V. Chemical kinetics and chemical equilibrium

Purpose of the work: to study the effect of temperature on the rate of concentration reaction to a shift in chemical equilibrium. Theoretical rationale: The rate of a chemical reaction is the amount of substance that reacts or is formed as a result of a reaction per unit time per unit volume for homogeneous reactions or per unit interface surface for heterogeneous reactions. If within a period of time...

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"Ufa State Petroleum Technical University"

Department of "General and Analytical Chemistry"

REPORT

For laboratory work No. 1

"Chemical kinetics and equilibrium"

Student of the _______________ group E.V. Beletskova

BTS-14-01

Associate Professor _______________S.B. Denisova

2014

Goal of the work : study of the influence of temperature on the reaction rate, concentration on the shift in chemical equilibrium.

Theoretical background:

Speed of chemical reactionis the amount of a substance that reacts or is formed as a result of a reaction per unit time per unit volume (for homogeneous reactions) or per unit interface surface (for heterogeneous reactions).

If over a period of time ∆τ = τ 2  τ 1 the concentration of one of the substances participating in the reaction decreases by ∆C = C 2 - C 1 , then the average rate of a chemical reaction over a specified period of time is equal to

V value expresses the rate of a chemical process over a certain period of time. Therefore, the smaller ∆τ, the closer the average speed will be to the true one.

The rate of a chemical reaction depends on the following factors:

nature and concentration of reacting substances;
reaction system temperature;
presence of a catalyst;
pressure,
the magnitude of the phase interface and the mixing rate of the system (for heterogeneous reactions);
type of solvent.

Effect of reagent concentrations. The rate of a reaction is proportional to the number of collisions of molecules of the reacting substances. The number of collisions, in turn, is greater, the higher the concentration of each of the starting substances.

A general formulation of the effect of concentration on the rate of a chemical reaction is given bylaw of mass action(1867, Guldberg, Waage, Beketov).

At a constant temperature, the rate of a chemical reaction is proportional to the product of the concentrations of the reacting substances, taken in powers of their equalizing (stoichiometric) coefficients.

For the reaction aA + bB = cC V = K[A]a[B]v,

where K proportionality coefficient or speed constant;

 reagent concentration in mol/l.

If [A] = 1 mol/l, [B] = 1 mol/l, then V=K , hence the physical meaning

rate constants K: the rate constant is equal to the reaction rate at concentrations of reactants equal to unity.

The effect of temperature on the reaction rate. As the temperature increases, the frequency of collisions of reacting molecules increases, and therefore the reaction rate increases.

The quantitative effect of temperature on the rate of homogeneous reactions can be expressed by Van't Hoff's rule.

In accordance with Van't Hoff's rule, when the temperature increases (decreases) by 10 degrees, the rate of a chemical reaction increases (decreases) by 2-4 times:

or ,

where V (t 2) and V (t 1) the rate of chemical reaction at appropriate temperatures;τ (t 2) and τ (t 1) duration of the chemical reaction at the appropriate temperatures;γ Van't Hoff temperature coefficient, which can take a numerical value in the range of 2-4.

Activation energy. The excess energy that molecules must have in order for their collision to lead to the formation of a new substance is called the activation energy of a given reaction (expressed in kJ/mol). One of the methods of activation is to increase the temperature: with increasing temperature, the number of active particles increases greatly, due to which the reaction rate sharply increases.

The dependence of the reaction rate on temperature is expressed by the Arrhenius equation:

where K is the rate constant of the chemical reaction; E a activation energy;

R universal gas constant; A constant; exp base of natural logarithms.

The activation energy can be determined if two values of the rate constant K are known 1 and K 2 at temperature respectively T 1 and T 2 , according to the following formula:

Chemical balance.

All chemical reactions can be divided into two groups: irreversible and reversible. Irreversible reactions proceed to completion until one of the reactants is completely consumed, i.e. flow in only one direction. Reversible reactions do not proceed to completion. In a reversible reaction, none of the reactants are completely consumed. A reversible reaction can occur in both the forward and reverse directions.

Chemical equilibrium is a state of a system in which the rates of forward and reverse reactions are equal.

For a reversible reaction

m A + n B ⇄ p C + q D

the chemical equilibrium constant is

In reversible chemical reactions, equilibrium is established at the moment when the ratio of the product of concentrations of products raised to powers equal to the stoichiometric coefficients to the product of concentrations of starting substances, also raised to the corresponding powers, is equal to some constant value called the chemical equilibrium constant.

The chemical equilibrium constant depends on the nature of the reactants and on temperature. The concentrations at which equilibrium is established are called equilibrium. A change in external conditions (concentration, temperature, pressure) causes a shift in the chemical equilibrium in the system and its transition to a new equilibrium state.

Such a transition of a reaction system from one state to another is called a displacement (or shift) of chemical equilibrium.

The direction of the shift in chemical equilibrium is determined by Le Chatelier’s principle:If any external influence is applied to a system that is in a state of chemical equilibrium (change concentration, temperature, pressure), then processes spontaneously arise in this system that tend to weaken the effect produced.

An increase in the concentration of one of the starting reagents shifts the equilibrium to the right (the direct reaction is enhanced); An increase in the concentration of reaction products shifts the equilibrium to the left (the reverse reaction intensifies).

If a reaction proceeds with an increase in the number of gas molecules (i.e., on the right side of the reaction equation, the total number of gas molecules is greater than the number of molecules of gaseous substances on the left side), then an increase in pressure prevents the reaction, and a decrease in pressure favors the reaction.

As the temperature increases, the equilibrium shifts towards the endothermic reaction, and as the temperature decreases, towards the exothermic reaction.

The catalyst changes the rate of both forward and reverse reactions by the same number of times. Therefore, the catalyst does not cause a shift in equilibrium, but only shortens or increases the time required to achieve equilibrium.

Experiment No. 1 Dependence of the rate of a homogeneous reaction on the concentration of the initial reagents.

Instruments, equipment: test tubes, stopwatch, solutions of sodium thiosulfate ( III ), div. sulfuric acid (1M), water.
Methodology: This dependence can be studied using the classic example of a homogeneous reaction of sodium thiosulfate with sulfuric acid, proceeding according to the equation

Na 2 S 2 O 3 + H 2 SO 4 = Na 2 SO 4 + S↓ + SO 2 + H 2 O.

At first, sulfur forms a colloidal solution with water (barely perceptible turbidity). It is necessary to measure the time from the moment of draining until a barely noticeable turbidity appears with a stopwatch. Knowing the reaction time (in seconds), you can determine the relative speed of the reaction, i.e. reciprocal of time: .

For the experiment, you should prepare three dry, clean test tubes and number them. Add 4 drops of sodium thiosulfate solution and 8 drops of water to the first; in the second 8 drops of sodium thiosulfate and 4 drops of water; in the third 12 drops of sodium thiosulfate. Shake the test tubes.

If we conditionally designate the molar concentration of sodium thiosulfate in test tube 1 as “c”, then accordingly in test tube 2 there will be 2 s mole, in test tube 3 3 s mol.

Add one drop of sulfuric acid into test tube 1, and at the same time turn on the stopwatch: shaking the test tube, watch for the appearance of turbidity in the test tube, holding it at eye level. When the slightest cloudiness appears, stop the stopwatch, note the reaction time and write it down in the table.

Perform similar experiments with the second and third test tubes. Enter the experimental data in the laboratory journal in the form of a table...

Test tube no.	Number of drops Na2S2O3	Number of water drops	Number of drops H2SO4	Concentration of Na 2 S 2 O 3 in moles	Reaction time τ, s	Relative speed V =1/ τ, c -1
					26,09	3,83
					12,19
					8,27	12,09

Graph of reaction rate versus sodium thiosulfate concentration.

Conclusion: with increasing concentration of sodium thiosulfate, the rate of this reaction increases. The dependence graph is a straight line passing through the origin.

Experiment No. 2. Study of the dependence of the rate of a homogeneous reaction on temperature.

Instruments and equipment: test tubes, stopwatch, thermometer, sodium thiosulfate solutions ( III ), sulfuric acid (1M)
Methodology:

Prepare three clean, dry test tubes and number them. Add 10 drops of sodium thiosulfate solution to each of them. Place test tube No. 1 in a glass of water at room temperature and after 1…2 minutes note the temperature. Then add one drop of sulfuric acid to the test tube, simultaneously turn on the stopwatch and stop it when a weak, barely noticeable turbidity appears. Record the time in seconds from the moment the acid is added to the test tube until turbidity appears. Record the result in the table.

Then increase the temperature of the water in the glass by exactly 10 0 either by heating on a hot plate or by mixing with hot water. Place test tube No. 2 in this water, hold for several minutes and add one drop of sulfuric acid, turning on the stopwatch at the same time, shake the test tube with its contents in a glass of water until turbidity appears. If a barely noticeable cloudiness appears, turn off the stopwatch and enter the stopwatch readings into the table. Carry out a similar experiment with the third test tube. First increase the temperature in the glass by another 10 0 , place test tube No. 3 into it, hold for several minutes and add one drop of sulfuric acid, while turning on the stopwatch and shaking the test tube.

Express the results of the experiments in a graph, plotting speed on the ordinate axis, and temperature on the abscissa axis.

Determine the reaction temperature coefficient γ

Test tubes

Temperature

t , 0 C

Reaction time

τ, s

Relative speed

reactions

1/τ,s -1

Temperature coefficient

26,09

17,22

10,74

3,83

5,81

9,31

1,51

1,55

Graph of reaction rate versus temperature.

Conclusion: during the experiment, the average temperature coefficient was calculated, which turned out to be equal to 1.55. Ideally it is

2-4. The deviation from the ideal can be explained by the error in measuring the time of turbidity of the solution. The graph of the reaction rate versus temperature has the form of a parabola branch that does not pass through 0. With increasing temperature, the reaction rate increases

Experiment No. 3 The influence of the concentration of reactants on chemical equilibrium.

Instruments and equipment: test tubes, potassium chloride (crystalline), ferric chloride solutions ( III ), potassium thiocyanate (saturated), distilled water, cylinder
Methodology:

A classic example of a reversible reaction is the interaction between ferric chloride and potassium thiocyanate:

FeCl 3 + 3 KCNS ⇄ Fe(CNS) 3 + 3 KCl.

Red

The resulting iron thiocyanate has a red color, the intensity of which depends on the concentration. By changing the color of the solution, one can judge the shift in chemical equilibrium depending on the increase or decrease in the content of iron thiocyanate in the reaction mixture. Create an equation for the equilibrium constant of this process.

Pour 20 ml of distilled water into a measuring cup or cylinder and add one drop of saturated ferric chloride solution ( III ) and one drop of a saturated solution of potassium thiocyanate. Pour the resulting colored solution equally into four test tubes. Number the test tubes.

Add one drop of saturated ferric chloride solution to the first test tube ( III ). Add one drop of a saturated solution of potassium thiocyanate to the second test tube. Add crystalline potassium chloride to the third test tubeand shake vigorously. Fourth test tube- for comparison.

Based on Le Chatelier's principle, explain what causes the color change in each individual case.

Write the results of the experiment in a table in the form

test tubes

What

added

Change

intensity

coloring

Direction of equilibrium shift

(right left)

In the first and second case, we increased the concentration of the starting substances, so a more intense color is obtained. Moreover, in the second case the color is darker, because the concentration KSCN changes at a cubic rate. In the third experiment, we increased the concentration of the final substance, so the color of the solution became lighter.

Conclusion: with an increase in the concentration of the starting substances, the equilibrium shifts towards the formation of reaction products. As the concentration of products increases, the equilibrium shifts towards the formation of starting substances.

General conclusions: during the experiments, we experimentally established the dependence of the reaction rate on the concentration of the starting substances (the higher the concentration, the higher the reaction rate); the dependence of the reaction rate on temperature (the higher the temperature, the greater the reaction rate); how the concentration of reacting substances affects the chemical equilibrium (with an increase in the concentration of starting substances, the chemical equilibrium shifts towards the formation of products; with an increase in the concentration of products, the chemical equilibrium shifts towards the formation of starting substances)

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2CHEMICAL KINETICS AND CHEMICAL EQUILIBRIUM

2.1 KINETICS OF CHEMICAL REACTIONS

Chemical reactions occur at different rates. Some of them are completed completely in small fractions of a second (explosion), others are carried out in minutes, hours, days and long periods of time. In addition, the same reaction can proceed quickly under some conditions (for example, at elevated temperatures), and slowly under others (for example, upon cooling). Moreover, the difference in the speed of the same reaction can be very large.

When considering the issue of reaction rates, it is necessary to distinguish between homogeneous and heterogeneous reactions. Closely related to these concepts is the concept of phase.

Phase is a part of a system separated from its other parts by an interface, during the transition through which the properties change abruptly.

A homogeneous reaction occurs in the volume of the phase [example - the interaction of hydrogen and oxygen with the formation of water vapor: H 2 (g) + O 2 (g) → H 2 O(g)], and if the reaction is heterogeneous, then it occurs at the phase interface [for example, carbon combustion: C(s) + O2(g) → CO 2 (g)].

The rate of a homogeneous reaction is the amount of substance that reacts or is formed during the reaction per unit time per unit volume of the phase:

Where n- amount of substance, mol; V- phase volume, l;τ - time; WITH- concentration, mol/l.

The rate of a heterogeneous reaction is the amount of substance that reacts or is formed during the reaction per unit time per unit surface area of the phase:

Where S- area of the phase interface.

The most important factors influencing the rate of a homogeneous reaction are the following: the nature of the reactants, their concentrations, temperature, and the presence of catalysts.

Dependence of the reaction rate on the concentrations of the reactants. A reaction between molecules occurs when they collide. Therefore, the rate of a reaction is proportional to the number of collisions that the molecules of the reacting substances undergo. The higher the concentration of each of the starting substances, the greater the number of collisions. For example, reaction rate A + B→ Proportional to the product of concentrations A and B:

v = k · [A] · [B],

Where k- proportionality coefficient, called reaction rate constant. Meaningful value k equal to the reaction rate for the case when the concentrations of the reactants are 1 mol/l.

This ratio expresses law of mass action This law is also called the law existing wt. : At constant temperature, the rate of a chemical reaction is directly proportional to the product of the concentrations of the reactants.

Much less often, a reaction occurs as a result of the simultaneous collision of three reacting particles. For example, reaction

2A+B → A 2 B

can proceed through a triple collision:

A+ A + B→ A 2 B

Then, in accordance with the law of mass action, the concentration of each of the reacting substances is included in the expression of the reaction rate to a degree equal to the coefficient in the reaction equation:

v = k · [A] · [A] · [B] = k · [A] 2 [B]

The sum of exponents in the equation of the law of mass action is called reaction order. For example, in the latter case, the reaction is third order (second - with respect to substance A and first - with respect to substance B.

Dependence of reaction rate on temperature. If we use the results of counting the number of collisions between molecules, the number of collisions will be so large that all reactions must occur instantly. This contradiction can be explained by the fact that only molecules with some energy enter into the reaction.

The excess energy that molecules must have in order for their collision to lead to the formation of a new substance is called activation energy (see Figure 2.1).

Figure 2.1 - Energy diagram for the reaction of formation of product AB from starting substances A and B. If the collision energy of molecules A and B is greater than or equal to activation energies E a , then the energy barrier is overcome, and movement occurs along the reaction coordinate r from starting materials to product. Otherwise, an elastic collision of molecules A and B takes place. The top of the energy barrier corresponds to the transition state (activated complex), in which the AB bond is partially formed.

As temperature increases, the number of active molecules increases Temperature is a measure of the average kinetic energy of molecules, so increasing the temperature leads to an increase in the average speed of their movement.. Therefore, the rate of a chemical reaction should increase with increasing temperature. The increase in reaction rate upon heating is usually characterized as temperature coefficient of reaction rate (γ ) - a number showing how many times the rate of a given reaction increases when the temperature increases by 10 degrees. Mathematically, this dependence is expressed rule van't Hoff :

Where v 1 - speed at temperature t 1 ; v 2 - speed at temperature t 2. For most reactions the temperature coefficientγ lies in the range from 2 to 4.

More strictly, the dependence of the reaction rate (or rather, the rate constant) on temperature is expressed Arrhenius equation :

Where A - pre-exponential a multiplier that depends only on the nature of the reactants; E a - activation energy, which is the height of the energy barrier separating the starting materials and reaction products (see Figure 2.1); R R=8.3144 J/(mol. K). In approximate calculations, R = 8.31 J/(mol. K) is often taken. - universal gas constant; T T - absolute temperature (in Kelvin scale). It is related to temperature in Celsius by the equation
T = t o C + 273.15.
In approximate calculations, the relation is used
T = t o C + 273. -

Chemical kinetics and equilibrium

Goal of the work: study of the influence of temperature on the reaction rate, concentration on the shift in chemical equilibrium.

Theoretical background:

Speed of chemical reaction is the amount of a substance that reacts or is formed as a result of a reaction per unit time per unit volume (for homogeneous reactions) or per unit interface surface (for heterogeneous reactions).

If over a period of time?f = f 2 f 1 the concentration of one of the substances participating in the reaction decreases by?C = C2C1, then the average rate of the chemical reaction for the specified period of time is equal to

The value V expresses the rate of a chemical process over a certain period of time. Therefore, the smaller?f, the closer the average speed will be to the true one.

The rate of a chemical reaction depends on the following factors:

1) the nature and concentration of the reacting substances;

2) temperature of the reaction system;

3) presence of a catalyst;

4) pressure,

5) the size of the phase interface and the mixing rate of the system (for heterogeneous reactions);

6) type of solvent.

A general formulation of the effect of concentration on the rate of a chemical reaction is given by law of mass action(1867, Guldberg, Waage, Beketov).

For the reaction aA + bB = cC V = K[A] a [B] b,

where K is the proportionality coefficient or speed constant;

If [A] = 1 mol/l, [B] = 1 mol/l, then V = K, hence the physical meaning

rate constants K: the rate constant is equal to the reaction rate at concentrations of reactants equal to unity.

The effect of temperature on the reaction rate. As the temperature increases, the frequency of collisions of reacting molecules increases, and therefore the reaction rate increases.

The quantitative effect of temperature on the rate of homogeneous reactions can be expressed by Van't Hoff's rule.

In accordance with Van't Hoff's rule, when the temperature increases (decreases) by 10 degrees, the rate of a chemical reaction increases (decreases) by 2-4 times:

where V (t 2 ) and V (t 1 ) - the rate of chemical reaction at appropriate temperatures; f(t 2 ) And f(t 1 ) - duration of the chemical reaction at appropriate temperatures; G - Van't Hoff temperature coefficient, which can take a numerical value in the range of 2-4.

The dependence of the reaction rate on temperature is expressed by the Arrhenius equation:

where K is the rate constant of the chemical reaction; E a - activation energy;

R - universal gas constant; A - constant; exp is the base of natural logarithms.

The magnitude of the activation energy can be determined if two values of the rate constant K 1 and K 2 are known at temperatures T 1 and T 2, respectively, according to the following formula:

Chemical balance.

All chemical reactions can be divided into two groups: irreversible and reversible. Irreversible reactions proceed to completion - until one of the reactants is completely consumed, i.e. flow in only one direction. Reversible reactions do not proceed to completion. In a reversible reaction, none of the reactants are completely consumed. A reversible reaction can occur in both the forward and reverse directions.

Chemical equilibrium is a state of a system in which the rates of forward and reverse reactions are equal.

For a reversible reaction

m A+ n B? p C+ q D

the chemical equilibrium constant is

Such a transition of a reaction system from one state to another is called a displacement (or shift) of chemical equilibrium.

The direction of the shift in chemical equilibrium is determined by Le Chatelier’s principle: If any external influence is applied to a system that is in a state of chemical equilibrium (change concentration, temperature, pressure), then processes spontaneously arise in this system that tend to weaken the effect produced.

When the temperature increases, the equilibrium shifts towards the endothermic reaction, and when the temperature decreases, it shifts towards the exothermic reaction.

Experiment No. 1 Dependence of the speed of a homogeneous reaction on the concentration of the initial reagents.

b Instruments, equipment: test tubes, stopwatch, solutions of sodium thiosulfate (III), dil. sulfuric acid (1M), water.

b Methodology: This dependence can be studied using the classic example of a homogeneous reaction between sodium thiosulfate and sulfuric acid, proceeding according to the equation

Na 2 S 2 O 3 + H 2 SO 4 = Na 2 SO 4 + Sv + SO 2 ^ + H 2 O.

chemical homogeneous kinetics

For the experiment, you should prepare three dry, clean test tubes and number them. Add 4 drops of sodium thiosulfate solution and 8 drops of water to the first; in the second - 8 drops of sodium thiosulfate and 4 drops of water; in the third - 12 drops of sodium thiosulfate. Shake the test tubes.

If we conditionally designate the molar concentration of sodium thiosulfate in test tube 1 as “c”, then accordingly in test tube 2 there will be 2 s mole, in test tube 3 - 3 s mol.

Perform similar experiments with the second and third test tubes. Enter the experimental data in the laboratory journal in the form of a table.

b Conclusion: with increasing concentration of sodium thiosulfate, the rate of this reaction increases. The dependence graph is a straight line passing through the origin.

Experience No. 2. Study of the dependence of the rate of a homogeneous reaction on temperature.

b Instruments and equipment: test tubes, stopwatch, thermometer, solutions of sodium thiosulfate (III), sulfuric acid (1M)

b Methodology:

Prepare three clean, dry test tubes and number them. Add 10 drops of sodium thiosulfate solution to each of them. Place test tube No. 1 in a glass of water at room temperature and after 1...2 minutes note the temperature. Then add one drop of sulfuric acid to the test tube, simultaneously turn on the stopwatch and stop it when a weak, barely noticeable turbidity appears. Record the time in seconds from the moment the acid is added to the test tube until turbidity appears. Record the result in the table.

Then increase the temperature of the water in the glass by exactly 10 0 either by heating it on a hot plate or by mixing it with hot water. Place test tube No. 2 in this water, hold for several minutes and add one drop of sulfuric acid, turning on the stopwatch at the same time, shake the test tube with its contents in a glass of water until turbidity appears. If a barely noticeable cloudiness appears, turn off the stopwatch and enter the stopwatch readings into the table. Carry out a similar experiment with the third test tube. First increase the temperature in the beaker by another 10 0, place test tube No. 3 in it, hold for several minutes and add one drop of sulfuric acid, while turning on the stopwatch and shaking the test tube.

Express the results of the experiments in a graph, plotting speed on the ordinate axis and temperature on the abscissa axis.

Determine the temperature coefficient of the reaction g

b Conclusion: during the experiment, the average temperature coefficient was calculated, which turned out to be equal to 1.55. Ideally it is

Experiment No. 3 The influence of the concentration of reactants on chemical equilibrium.

b Instruments and equipment: test tubes, potassium chloride (crystal), solutions of iron (III) chloride, potassium thiocyanate (saturated), distilled water, cylinder

b Methodology:

A classic example of a reversible reaction is the interaction between ferric chloride and potassium thiocyanate:

FeCl3+ 3 KCNS D Fe(CNS) 3 + 3 KCl.

Red

Pour 20 ml of distilled water into a measuring cup or cylinder and add one drop of a saturated solution of iron (III) chloride and one drop of a saturated solution of potassium thiocyanate . Pour the resulting colored solution equally into four test tubes. Number the test tubes.

Add one drop of a saturated solution of iron (III) chloride to the first test tube. Add one drop of a saturated solution of potassium thiocyanate to the second test tube. Add crystalline potassium chloride to the third test tube and shake vigorously. The fourth test tube is for comparison.

Based on Le Chatelier's principle, explain what causes the color change in each individual case.

Write the results of the experiment in a table in the form

In the first and second case, we increased the concentration of the starting substances, so a more intense color is obtained. Moreover, in the second case the color is darker, because the concentration of KSCN changes at a cubic rate. In the third experiment, we increased the concentration of the final substance, so the color of the solution became lighter.

General conclusions: during the experiments, we experimentally established the dependence of the reaction rate on the concentration of the starting substances (the higher the concentration, the higher the reaction rate); dependence of the reaction rate on temperature (the higher the temperature, the greater the reaction rate); how the concentration of reacting substances affects the chemical equilibrium (with an increase in the concentration of starting substances, the chemical equilibrium shifts towards the formation of products; with an increase in the concentration of products, the chemical equilibrium shifts towards the formation of starting substances)

Glava 6

Chemical kinetics. Chemical balance.

6.1.Chemicalkinetics.

Chemical kinetics- a branch of chemistry that studies the rates and mechanisms of chemical processes, as well as their dependence on various factors.

The study of the kinetics of chemical reactions allows both to determine the mechanisms of chemical processes and to control chemical processes in their practical implementation.

Any chemical process is the transformation of reagents into reaction products:

reactants→ transition state→ reaction products.

Reagents (starting materials) – substances that enter into the process of chemical interaction.

Reaction products– substances formed at the end of a chemical transformation process. In reversible processes, the products of the direct reaction are reagents of the reverse reaction.

Irreversible reactions– reactions occurring under given conditions in almost the same direction (denoted by the sign →).

For example:

CaCO 3 → CaO + CO 2

Reversible reactions– reactions occurring simultaneously in two opposite directions (indicated by a sign).

Transition state (activated complex) is a state of a chemical system that is intermediate between the starting substances (reagents) and the reaction products. In this state, old chemical bonds are broken and new chemical bonds are formed. The activated complex is then converted into reaction products.

Most chemical reactions are complex and consist of several stages called elementary reactions .

Elementary reaction– a single act of formation or rupture of a chemical bond. The set of elementary reactions that make up a chemical reaction determines mechanism of chemical reaction.

The equation of a chemical reaction usually indicates the initial state of the system (starting substances) and its final state (reaction products). At the same time, the actual mechanism of a chemical reaction can be quite complex and include a number of elementary reactions. Complex chemical reactions include reversible, parallel, sequential And other multi-step reactions (chain reactions , coupled reactions etc.).

If the rates of different stages of a chemical reaction differ significantly, then the rate of a complex reaction as a whole is determined by the rate of its slowest stage. This stage (elementary reaction) is called limiting stage.

Depending on the phase state of the reacting substances, two types of chemical reactions are distinguished: homogeneous And heterogeneous.

Phase is a part of a system that differs in its physical and chemical properties from other parts of the system and is separated from them by an interface. Systems consisting of one phase are called homogeneous systems, from several phases – heterogeneous. An example of a homogeneous system would be air, which is a mixture of substances (nitrogen, oxygen, etc.) in the same gas phase. A suspension of chalk (solid) in water (liquid) is an example of a heterogeneous system consisting of two phases.

Accordingly, reactions in which the interacting substances are in the same phase are called homogeneous reactions. The interaction of substances in such reactions occurs throughout the entire volume of the reaction space.

Heterogeneous reactions include reactions occurring at the interface. An example of a heterogeneous reaction is the reaction of zinc (solid phase) with a solution of hydrochloric acid (liquid phase). In a heterogeneous system, a reaction always occurs at the interface between two phases, since only here can reactants located in different phases collide with each other.

Chemical reactions are usually distinguished by their molecularity, those. by the number of molecules participating in each elementary act of interaction . On this basis, reactions are distinguished between monomolecular, bimolecular and trimolecular.

Monomolecular are called reactions in which the elementary act is a chemical transformation of one molecule , For example:

Bimolecular are considered reactions in which the elementary act occurs when two molecules collide, for example:

IN trimolecular In reactions, an elementary act occurs during the simultaneous collision of three molecules, for example:

The collision of more than three molecules at the same time is almost impossible, so reactions of greater molecularity do not occur in practice.

The rates of chemical reactions can vary significantly. Chemical reactions can proceed extremely slowly, over entire geological periods, such as the weathering of rocks, which is the transformation of aluminosilicates:

K 2 O Al 2 O 3 6SiO 2 + CO 2 + 2H 2 O → K 2 CO 3 + 4SiO 2 + Al 2 O 3 2SiO 2 2H 2 O.

orthoclase – feldspar, potash quartz. sand kaolinite (clay)

Some reactions occur almost instantly, for example, the explosion of black powder, which is a mixture of coal, sulfur and saltpeter:

3C + S + 2KNO3 = N2 + 3CO2 + K2S.

The rate of a chemical reaction serves as a quantitative measure of the intensity of its occurrence.

In general under the speed of a chemical reaction understand the number of elementary reaction acts occurring per unit of time in a unit of reaction space.

Since for homogeneous processes the reaction space is the volume of the reaction vessel, then

for homogeneous reactions With The speed of a chemical reaction is determined by the amount of substance reacted per unit time in a unit volume.

Considering that the amount of a substance contained in a certain volume characterizes the concentration of the substance, then

the reaction rate is a value indicating the change in the molar concentration of one of the substances per unit time.

If, at constant volume and temperature, the concentration of one of the reactants decreases from With 1 to With 2 for the period from t 1 to t 2, then, in accordance with the definition, the reaction rate for a given period of time (average reaction rate) is equal to:

Typically for homogeneous reactions the rate dimension V[mol/l·s].

Since for heterogeneous reactions the reaction space is surface , on which the reaction occurs, then for heterogeneous chemical reactions, the reaction rate refers to the unit surface area on which the reaction occurs. Accordingly, the average rate of a heterogeneous reaction has the form:

Where S– surface area on which the reaction occurs.

The speed dimension for heterogeneous reactions is [mol/l·s·m2].

The speed of a chemical reaction depends on a number of factors:

the nature of the reacting substances;

concentrations of reactants;

pressure (for gas systems);

system temperature;

surface area (for heterogeneous systems);

the presence of a catalyst and other factors in the system.

Since every chemical interaction is the result of a collision of particles, an increase in concentration (the number of particles in a given volume) leads to more frequent collisions and, as a consequence, an increase in the reaction rate. The dependence of the rate of chemical reactions on the molar concentrations of the reactants is described by the basic law of chemical kinetics - law of mass action , which was formulated in 1865 by N.N. Beketov and in 1867 by K.M. Guldberg and P. Waage.

Law of mass action reads: the rate of an elementary chemical reaction at a constant temperature is directly proportional to the product of the molar concentrations of the reactants in powers equal to their stoichiometric coefficients.

The equation expressing the dependence of the reaction rate on the concentration of each substance is called kinetic equation of the reaction .

It should be noted that the law of mass action is fully applicable only to the simplest homogeneous reactions. If a reaction occurs in several stages, then the law is valid for each stage, and the speed of a complex chemical process is determined by the speed of the slowest reaction, which is limiting stage the whole process.

In general, if an elementary reaction occurs simultaneously T molecules of matter A And n molecules of matter IN:

mA + nIN = WITH,

then the equation for the reaction rate is (kinetic equation) has the form:

Where k– proportionality coefficient, which is called rate constant chemical reaction; [ A A; [B] – molar concentration of the substance B; m And n– stoichiometric coefficients in the reaction equation.

To understand physical meaning of the reaction rate constant , it is necessary to take in the equations written above the concentrations of the reacting substances [ A] = 1 mol/l and [ IN] = 1 mol/l (or equate their product to unity), and then:

From here it is clear that reaction rate constant k is numerically equal to the reaction rate in which the concentrations of reactants (or their product in kinetic equations) are equal to unity.

Reaction rate constant k depends on the nature of the reactants and temperature, but does not depend on the concentration of the reagents.

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

For example, in the methane synthesis reaction:

If a reaction occurs in the gas phase, then its rate is significantly affected by a change in pressure in the system, since a change in pressure in the gas phase leads to a proportional change in concentration. Thus, an increase in pressure leads to a proportional increase in concentration, and a decrease in pressure, accordingly, reduces the concentration of the gaseous reactant.

Changes in pressure have virtually no effect on the concentration of liquid and solid substances (condensed state of matter) and have no effect on the rate of reactions occurring in the liquid or solid phases.

Chemical reactions are carried out due to the collision of particles of reacting substances. However, not every collision of reactant particles is effective , i.e. leads to the formation of reaction products. Only particles with increased energy - active particles , are capable of carrying out a chemical reaction. With increasing temperature, the kinetic energy of particles increases and the number of active ones increases, therefore, the rate of chemical processes increases.

The dependence of the reaction rate on temperature is determined van't Hoff's rule : for every 10 0 C increase in temperature, the rate of a chemical reaction increases two to four times.

V 1 – reaction rate at the initial temperature of the system t 1 , V 2 – reaction rate at the final temperature of the system t 2 ,

γ – temperature coefficient of reaction (van’t Hoff coefficient), equal to 2÷4.

Knowing the value of the temperature coefficient γ makes it possible to calculate the change in reaction rate with increasing temperature from T 1 to T 2. In this case, you can use the formula:

Obviously, as the temperature increases in arithmetic progression, the reaction rate increases exponentially. The greater the value of the reaction temperature coefficient g, the greater the effect of temperature on the reaction rate.

It should be noted that Van't Hoff's rule is approximate and is applicable only for an approximate assessment of the effect of small changes in temperature on the reaction rate.

The energy required for reactions to occur can be provided by various influences (heat, light, electric current, laser radiation, plasma, radioactive radiation, high pressure, etc.).

Reactions can be divided into thermal, photochemical, electrochemical, radiation-chemical etc. With all these influences, the proportion of active molecules that have energy equal to or greater the minimum energy required for a given interaction E min.

When active molecules collide, a so-called activated complex , within which the redistribution of atoms occurs.

The energy required to increase the average energy of the molecules of the reacting substances to the energy of the activated complex is called the activation energy Ea.

Activation energy can be considered as a certain additional energy that reagent molecules must acquire in order to overcome a certain energy barrier . Thus, E a ra in the difference between the average energy of the reacting particles E ref and energy of the activated complex E min. The activation energy is determined by the nature of the reagents. Meaning E a ranges from 0 to 400 kJ. If the value E a exceeds 150 kJ, then such reactions practically do not occur at temperatures close to the standard one.

The change in energy of a system during a reaction can be graphically represented using the following energy diagram (Figure 6.1).

Reaction Path

Rice. 6.1. Energy diagram of an exothermic reaction:

E out is the average energy of the starting substances; Econd – average energy of reaction products; E min – energy of the activated complex; E act – activation energy; ΔH р – thermal effect of a chemical reaction

From the energy diagram it is clear that the difference between the energy values of the reaction products and the energy of the starting substances will represent the thermal effect of the reaction.

E cont. – E ref. = ΔН р.

According to Arrhenius equation, the higher the activation energy value E act, the greater the rate constant of the chemical reaction k depends on temperature:

E- activation energy (J/mol),

R - universal gas constant,

T– temperature in K,

A- Arrhenius constant,

e= 2.718 – the base of natural logarithms.

Catalysts- These are substances that increase the rate of a chemical reaction. They react with reagents to form a chemical intermediate and are released at the end of the reaction. The effect that catalysts have on chemical reactions is called catalysis.

For example, a mixture of aluminum powder and crystalline iodine at room temperature shows no noticeable signs of interaction, but a drop of water is enough to cause a violent reaction:

Distinguish homogeneous catalysis (the catalyst forms a homogeneous system with the reacting substances, for example, a gas mixture) and heterogeneous catalysis (the catalyst and reactants are in different phases and the catalytic process occurs at the phase interface).

To explain the mechanism of homogeneous catalysis, the most widely used intermediate theory (proposed by the French researcher Sabatier and developed in the works of the Russian scientist N.D. Zelinsky). According to this theory, a slow process, for example, the reaction:

in the presence of a catalyst occurs quickly, but in two stages. In the first stage of the process, an intermediate compound of one of the reagents with the catalyst is formed A...kat.

First stage:

A + kat = A.∙. kat.

At the second stage, the resulting compound forms an activated complex with another reagent [ A.∙.kat.∙.B], which turns into the final product AB with catalyst regeneration kat.

Second stage:

A.∙.kat + B = = AB + kat.

The intermediate interaction of the catalyst with the reagents directs the process onto a new path, characterized by a lower energy barrier. Thus, The mechanism of action of catalysts is associated with a decrease in the activation energy of the reaction due to the formation of intermediate compounds.

An example would be a slow reaction:

2SO2 + O2 = 2SO3 slowly.

In the industrial nitrous method for producing sulfuric acid, nitrogen oxide (II) is used as a catalyst, which significantly speeds up the reaction:

Heterogeneous catalysis is widely used in oil refining processes. Catalysts include platinum, nickel, aluminum oxide, etc. Hydrogenation of vegetable oil occurs on a nickel catalyst (nickel on kieselguhr), etc.

An example of heterogeneous catalysis is the oxidation of SO 2 to SO 3 on a V 2 O 5 catalyst in the production of sulfuric acid by the contact method.

Substances that increase the activity of the catalyst are called promoters (or activators). At the same time, promoters themselves may not have catalytic properties.

Catalytic poisons - foreign impurities in the reaction mixture leading to partial or complete loss of catalyst activity. Thus, traces of phosphorus and arsenic cause a rapid loss of activity by the V 2 O 5 catalyst in the oxidation reaction of SO 2 to SO 3.

Many important chemical productions, such as the production of sulfuric acid, ammonia, nitric acid, synthetic rubber, a number of polymers, etc., are carried out in the presence of catalysts.

Biochemical reactions in plant and animal organisms accelerate biochemical catalysts – enzymes.

Sharp It is possible to slow down the occurrence of undesirable chemical processes by adding special substances to the reaction medium - inhibitors. For example, to inhibit undesirable processes of corrosion destruction of metals, various metal corrosion inhibitors .

6.1.1. Questions for self-control of theory knowledge

on the topic "Chemical kinetics"

1. What does chemical kinetics study?

2. What is commonly understood by the term “reagents”?

3. What is commonly understood by the term “reaction products”?

4. How are reversible processes designated in chemical reactions?

5. What is commonly understood by the term “activated complex”?

6. What is an elementary reaction?

7. What reactions are considered complex?

8. Which stage of reactions is called the limiting stage?

9. Define the concept of “phase”?

10. What systems are considered homogeneous?

11. What systems are considered heterogeneous?

12. Give examples of homogeneous systems.

13. Give examples of heterogeneous systems.

14. What is considered the “molecularity” of a reaction?

15. What is meant by the term “rate of a chemical reaction”?

16. Give examples of fast and slow reactions.

17. What is meant by the term “rate of a homogeneous chemical reaction”?

18. What is meant by the term “rate of a heterogeneous chemical reaction”?

19. On what factors does the rate of a chemical reaction depend?

20. Formulate the basic law of chemical kinetics.

21. What is the rate constant of chemical reactions?

22.What factors does the rate constant of chemical reactions depend on?

23. The concentrations of which substances are not included in the kinetic equation of chemical reactions?

24. How does the rate of a chemical reaction depend on pressure?

25. How does the rate of a chemical reaction depend on temperature?

26. How is the “Van’t Hoff Rule” formulated?

27. What is the “temperature coefficient of a chemical reaction”?

28. Define the concept of “activation energy”.

29. Define the concept of “catalyst for a chemical reaction”?

30. What is homogeneous catalysis?

31. What is heterogeneous catalysis?

32. How is the mechanism of action of a catalyst in homogeneous catalysis explained?

33. Give examples of catalytic reactions.

34. What are enzymes?

35. What are promoters?

6.1.2. Examples of solving typical problems

on the topic "Chemical kinetics"

Example 1. The reaction rate depends on the contact surface area of the reactants:

1) sulfuric acid with barium chloride solution,

2) combustion of hydrogen in chlorine,

3) sulfuric acid with potassium hydroxide solution,

4) combustion of iron in oxygen.

The rate of heterogeneous reactions depends on the contact surface area of the reacting substances. Among the above reactions, a heterogeneous reaction, i.e. characterized by the presence of different phases is the combustion reaction of iron (solid phase) in oxygen (gas phase).

Answer. 3.

Example 2. How will the reaction rate change?

2H 2(g) + O 2(G) = 2H 2 O (g)

when the concentration of the starting substances doubles?

Let us write the kinetic equation of the reaction, which establishes the dependence of the reaction rate on the concentration of the reactants:

V 1 = k [N 2 ] 2 · [O 2 ].

If the concentrations of the starting substances are doubled, the kinetic equation will take the form:

V 2 = k (2 [N 2 ]) 2 2 [O 2 ] = 8 k [N 2 ] 2 · [O 2 ], i.e.

When the concentration of the starting substances was doubled, the rate of this reaction increased 8 times.

Answer. 8.

Example 3. How will the reaction rate change if the total pressure in the system CH 4 (G) + 2O 2 (G) = CO 2 (G) + 2H 2 O (G) is reduced by 5 times?

In accordance with the kinetic equation of the reaction, the rate of this reaction will be determined:

V 1 = k[CH 4 ] · [O 2 ] 2 .

When the pressure decreases by a factor of five, the concentration of each gaseous substance will also decrease by a factor of five. The kinetic equation of the reaction under these conditions will be as follows:

it can be determined that the reaction rate has decreased by 125 times.

Answer. 125.

Example 4. How will the rate of a reaction characterized by a reaction temperature coefficient of 3 change if the temperature in the system increases from 20 to 60°C?

Solution. According to van't Hoff's rule

When the temperature increased by 40 0 C, the rate of this reaction increased 81 times

Answer. 81.

6.1.3. Questions and exercises for self-study

Rate of chemical reactions

1. Depending on the physical state of the reacting substances, chemical reactions are divided into:

1) exothermic and endothermic,

2) reversible and irreversible,

3) catalytic and non-catalytic,

4) homogeneous and heterogeneous.

2. Indicate the number or sum of conventional numbers under which homogeneous reactions are given:

3. Indicate the number or sum of conventional numbers under which expressions are given that can be used to calculate the rate of a homogeneous reaction:

4. The unit of measurement for the rate of a homogeneous reaction can be:

1) mol/l s,

3) mol/l·,

4) l/mol·s.

5. Indicate the number or sum of conditional numbers under which fair expressions are given. During a homogeneous reaction

A + 2B® 2 C + D:

1) concentration A And IN are decreasing

2) concentration WITH increases faster than concentration D,

4) concentration IN decreases faster than concentration A,

8) the reaction rate remains constant.

6. What number is shown on the line that correctly reflects the change over time in the concentration of the substance formed in the reaction:

7. Change over time in the concentration of the starting substance in a reaction that proceeds to completion, right describes the curve:

9. Indicate the number or sum of conventional numbers under which reactions are given, the speed of which does not depend on what substance it is calculated from?

10. Indicate the number or sum of conditional numbers under which the factors influencing the reaction rate are given:

1) the nature of the reacting substances,

2) concentration of reacting substances,

4) temperature of the reaction system,

8) the presence of a catalyst in the reaction system.

11. The basic law of chemical kinetics establishes the dependence of the reaction rate on:

1) temperatures of reacting substances,

2) concentrations of reacting substances,

3) the nature of the reacting substances,

4) reaction time.

12. Indicate the number or sum of conditional numbers under which the correct statements are given. Chemical kinetics:

1) section of physics,

2) studies the rate of a chemical reaction,

4) uses the law of mass action,

8) studies the dependence of the rate of reactions on the conditions of their occurrence.

13. Ya.Kh. Van't Hoff:

1) the first Nobel Prize laureate in chemistry,

2) studied the dependence of the reaction rate on temperature,

4) studied the dependence of the reaction rate on the concentration of substances,

8) formulated the law of mass action.

14. Under the same conditions, the reaction proceeds faster:

1) Ca + H 2 O ®

3) Mg + H 2 O ®

4) Zn + H 2 O ®

15. The rate of hydrogen evolution is highest in the reaction:

1) Zn + HCl (5% solution) ®

2) Zn + HCl (10% solution) ®

3) Zn + HCl (15% solution) ®

4) Zn + HCl (30% solution) ®

16. Concentration of reactant does not affect on the reaction rate if this substance is taken into the reaction in:

1) solid state,

2) gaseous state,

3) dissolved state.

17. Calculate the average rate of reaction A + B = C (mol/l×s), if it is known that the initial concentration of A was 0.8 mol/l, and after 10 seconds it became 0.6 mol/l.

1) 0,2, 2) 0,01, 3) 0,1, 4) 0,02.

18. By how much mol/l did the concentrations of substances A and B decrease in the reaction? A + 2B® 3 C, if it is known that during the same time the concentration WITH increased by 4.5 mol/l?

D WITH A D WITH B

19. Calculate the average rate of the reaction 2CO + O 2 ® 2CO 2 (mol/l×s), if it is known that the initial concentration of CO was 0.60 mol/l, and after 10 seconds it became 0.15 mol/l. By how many mol/l did the CO 2 concentration change over this period of time?

3) 0,045; 0,045,

20. How many degrees does the system need to be heated so that the rate of the reaction occurring in it increases by 2–4 times?

1) 150, 2) 10, 3) 200, 4) 50.

21. The reaction rate at 20°C is 0.2 mol/l×s. Determine the reaction rate at 60°C (mol/l×s) if the temperature coefficient of the reaction rate is 3.

1) 16,2, 2) 32,4, 3) 8,1, 4) 4,05.

22. Empirical dependence of reaction rate on temperature right reflects the equation:

23. The reaction rate at 20°C is 0.08 mol/l×s. Calculate the reaction rate at 0°C (mol/l×s), if the temperature coefficient of the reaction rate is 2.

1) 0,16, 2) 0,04, 3) 0,02, 4) 0,002.

24. How many times will the reaction rate increase when the temperature increases by 40°C, if the temperature coefficient of the reaction rate is 3?

1) 64, 2) 243, 3) 81, 4) 27.

25. By how many degrees should the temperature be increased to increase the reaction rate by 64 times, if the temperature coefficient of the reaction rate is 4?

1) 60, 2) 81, 3) 27, 4) 30.

26. Calculate the temperature coefficient of the reaction rate if it is known that when the temperature increases by 50 o C, the reaction rate increases by 32 times.

1) 3, 2) 2, 3) 4, 4) 2,5.

27. The reason for the increase in reaction rate with increasing temperature is an increase in:

1) the speed of movement of molecules,

2) the number of collisions between molecules,

3) the proportion of active molecules,

4) stability of the molecules of the reaction products.

28. Indicate the number or sum of conventional numbers under which the reactions for which MnO 2 is a catalyst are given:

1) 2KClO 3 ® 2KCl + 3O 2,

2) 2Al + 3I 2 ® 2AlI 3,

4) 2H 2 O 2 ® 2H 2 O + O 2,

8) 2SO 2 + O 2 ® 2SO 3 .

29. Indicate the number or sum of conventional numbers under which the correct answers are given. Using catalytic reactions in industry, the following is obtained:

1) hydrochloric acid,

2) sulfuric acid,

4) ammonia,

8) nitric acid.

30. Indicate the number or sum of conventional numbers under which the correct answers are given. Catalyst:

1) participates in the reaction,

2) used only in solid state,

4) is not consumed during the reaction,

8) necessarily contains a metal atom in its composition.

31. Indicate the number or sum of conventional numbers under which the correct answers are given. The following are used as catalysts:

32. Substances that reduce the activity of a catalyst are called:

1) promoters,

2) regenerators,

3) inhibitors,

4) catalytic poisons.

33. Catalytic is not reaction:

1) (C 6 H 10 O 5) n + n H2O® n C6H12O6,

cellulose

2) 2SO 2 + O 2 ® 2SO 3,

3) 3H 2 + N 2 ® 2NH 3 ,

4) NH 3 + HCl ® NH 4 Cl.

34. Under what number is the equation of homogeneous catalysis given:

35. The mechanism of action of the catalyst correctly reflects the statement. Catalyst:

1) increasing the kinetic energy of the initial particles, increases the number of their collisions,

2) forms intermediate compounds with starting substances that are easily converted into final substances,

3) without interacting with the starting substances, it directs the reaction along a new path,

4) decreasing the kinetic energy of the initial particles, increases the number of their collisions.

36. The role of a promoter in a catalytic reaction is that it:

1) reduces the activity of the catalyst,

2) increases the activity of the catalyst,

3) leads the reaction in the desired direction,

4) protects the catalyst from catalytic poisons.

37. Enzymes:

1) biological catalysts,

2) have a protein nature,

4) do not differ in the specificity of action,

8) accelerate biochemical processes in living organisms.

38. The reaction is heterogeneous:

39. Indicate the number or sum of conventional numbers under which the correct answers are given. To increase the burning rate of coal: C + O 2 ® CO 2, you need to:

1) increase the concentration of O 2,

2) increase the concentration of coal,

4) grind the coal,

8) increase the concentration of carbon dioxide.

40. If reactant A is taken into the reaction: A t + X gas ® in the solid state, then the reaction rate is affected by:

1) concentration A,

2) surface area of contact between A and X,

4) molar mass A,

8) concentration of substance X.

41. The dimension of the rate of a heterogeneous reaction is:

1) mol/l, 2) mol/cm 3 ×s,

3) mol/l×s 4) mol/cm 2 ×s.

42. Indicate the number or sum of conventional numbers under which the correct answers are given. The fluidized bed principle is used:

1) to increase the contact surface of the reagents,

2) when firing pyrites,

4) during catalytic cracking of petroleum products,

8) to regenerate catalyst activity.

43. Least

1) Na + H 2 O ® 2) Ca + H 2 O ®

3) K + H 2 O ® 4) Mg + H 2 O ®

44. The graph shows energy diagrams of the non-catalytic and catalytic reaction of hydrogen iodide decomposition. The change in activation energy reflects the energy segment:

1) b, 2) c, 3) d, 4) b–c.

45. The greatest The reaction described by the scheme has activation energy:

1) AgNO 3 + KCl ® AgCl + KNO 3,

2) BaCl 2 + K 2 SO 4 ® BaSO 4 + 2KCl,

3) 2Na + 2H 2 O ® 2NaOH + 2H 2,

6.2. Chemical balance.

Along with practically irreversible chemical reactions:

СaCl 2 + 2AgNO 3 = Ca(NO 3) 2 + 2AgCl↓, etc.

Numerous processes are known when the chemical transformation does not reach completion, but an equilibrium mixture of all participants and products of the reaction occurs, located both on the left and on the right sides of the stoichiometric reaction equation. Thus, under standard conditions the system is reversible:

Let us consider the features of the occurrence of reversible processes using the example of a system, which, in general, has the form:

Provided that the forward → and reverse ← reactions occur in one stage, according to the law of mass action, the speed values for the forward reaction ( V direct) and reverse ( V the reactions are described by the following kinetic equations:

Where k straight And k arr - rate constants, respectively, of forward and reverse reactions.

At the initial moment of time (see Fig. 6.2), the concentrations of the starting substances [A] and [B], and therefore the rate of the direct reaction, have a maximum value. The concentrations of the reaction products [C] and [D] and the rate of the reverse reaction at the initial moment are zero. During the reaction, the concentrations of the starting substances decrease, which leads to a decrease in the rate of the forward reaction. The concentrations of the reaction products, and, consequently, the rate of the reverse reaction increase. Finally, there comes a point at which the rates of the forward and reverse reactions become equal.

The state of the system in which V straight = V arr. called chemical equilibrium. This balance is dynamic , since a two-way reaction takes place in the system - in direct ( A And B– reagents, C And D– products) and in reverse ( A And B– products, C and D– reagents) directions.

V arr.

Reaction time

Rice. 6.2. Dependence of the rates of forward and reverse reactions

from the time of their occurrence.

In a reversible system in equilibrium, the concentrations of all participants in the process are called equilibrium concentrations, since in this case both forward and reverse reactions occur constantly and at the same speed.

A quantitative characteristic of chemical equilibrium can be derived using the appropriate kinetic equations :

Since the reaction rate constants at a fixed temperature are constant, the ratio will also be constant

called chemical equilibrium constant. Equating the right-hand sides of the kinetic equations for the forward and reverse reactions, we can obtain:

Where K r– chemical equilibrium constant, expressed in terms of equilibrium concentrations of reaction participants.

The chemical equilibrium constant is the ratio of the product of the equilibrium concentrations of reaction products to the product of the equilibrium concentrations of the starting substances in powers of their stoichiometric coefficients.

For example, for a reversible reaction

The expression for the equilibrium constant has the form:

If two or more phases are involved in the process of a chemical transformation, then the expression for the equilibrium constant should take into account only those of them in which changes in the concentrations of the reactants occur. For example, in the expression for the equilibrium constant for the system

the total number of moles of gaseous substances before and after the reaction remains constant and the pressure in the system does not change. The equilibrium in this system does not shift when pressure changes.

The influence of temperature changes on the shift of chemical equilibrium.

In each reversible reaction, one of the directions corresponds to an exothermic process, and the other to an endothermic process. So in the reaction of ammonia synthesis, the forward reaction is exothermic, and the reverse reaction is endothermic.

1) the concentrations of H 2, N 2 and NH 3 do not change over time,

3) the number of NH 3 molecules decaying per unit time is equal to half the total number of H 2 and N 2 molecules formed during this time,

4) the total number of H 2 and N 2 molecules converted into NH 3 per unit time is equal to the number of NH 3 molecules formed during the same time.

49. Indicate the number or sum of conventional numbers under which the correct answers are given. Chemical equilibrium in the system: 2SO 2 + O 2 2SO 3 ∆Н ˂0 will be disrupted by:

1) reducing the pressure in the system,

2) heating,

4) increase in oxygen concentration.

50. Indicate the number or sum of conventional numbers under which the correct answers are given. To shift the equilibrium in the system N 2 + 3H 2 2NH 3 ∆H ˂0 to the left, you need to:

1) introduce H 2 into the system,

2) remove NH 3 from the system,

4) increase blood pressure,

8) increase the temperature.

51. To shift the equilibrium of the reaction 2SO 2 + O 2 2SO 3 ∆Н ˂0 to the right, it is necessary:

1) heat the system,

2) introduce O 2 into the system,

4) introduce SO 3 into the system,

8) reduce the pressure in the system.

52. Le Chatelier's rule (principle) does not match statement:

1) an increase in temperature shifts the equilibrium towards an endothermic reaction;

2) a decrease in temperature shifts the equilibrium towards an exothermic reaction;

3) an increase in pressure shifts the equilibrium towards a reaction leading to an increase in volume;

N 2 + O 2 ∆H ˂0.2H 2 O (steam), 2NH 3 cat. 3H 2 + N 2 . B,

2) k 1 H = k 2 2 ,

67. For the equilibrium constant ( Kp) affects:

1) pressure,

2) temperature,

3) concentration,

4) catalyst.