount of sulfur substance decreased from 3 mol to 1 mol. Find the rate of this chemical reaction.

8. A piece of carbon with a reacting surface area of 10 cm 2 was burned in oxygen to form carbon monoxide (IV). In 15 s, the amount of carbon substance decreased from 5 mol to 1.5 mol. Find the rate of this chemical reaction.

9. A cube of magnesium with a total reacting surface area of 15 cm 2 and the amount of substance
6 moles burned in excess oxygen. Moreover, 7 s after the start of the reaction, the amount of magnesium substance turned out to be equal to 2 mol. Find the rate of this chemical reaction.

10. A calcium bar with a total reacting surface area of 12 cm 2 and an amount of substance of 7 mol was burned in excess oxygen. Moreover, 10 s after the start of the reaction, the amount of calcium substance turned out to be 2 times less. Find the rate of this chemical reaction.

Solutions and Answers

1 (NO) = 4 mol,

O 2 – excess,

t 2 = 10 s,

t 1 = 0 s,

2 (NO) = 1.5 mol,

Find:

Solution

2NO + O 2 = 2NO 2.

Using the formula:

R-tions = (4 – 1.5)/(50 (10 – 0)) = 0.005 mol/(l s).

Answer. r-tion = 0.005 mol/(l s).

1 (CH 4) = 7 mol,

O 2 – excess,

t 2 = 5 s,

t 1 = 0 s,

2 (CH 4) = 3.5 mol,

Find:

Solution

CH 4 + 2O 2 = CO 2 + 2H 2 O.

Using the formula:

Let's find the rate of this chemical reaction:

R-tions = (7 – 3.5)/(20 (5 – 0)) = 0.035 mol/(l s).

Answer. r-tion = 0.035 mol/(l s).

s 1 (H 2 S) = 3.5 mol/l,

O 2 – excess,

t 2 = 15 s,

t 1 = 0 s,

With 2 (H 2 S) = 1.5 mol/l.

Find:

Solution

2H 2 S + 3O 2 = 2SO 2 + 2H 2 O.

Using the formula:

Let's find the rate of this chemical reaction:

R-tions = (3.5 – 1.5)/(15 – 0) = 0.133 mol/(l s).

Answer. r-tion = 0.133 mol/(l s).

c 1 (C 2 H 6) = 5 mol/l,

O 2 – excess,

t 2 = 12 s,

t 1 = 0 s,

c 2 (C 2 H 6) = 1.4 mol/l.

Find:

Solution

2C 2 H 6 + 7O 2 = 4CO 2 + 6H 2 O.

Let's find the rate of this chemical reaction:

R-tions = (6 – 2)/(15 (7 – 0)) = 0.0381 mol/(cm 2 s).

Answer. r-tion = 0.0381 mol/(cm 2 s).

10. Answer. r-tion = 0.0292 mol/(cm 2 s).

Literature

Glinka N.L. General Chemistry, 27th ed. Ed. V.A. Rabinovich. L.: Chemistry, 1988; Akhmetov N.S. General and inorganic chemistry. M.: Higher. school, 1981; Zaitsev O.S. General chemistry. M.: Higher. shk, 1983; Karapetyants M.Kh., Drakin S.I. General and inorganic chemistry. M.: Higher. school, 1981; Korolkov D.V. Fundamentals of inorganic chemistry. M.: Education, 1982; Nekrasov B.V. Fundamentals of general chemistry. 3rd ed., M.: Khimiya, 1973; Novikov G.I. Introduction to inorganic chemistry. Part 1, 2. Minsk: Higher. school, 1973–1974; Shchukarev S.A.. Inorganic chemistry. T. 1, 2. M.: Vyssh. school, 1970–1974; Schröter W., Lautenschläger K.-H., Bibrak H. et al. Chemistry. Reference ed. Per. with him. M.: Khimiya, 1989; Feldman F.G., Rudzitis G.E. Chemistry-9. Textbook for 9th grade of secondary school. M.: Education, 1990; Feldman F.G., Rudzitis G.E. Chemistry-9. Textbook for 9th grade of secondary school. M.: Education, 1992.

The rate of chemical reactions. Definition of the concept. Factors affecting the rate of a chemical reaction: concentration of the reagent, 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;

- concentration of the reagent;

- 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.

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.

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.

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 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,

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.

During the development of the chain, radicals interact with the original molecules, and new radicals are formed in each act of interaction.

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.

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.

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 made 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.