What can be cooked from squid: quick and tasty

38. Influence of the catalyst on the rate of chemical reactions. Reasons for the influence of the catalyst.

Substances that are not consumed as a result of the reaction, but affect its rate, are called catalysts. Catalysts that reduce the rate of reaction are called inhibitors. The effect of catalysts on chemical reactions is called catalysis ... The essence of catalysis lies in the fact that in the presence of a catalyst, the path along which the overall reaction takes place changes, other transition states with different activation energies are formed, and therefore the rate of the chemical reaction also changes. Distinguish between homogeneous and heterogeneous catalysis. In heterogeneous catalysis, the reaction proceeds on the catalyst surface. Hence it follows that the activity of the catalyst depends on the size and properties of its surface. In order to have a large surface, the catalyst must have a porous structure or be in a highly fragmented state. Catalysts are distinguished by their selectivity: they act selectively on the processes, directing it in a certain direction. Negative catalysis is used to inhibit corrosion.

39. Reversible processes. Chemical equilibrium. Equilibrium constant.

Reactions that proceed in only one direction and end with the complete transformation of the initial reacting substances into final substances are called irreversible... 2KClO 3 = 2KCl + 3O 2 . Reversible such reactions are called that simultaneously proceed in two mutually opposite directions. 3H 2 + N 2 ⇆ 2NH 3

Reversible reactions do not proceed completely: none of the reactants is completely consumed. Reversible processes: at first, when mixing the initial substances, the rate of the direct reaction is high, and the rate of the reverse is equal to zero. As the reaction proceeds, the starting materials are consumed and their concentrations fall, as a result of which the reaction rate decreases. At the same time, reaction products appear, whose concentration increases, and, accordingly, the rate of the reverse reaction increases. When the rates of the forward and reverse reactions become the same, chemical equilibrium occurs. It is called dynamic equilibrium, since direct and reverse reactions proceed, but due to the same rates, changes in the system are not noticeable. The quantitative characteristic of chemical equilibrium is a quantity called the constant of chemical equilibrium. In equilibrium, the rates of the forward and reverse reactions are equal, while constant concentrations of the initial substances and reaction products, called equilibrium concentrations, are established in the system. For 2CO + O 2 = 2CO 2, the equilibrium constant can be calculated by the equation: The numerical value of the equilibrium constant in the first approximation characterizes the yield of this reaction. The reaction yield is the ratio of the amount of the obtained substance to the amount that would be obtained if the reaction proceeded to the end. K >> 1 the reaction yield is large, K<10-6). В случае гетерогенных реакций в выражение константы равновесия входят концентрации только тех веществ, которые находятся в наиболее подвижной фазе. Катализатор не влияет на константу равновесия. Он может только ускорить наступление равновесия. K=e^(-ΔG/RT).

40. The influence of various factors on the displacement of the balance. Le Chatelier principle.

If the system is in equilibrium, then it will remain in it as long as the external conditions remain constant. The process of changing any conditions that affect equilibrium is called equilibrium displacement.

Le Principle: If the system. find. in equilibrium to exert external influence, then the system of betrayals. in a way that compensates for this impact.

Consequences: 1) With an increase in temperature. balance shifted. in favor of an endothermic reaction.

2) With increasing pressure, the equilibrium is displaced. towards a smaller volume (or less.number of moles)

3) With an increase in the concentration of one of the initial substances, the equilibrium shifts towards an increase in the concentration of the reaction products, and vice versa.

In the 9th-10th grades of secondary school, they continue to form concepts about the rate of a chemical reaction, about the influence of various factors on the rate of chemical transformations, expand and deepen knowledge about catalysis and catalysts, and give some ideas about the mechanism of catalytic phenomena.

In the topic "Alkali metals", demonstrating such experiments as the interaction of sodium with water and hydrochloric acid, the interaction of potassium and sodium with water, the teacher emphasizes that some of these reactions proceed faster under the same conditions than others. For example, sodium reacts more vigorously with hydrochloric acid than with water; potassium reacts more vigorously with water than sodium. After experiments on combustion in chlorine of sodium, copper, antimony, hydrogen, organic substances, one can ask questions: "Why was antimony powder taken for combustion in chlorine, and not pieces? Why a bundle of thin copper wire burns in chlorine, but a thick wire does not burn?" In these cases, the difference in the interaction of substances is explained either by the nature of the substances themselves and the structure of atoms, or by a different contact surface.

In the same topic, when introducing students to the properties of hydrochloric acid, it is useful to find out why the reactions between this acid and metals (zinc, magnesium) accelerate over time. Acceleration depends, in particular, on the fact that during these reactions a large amount of heat is released, and with the heating of substances, the speed of interaction increases.

Using the example of the reaction of interaction of aluminum with iodine, one should recall what a catalyst is and show that water can be a catalyst. A mixture of iodine and aluminum powders is poured onto the asbestos mesh in a slide and a few drops of water are poured. The interaction of substances under the influence of water is accelerated, a flame breaks out. The teacher draws attention to the fact that in the mixture that was not poured out of the porcelain cup onto the grid, the flash did not occur, but it can occur after a while and without water.

It should be noted that water not only accelerates the interaction of aluminum with iodine, but also plays a catalytic role in many chemical processes. The catalytic action of water in the combustion of various gases used in technology is very important.

When considering the properties of hydrogen peroxide, it is indicated that hydrogen peroxide is a very fragile substance. When stored in glassware, it slowly decomposes with the release of heat:

2H 2 O 2 = 2H 2 O 4 + O 2 + 46 kcal

The teacher asks the students to list the conditions that accelerate the decomposition of hydrogen peroxide. They can

indicate in this case: 1) heating, 2) the effect of catalysts, 3) an increase in the concentration of the solution. It can be added that the decomposition of hydrogen peroxide is also faster in the light, this can be confirmed by experience in extracurricular activities. Pour hydrogen peroxide into two flasks fixed in tripods, close them with stoppers with gas pipes. Place the tubes under overturned cylinders or test tubes filled with water and lowered into a wide vessel with water. Wrap one of the flasks with black paper. Put the devices on a window illuminated by the sun or illuminate them with an electric lamp of 75-100 V. The experiment will show the rapid decomposition of hydrogen peroxide under the influence of light.

Then the students in the lesson independently study the change in the rate of decomposition of hydrogen peroxide under the action of catalysts. For work, they give out a 3-5% solution of hydrogen peroxide, manganese dioxide, concentrated hydrochloric acid, a splinter, a funnel, filter paper, several test tubes.

Tasks: 1) Check if there is decomposition of hydrogen peroxide in the solution that is issued? 2) Using manganese dioxide, accelerate the decomposition reaction of hydrogen peroxide. 3) Prove that manganese dioxide has not chemically changed as a result of the reaction * 4) Prove that manganese dioxide, already used as a catalyst, can again accelerate the decomposition of hydrogen peroxide.

* (Sample with hydrochloric acid when heated.)

After completing independent work, the teacher shows that various catalysts can be used to accelerate the same chemical reaction, that the decomposition of an inorganic substance (hydrogen peroxide) is accelerated by organic catalysts - enzymes. A 3% hydrogen peroxide solution is poured into a small beaker, then a small piece of raw meat is placed in it. Oxygen is intensively released from the solution, since the enzyme catalase is contained in the blood and tissues of animals. It should be emphasized that enzymes are excellent natural accelerators of reactions. One of the important tasks of chemistry of the future is the artificial production and industrial application of catalysts that will resemble enzymes in their composition and catalytic properties.

An experiment was carried out to explain why the decomposition of hydrogen peroxide proceeds faster when stored in glassware. A solution of hydrogen peroxide is poured into three test tubes, a solution of sulfuric acid is added to one of them, caustic soda is added to the other, and the third is left for comparison (control solution). All three solutions are heated (not to boiling). Oxygen will be released strongly from the test tube with hydrogen peroxide and sodium hydroxide solutions, less strongly - from the test tube with control solution. In the presence of sulfuric acid (hydrogen ions), hydrogen peroxide does not decompose. OH ions catalyze the decomposition of hydrogen peroxide, therefore, in glassware, the walls of which release hydroxyl ions into the solution, hydrogen peroxide easily decomposes.

Consolidation and development of knowledge about the rate of a chemical reaction continues further. Passing a mixture of sulfur dioxide and oxygen through a heated glass tube without a catalyst, the teacher shows that the formation of sulfuric anhydride under these conditions is not noticeable, and asks the students how the interaction of gases can be accelerated. In the conversation, it turns out that such methods of accelerating reactions as increasing the concentration of reagents, raising the temperature, without the use of a catalyst, do not give the necessary results. The oxidation reaction of sulfur dioxide to sulfuric acid is reversible:

2SO 2 + O 2 ↔ 2SO 3 + Q,

and an increase in temperature accelerates decomposition of sulfuric anhydride to a greater extent than its formation.

It is checked whether iron oxide will be a catalyst for the oxidation reaction of sulfur dioxide. When demonstrating contact oxidation of sulfur dioxide to sulfuric anhydride in the presence of iron oxide, the production of sulfuric anhydride, fuming in air, is observed. It is then found that the reaction does not chemically change the iron oxide. To do this, repeat the experience of contact oxidation of sulfur dioxide to sulfuric anhydride with the same portion of iron oxide. It is further noted that various catalysts can be used to accelerate the oxidation of sulfur dioxide. In addition to iron oxide, platinum was used in the chemical industry, and now vanadium pentoxide V 2 O 5 * is used.

* (The vanadium catalyst currently used has a complex composition (see: D. A. Epshtein. Chemistry teacher about chemical technology, Moscow, Publishing House of the Academy of Pedagogical Sciences of the RSFSR, 1961).)

It is also important to emphasize the property of the catalyst, accelerating the reaction, not to influence its reversibility: the oxidation reaction of sulfur dioxide to sulfuric acid, and in the case of using a catalyst, remains reversible.

When studying the contact method for the production of sulfuric acid, it is necessary to consider the use of the catalyst in industry. Without a catalyst, the rapid production of large amounts of sulfuric anhydride would not be possible, but its use raises some additional requirements for the process conditions. The fact is that impurities to the reactants have a negative effect on the catalyst. Arsenic trioxide, as they say, "poisons" it negatively on the vanadium catalyst. Therefore, it is necessary to thoroughly clean the reacting gases from impurities.

If the students have a question why the catalyst is poisoned, then the teacher first explains its action using the theory of the formation of intermediates, and then considers the poisonous effect of impurities.

The acceleration of reactions with the help of a catalyst occurs due to the fact that it forms fragile compounds with the starting materials, and then again separates in a free form. These reactions are much faster than the reaction between sulfur dioxide and oxygen. If there are impurities in the gas mixture that enter into irreversible reactions with the catalyst, then it is poisoned. Despite careful cleaning of gases, the activity of catalysts used in the production of sulfuric acid decreases over time. Its "aging" is caused not only by gradual poisoning, but also by prolonged heating and mechanical destruction, which change the state of the catalyst surface. Not the entire surface of the catalyst participates in the catalyzed reaction, but only its trimmed sections - active centers, and the number of these centers decreases during "aging".

In the previous section, it was discussed how, in the light of the theory of atomic structure, the effect of energy on the excitation of a chemical reaction should be explained to students. This will make it possible to solve the question of why chemical reactions are accelerated when heated. Students know that as the temperature rises in substances, the number of active molecules increases, the speed of movement of molecules and the number of their encounters per unit of time increase. In the atoms of active molecules, electrons are moved to higher energy levels; such molecules are unstable and can more easily react with molecules of other substances.

The theory of electrolytic dissociation explains why the reactions between solutions of acids, salts and bases occur almost instantaneously. The solutions of these substances already contain active particles - oppositely charged ions. Therefore, the reactions between aqueous solutions of acids, salts and bases proceed very quickly and differ significantly from the reactions between the same substances, but taken in dry form.

Starting the lesson on the topic "The rate of a chemical reaction", the teacher reminds that chemical reactions can proceed at different rates, the study of the conditions that affect it is of great practical importance.

How can you measure the rate of a chemical reaction?

Students already know that the rate of chemical transformation can be judged by the amount of a substance that has entered into a reaction or obtained in a certain time, that the speed of mechanical movement is measured by the path that a body travels per unit of time; to calculate this speed, use the formula

where v is speed, S is path and t is time.

With this in mind, students write by analogy the formula for calculating the rate of a chemical reaction

where m is the amount of a substance that has entered into the reaction or is obtained as a result of it during time t.

Consider what the disadvantage of this formula is. It turns out that when using it, the calculated reaction rate will be different even for two portions of the same substance, taken under the same conditions.

Suppose that 15 g of a substance is decomposed in a vessel every second. It turns out that when a partition is introduced into this vessel, which will divide the substance in it into two parts in a ratio of 1: 2, in the first (smaller) part the reaction will proceed at a rate of 5 g / sec, and in the second - 10 g / sec.

In order for the calculated rate to characterize the reaction itself, and not how much of the initial substance is taken, it is necessary to take into account the change in the mass of the reacting substance, referred to the volume, i.e., the change in the concentration of the reacting substance. Therefore, the rate of a chemical reaction can be calculated by the formula:

v = c 0 -c t / t

where c 0 is the initial concentration of any of the reacting substances, c t is the concentration of the same substance after t seconds. When calculating speed, concentration is usually expressed in moles per liter and time in seconds.

In this lesson, attention is paid to the most important ways to accelerate chemical reactions. For this purpose, a laboratory experiment is carried out, showing that the rate of a chemical reaction depends on the concentration of the reacting substances.

For the experiment, the following equipment is used, placed on student tables: 1) a tripod with three test tubes, in one of which there is a crystal of sodium iodide or potassium iodide (2 - 3 pin heads), in the other - a solution of ferric chloride, and the third - empty; 2) a flask or glass of water; 3) two identical glass tubes; 4) glass rod.

The teacher invites students to prepare for the experiment: 1) add water to sodium iodide to form 1/2 test tube of the solution, and mix the liquid with a stick, 2) pour 1/3 of the resulting solution into another test tube, 3) add to the cast solution a test tube with a solution of water so that the volumes of solutions of sodium iodide (or potassium iodide) in the test tubes are the same.

The teacher asks questions to check how students understand the directions:

1) How many times is the sodium iodide solution diluted in the second test tube?

2) How many times is the salt concentration in the first tube higher than in the second?

It is noted that the concentration of one of the solutions is twice the concentration of the other. After that, in two prepared solutions, the reaction of ferric chloride with sodium iodide is carried out, which goes with the release of free iodine:

2NaI + 2FeCl 3 = 2NaCl + 2FeCl 2 + I 2,

2I - + 2Fe 3+ = 2Fe 2+ + I 2.

Students decide in which test tube the rate of interaction of salts is greater and on what basis this can be judged. The assumption is tested by experience.

In both test tubes with solutions of sodium iodide (or potassium iodide) poured first the same amount of starch paste (1-2 ml), and then, after stirring, a few drops of 5-10% ferric chloride solution. It is advisable to add a solution of ferric chloride to both test tubes at the same time. A blue color appears more likely in a test tube with a solution of higher concentration. In the test tube where the concentration of the solution is higher, iodine ions are more often found with ferric ions, and therefore more often interact with them - the reaction proceeds faster.

The teacher shows the burning of sulfur in the air and asks the students how this reaction can be accelerated. Students suggest placing burning sulfur in oxygen and doing this experiment. Based on the analysis of the experiments, a general conclusion is made: the rate of a chemical reaction depends on the concentration of the reacting substances (on the number of ions or molecules per unit volume).

They turn to the question of the effect on the rate of chemical reaction of the surface of the reacting substances. Students recall the reactions with stirring and grinding of reactants: grinding a mixture of ammonia with slaked lime, interaction of small pieces of marble or zinc with hydrochloric acid, burning pulverized fuel in nozzles, using crushed ores in the smelting of metals and pyrite in the production of sulfuric acid. The conditions for roasting pyrite in the production of sulfuric acid are discussed in more detail. To obtain sulfur dioxide, crushed pyrite is used, since it burns faster than pyrite taken in large pieces. The combustion of pulverized pyrite proceeds especially quickly if it is thrown out with a stream of air from the nozzle, as well as when it is burned in a fluidized bed, when the entire surface of the pieces of pyrite comes into contact with air.

It should be borne in mind that chemical reactions with highly crushed combustible substances can be accompanied by an explosion. There have been, for example, explosions of sugar dust in sugar factories.

It is concluded that the more the solid is ground, the greater the rate of the chemical reaction in which it participates.

Then the effect of temperature on the rate of the chemical reaction is analyzed. The same amount of sulfuric acid solution is poured into a test tube with 1/4 of the hyposulfite solution; in parallel with this experiment, heated solutions of hyposulfite and sulfuric acid are drained:

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

The time until the solutions become cloudy are noted. The teacher says that when the temperature rises by 10 ° C, the rate of most reactions increases 2-3 times.

On the basis of the acquired knowledge, students are given the opportunity to explain the acceleration of chemical reactions when substances are heated.

In this lesson, it is not necessary to demonstrate the experience of the catalytic action of substances, since the students got acquainted with it through examples of the decomposition of hydrogen peroxide and the oxidation of sulfur dioxide. They list the catalytic reactions known to them, give definitions of catalysis and catalyst.

To consolidate knowledge in this lesson, they offer questions:

1. What determines the rate of a chemical reaction? Give examples.
2. Under what conditions does the rate of a chemical reaction increase?
3. How, in the light of the theory of electrolytic dissociation, can it be explained that the evolution of hydrogen during the interaction of zinc with acetic acid occurs much more slowly than during the interaction of zinc with hydrochloric acid?
4. What methods can be used to accelerate the reaction of interaction of zinc with hydrochloric acid?
5. Why does a splinter smoldering in air flare up in oxygen?
6. You are given two test tubes in which the interaction of calcium carbonate with hydrochloric acid is slowly going on. Try to speed up the chemical reaction in each tube using different techniques.
7. Why does the rate of a chemical reaction increase with increasing temperature?
8. What methods of accelerating chemical reactions are used in the production of sulfuric acid?
9. List which chemical reactions you know are accelerated by catalysts.

When studying the reaction of ammonia synthesis, students again encounter the use of a catalyst, and, along with consolidating previously obtained information about catalysis and a catalyst, this knowledge can be developed somewhat.

The teacher draws attention to the fact that both reactions - the synthesis of ammonia and its decomposition into nitrogen and hydrogen - proceed in the presence of the same catalyst - reduced iron, which accelerates both direct and reverse reactions to the same extent. Therefore, the catalyst does not shift the chemical equilibrium, but only contributes to the faster achievement of this state. To test their understanding of this position, the teacher asks them questions:

1. Is it possible to obtain ammonia in production from a mixture of nitrogen and hydrogen under high pressure and heating, but without a catalyst? Why?
2. The ammonia synthesis reaction is accelerated by heating and a catalyst. What is the difference in the effect of these conditions on chemical equilibrium?

Introducing students to the synthesis of ammonia in production, the teacher points out that the catalyst quickly loses its activity if the gases (hydrogen and nitrogen) are not previously freed from impurities. In this process, oxygen, water vapor, carbon monoxide, hydrogen sulfide and other sulfur compounds have a toxic effect.

As in the case of the catalytic oxidation of sulfur dioxide to trioxide, in the synthesis of ammonia, the catalyst exerts its accelerating effect only within certain temperature ranges. At temperatures above 600 ° C, reduced iron decreases its catalytic activity.

Using the example of ammonia synthesis, the mechanism of action of the catalyst can be considered. It is noted that iron nitride is formed on the surface of the iron catalyst:

Hydrogen reacts with nitride, ammonia is obtained:

FeN 2 + 3H 2 → Fe + 2NH 3.

Then the process is repeated.

The reactions of the formation of iron nitride and its interaction with hydrogen are very fast.

When studying the reactions of ammonia oxidation, after demonstrating experiments on the combustion of ammonia in oxygen and catalytic oxidation of ammonia, the attention of students is drawn to the fact that the starting materials in these two cases were the same, but depending on the conditions (the use of a catalyst), different products are obtained. ...

Oxidation of ammonia can go with the formation of different substances according to the equations:

4NH 3 + 3O 2 = 2N 2 + 6H 2 O;

4NH 3 + 4O 2 = 2N 2 O + 6H 2 O;

4NH 3 + 5O 2 = 4NO + 6H 2 O.

The catalyst, platinum, accelerates only the last of these reactions. Therefore, using a catalyst, it is possible to direct the interaction of ammonia and oxygen in the desired direction. It finds application in the chemical industry in the production of nitric acid.

Formation in the 9th grade of the concept of chemical production provides great opportunities for acquainting students with the practical control of the rate of chemical reactions in chemical plants.

Based on the generalization of knowledge about previously studied industries (hydrochloric, sulfuric, nitric acids, ammonia), the teacher forms in students the concept of the best conditions for carrying out chemical reactions in production: the use of optimal temperatures, an increase in the concentration of reactants, an increase in the contact surface of reactants, the use of catalysts. After that, in order to identify the circumstances limiting the application of each condition, the students are asked the question: "Is it possible to increase the temperature indefinitely to accelerate chemical reactions in production?" It is found out that strong heating can shift the chemical equilibrium in an undesirable direction, and in the case of using a catalyst, reduce its activity. Taking this into account, in production, not maximum, but optimal temperatures are used.

Other conditions for carrying out chemical reactions in production are analyzed in the same way.

The study of new factual material on chemistry in the IX-X classes is used to further consolidate knowledge about the rate of a chemical reaction.

When studying the properties of white phosphorus, the teacher says that the glow of white phosphorus in the dark indicates its slow oxidation in air. Next, it is considered under what conditions the oxidation of white phosphorus can be accelerated. Heating, fragmentation of phosphorus, the use of oxygen actually accelerate the oxidation of phosphorus, causing it to flash.

Students use their knowledge of ways to accelerate chemical processes to predict the conditions for superphosphate formation. They say that the reaction between tertiary calcium phosphate and sulfuric acid can be accelerated by heating, grinding calcium phosphate, stirring, increasing the concentration of sulfuric acid. The teacher, summarizing what has been said, adds that in

In this production, heating is really used, but for this they use the heat released during the reaction itself, when ground tertiary calcium phosphate is thoroughly mixed with sulfuric acid.

As they study organic substances, students are faced with very many processes that lead with the participation of catalysts, for example, the production of aviation gasoline, rubber, aromatic hydrocarbons.

The role of sulfuric acid in the hydration of ethylene can be considered. In the presence of sulfuric acid, instead of the slow reaction of the addition of water to ethylene (C 2 H 4 + H 2 O → C 2 H 5 OH), the following processes proceed quickly one after the other: 1) sulfuric acid is added to ethylene, forming ethyl sulfate:

2) ethyl sulfate is saponified to form ethyl alcohol and sulfuric acid.

After distilling off the alcohol, sulfuric acid appears in the same amount, but it took part in the formation of an intermediate product. Other examples of the catalytic action of sulfuric acid (the formation of ethylene and ethyl ether from ethyl alcohol) are analyzed independently by the students when doing their homework.

The same substances, with the same catalyst, but at different temperatures, react to form different products. This should be emphasized when familiarizing yourself with the properties of alcohols.

The interaction of carbon monoxide with hydrogen shows that using different catalysts, different organic products can be obtained from the same substances. This interaction can proceed with the formation of methyl alcohol, hydrocarbons or higher alcohols. The desired direction of the interaction of substances is achieved through the use of a catalyst that accelerates the corresponding reaction, but does not significantly affect others. To accelerate the reaction of the formation of methyl alcohol, a mixture of chromium oxides with zinc oxide is used as a catalyst.

After studying hydrocarbons and oxygen-containing organic compounds, to generalize knowledge, students are offered an assignment for independent work in the classroom or at home: choose from such and such a section of the textbook all cases of catalytic reactions, and each student is given only such a part of the textbook material that he can view in the allotted time ...

Analyzing industrial methods of obtaining organic substances, it is useful to draw the attention of students to the fact that the same techniques are used to control the rates of chemical reactions that are used in the production of inorganic

Substances that participate in the reactions and increase its speed, remaining unchanged by the end of the reaction, are called catalysts.

The phenomenon of a change in the rate of reaction under the influence of such substances is called catalysis... The reactions proceeding under the action of catalysts are called catalytic.

In most cases, the effect of a catalyst is explained by the fact that it reduces the activation energy of the reaction. In the presence of a catalyst, the reaction passes through other intermediate stages than without it, and these stages are energetically more accessible. In other words, in the presence of a catalyst, other activated complexes arise, and less energy is required for their formation than for the formation of activated complexes that arise without a catalyst. Thus, the activation energy drops sharply: some molecules, the energy of which was insufficient for active collisions, now turn out to be active.

Intermediates have been studied for a number of reactions; as a rule, they are very active, unstable products.

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. Catalysis can be represented as follows:

A + K = A ... K

A ... K + B = AB + K,

where A ... K is an intermediate activated compound.

Figure 13.5 - Image of the reaction path of non-catalytic A + B → AB reaction (curve 1) and homogeneous catalytic reaction (curve 2).

In the chemical industry, catalysts are widely used. Under the influence of catalysts, reactions can be accelerated by a factor of millions or more. In some cases, under the action of catalysts, such reactions can be initiated, which practically do not proceed without them under these conditions.

Distinguish homogeneous and heterogeneous catalysis.

When homogeneous catalysis the catalyst and reactants form one phase (gas or solution). When heterogeneous catalysis the catalyst is in the system as an independent phase.

Examples of homogeneous catalysis:

1) oxidation of SO 2 + 1 / 2O 2 = SO 3 in the presence of NO; NO readily oxidizes to NO 2, and NO 2 already oxidizes SO 2;

2) decomposition of hydrogen peroxide in an aqueous solution into water and oxygen: ions Cr 2 O 2 = 7, WO 2-4, MoO 2-4, catalyzing the decomposition of hydrogen peroxide, form intermediate compounds with it, which further decompose with the release of oxygen.

Homogeneous catalysis is carried out through intermediate reactions with a catalyst, and as a result, one reaction with a high activation energy is replaced by several, for which the activation energies are lower, their rate is higher:

CO + 1 / 2O 2 = CO 2 (catalyst - water vapor).

Heterogeneous catalysis is widely used in the chemical industry. Most of the products currently produced by this industry come from heterogeneous catalysis. In heterogeneous catalysis, the reaction proceeds on the catalyst surface. Hence it follows that the activity of the catalyst depends on the size and properties of its surface. In order to have a large ("developed") surface, the catalyst must have a porous structure or be in a highly fragmented (highly dispersed) state. In practical application, the catalyst is usually applied to a carrier having a porous structure (pumice, asbestos, etc.).

As in the case of homogeneous catalysis, in heterogeneous catalysis, the reaction proceeds through active intermediates. But here these compounds are surface compounds of the catalyst with reactants. Passing through a series of stages in which these intermediates are involved, the reaction ends with the formation of final products, and the catalyst is not consumed as a result.

All catalytic heterogeneous reactions include adsorption and desorption steps.

The catalytic action of the surface is reduced to two factors: an increase in the concentration at the interface and the activation of adsorbed molecules.

Examples of heterogeneous catalysis:

2H 2 O = 2H 2 O + O 2 (catalyst - MnO 2);

H 2 + 1/2 O 2 = H 2 O (catalyst - platinum).

Catalysis plays a very important role in biological systems. Most of the chemical reactions that take place in the digestive system, in the blood and in the cells of animals and humans, are catalytic reactions. The catalysts, in this case called enzymes, are simple or complex proteins. So, saliva contains the enzyme ptyalin, which catalyzes the conversion of starch into sugar. An enzyme found in the stomach, pepsin, catalyzes the breakdown of proteins. There are about 30,000 different enzymes in the human body: each of them serves as an effective catalyst for the corresponding reaction.

Catalysts are substances that can accelerate a chemical reaction, while the catalysts themselves are not consumed in the chemical reaction. Found that catalysts change the mechanism of a chemical reaction. In this case, other, new transition states appear, characterized by a lower height of the energy barrier. Thus, under the action of the catalyst, the

the activation energy of the process (Fig. 3). Entering into various kinds of interactions with intermediate particles, the catalysts remain unchanged at the end of the reaction. The catalysts affect only thermodynamically allowed reactions. The catalyst cannot cause a reaction, because does not affect its driving forces. The catalyst does not affect the chemical equilibrium constant, because equally reduces the activation energy of both direct and reverse reactions.

Fig.3 Energy diagram of the course of the reaction A + B = AB a) without a catalyst and b) in the presence of a catalyst. Ea is the activation energy of a non-catalytic reaction; Еа 1 and Еа 2 - activation energy of the catalytic reaction; AK - intermediate reactive compound of the catalyst with one of the reagents; A ... K, AK ... B - activated complexes of the catalytic reaction; А ... В - activated complex of non-catalytic reaction; ∆E cat. - a decrease in the activation energy under the influence of the catalyst.

Distinguish between homogeneous and heterogeneous catalysis. In the first case, the catalyst is in the same phase with the reactants, and in the second, the catalyst is a solid, on the surface of which a chemical reaction takes place between the reactants.

Chemical equilibrium

Chemical reactions are usually subdivided into reversible and irreversible. Irreversible chemical reactions proceed until the complete consumption of at least one of the initial substances, i.e. the reaction products either do not interact with each other at all, or form substances that are different from the original ones. There are very few such reactions. For example:

2KСlO 3 (tv) = 2KCl (tv) + 3О 2 (g)

In electrolyte solutions, reactions involving the formation of precipitates, gases and weak electrolytes (water, complex compounds) are considered to be practically irreversible.

Most chemical reactions are reversible, i.e. they go both forward and backward. This becomes possible when the activation energies of the forward and reverse processes differ insignificantly from each other, and the reaction products are capable of converting into initial substances. For example, the HI synthesis reaction is a typically reversible reaction:

H 2 (g) + I 2 (g) ⇄ 2HI (g)

The law of mass action (expression of the reaction rate) for the forward and reverse processes, respectively, will have the form: = ∙; = 2

At some point in time, a state occurs when the rates of the forward and reverse reactions become equal = (fig. 4).

Fig. 4 Changes in the velocities of the forward (and reverse (reactions over time t

This state is called chemical equilibrium. It is dynamic (mobile) in nature and can shift in one direction or another depending on changes in external conditions. Starting from the moment of equilibrium, under constant external conditions, the concentrations of the initial substances and reaction products do not change over time. The concentrations of the reagents corresponding to the equilibrium state are called equilibrium... To determine the equilibrium concentration of a reagent, it is necessary to subtract from its initial concentration the amount of a substance that has reacted by the time of the onset of an equilibrium state: WITH equal = C ref. - WITH proreagir... The amount of reagents that entered into the reaction and formed from them by the time of equilibrium of the products is proportional to the stoichiometric coefficients in the reaction equation.

The state of equilibrium under unchanged external conditions can exist for an arbitrarily long time. In a state of equilibrium

∙ = [2, whence / [= 2 / ∙.

At a constant temperature, the rate constants of the forward and reverse process are constant values.

The ratio of two constants is also the value of the constant K = / and is called chemical equilibrium constant... It can be expressed

either through the concentrations of the reactants =, or through their partial pressures if the reaction proceeds with the participation of gases.

In the general case, for the reaction aA + bB +… ⇄cC + dD +… the chemical equilibrium constant is equal to the ratio of the product of the concentrations of the reaction products to the product of the concentrations of the starting substances in powers equal to their stoichiometric coefficients.

The constant of chemical equilibrium does not depend on the path of the process and determines the depth of its course by the time the equilibrium state is reached. The larger this value, the greater the degree of conversion of reagents into products.

The chemical equilibrium constant, as well as the reaction rate constants, are functions only of the temperature and the nature of the reacting substances and do not depend on their concentration.

For heterogeneous processes, the concentration of solids is not included in the expression for the reaction rate and chemical equilibrium constant, because the reaction proceeds on the surface of the solid phase, the concentration of which remains constant over time. For example, for a reaction:

FeO (tv) + CO (g) ⇄ Fe (tv) + CO 2 (g)

the expression for the equilibrium constant will be:

K p and K with are related by the relation K p = K c (RT) ∆ n, where n = n prod. -n of original substances - change in the number of moles gaseous substances during the reaction. For this reaction, K p = K s, since n gaseous substances is equal to zero.

Why do catalysts increase the rate of a chemical reaction? It turns out that they act in full accordance with the popular wisdom: "The clever will not go up the hill, the clever will bypass the mountain." In order for substances to begin to interact, their particles (molecules, atoms, ions) must be given a certain energy, called the activation energy (Fig. 13, a). Catalysts lower this energy by combining with one of the reacting substances and conducting it along the "energy mountain" to meet with another substance with less energy. Therefore, in the presence of a catalyst, chemical reactions proceed not only faster, but also at a lower temperature, which reduces the cost of production processes.

Rice. 13.
Energy diagrams of catalytic reactions using conventional (a) and selective (b) catalysts

And not only. The use of catalysts can lead to the fact that the same substances will react in different ways, that is, with the formation of different products (Fig. 13, b). For example, ammonia is oxidized by oxygen to nitrogen and water, and in the presence of a catalyst - to nitric oxide (II) and water (write down the reaction equations and consider the oxidation and reduction processes).

The process of changing the rate of a chemical reaction or the path it takes is called catalysis. As well as reactions, homogeneous and heterogeneous types of catalysis are distinguished. In the case of using enzymes, catalysis is called enzymatic. This type of catalysis has been known to man since ancient times. Thanks to the enzymatic breakdown of organic substances, man learned how to bake bread, brew beer, make wine and cheese (Fig. 14).

Rice. fourteen.
Since ancient times, man has used catalysis, which occurs when baking bread, brewing beer, making wine, making cheese.

The most famous in everyday life are enzymes that are part of washing powders. It is they that allow you to rid the laundry of stains and unpleasant odors during washing.

Let's take a closer look at catalysts using a chemical experiment.

Hydrogen peroxide (in everyday life it is often called hydrogen peroxide) is a drug necessary in any home medicine cabinet (Fig. 15).

Rice. 15.
Hydrogen peroxide solution

The expiration date must be indicated on the packaging with this drug, since it decomposes during storage:

However, under normal conditions, this process proceeds so slowly that we do not notice the release of oxygen, and only by opening a bottle in which hydrogen peroxide has been stored for a long time, you can see how a little gas is released from it. How can this process be accelerated? Let's carry out a laboratory experiment.

Laboratory experiment No. 9 Decomposition of hydrogen peroxide using manganese (IV) oxide

Laboratory experiment No. 10
Detection of catalase in food

Catalysts not only make production processes more economical, but also make a significant contribution to environmental protection. Thus, modern passenger cars are equipped with a catalytic device, inside of which there are cellular ceramic catalyst carriers (platinum and rhodium). Passing through them, harmful substances (oxides of carbon, nitrogen, unburned gasoline) are converted into carbon dioxide, nitrogen and water (Fig. 16).

Rice. 16.
A catalytic converter in a car that converts the nitrogen oxides of its exhaust gases into harmless nitrogen

However, for chemical reactions, not only catalysts are important, which speed up the reaction, but also substances that can slow them down. Such substances are called inhibitors. The best known are metal corrosion inhibitors.

Laboratory experiment No. 11
Inhibition of the interaction of acids with metals with urotropin

In the vocabulary of an ordinary person, words are often found that are borrowed from chemistry. For example, antioxidants, or antioxidants. What are substances called antioxidants? You have probably noticed that if you store butter for a long time, then it changes color, taste, acquires an unpleasant odor - it oxidizes in air. To prevent food spoilage, antioxidants are added to them. They play an important role in maintaining human health, because unwanted oxidation processes also occur in his body, as a result of which a person gets sick, gets tired and ages faster. The human body receives antioxidants by eating foods containing, for example, carotene (vitamin A) and vitamin E (Fig. 17).

Rice. 17.
Antioxidants: a - β-carotene; b - vitamin E

So, the rate of a chemical reaction can be controlled with the help of catalysts and inhibitors, changes in temperature, concentration of reactants, pressure (for homogeneous gas reactions), contact area of ​​reactants (for heterogeneous processes). And of course, the rate of chemical reactions depends on the nature of the reacting substances.

New words and concepts

1. Catalysts.
2. Enzymes.
3. Catalysis (homogeneous, heterogeneous, enzymatic).
4. Inhibitors.
5. Antioxidants

Self-study assignments

1. What are catalysts? What role do they play in chemical reactions? Why do catalysts accelerate chemical reactions?
2. What role did enzymatic catalysis play in the history of human civilization?
3. Prepare a report on the role of catalysts in modern manufacturing.
4. Prepare a report on the role of inhibitors in modern manufacturing.
5. Prepare a paper on the role of antioxidants in medicine and the food industry.