Definition of catalysis. Catalytic reactions: examples. Homogeneous and heterogeneous catalysis. How do catalysts work?

Catalysis has found wide application in the chemical industry, in particular in the technology of inorganic substances. Catalysis– excitation of chemical reactions or changes in their speed under the influence of substances - catalysts, which repeatedly enter into chemical interaction with reaction participants and restore their chemical composition after each cycle of interaction. There are substances that reduce the rate of reaction, called inhibitors or negative catalysts. Catalysts do not change the state of equilibrium in the system, but only facilitate its achievement. A catalyst can simultaneously accelerate both forward and reverse reactions, but the equilibrium constant remains constant. In other words, the catalyst cannot change the equilibrium of thermodynamically unfavorable reversible reactions in which the equilibrium is shifted towards the starting substances.

The essence of the accelerating effect of catalysts is to reduce the activation energy Ea of a chemical reaction by changing the reaction path in the presence of a catalyst. For the reaction of converting A into B, the reaction path can be represented as follows:

A + K  AK

VK  V + K

As can be seen from Figure 1, the second stage of the mechanism is limiting, since it has the highest activation energy E cat, but significantly lower than for the non-catalytic process E necat. The activation energy decreases due to the compensation of the energy of breaking the bonds of the reacting molecules with the energy of the formation of new bonds with the catalyst. A quantitative characteristic of the decrease in activation energy, and therefore the efficiency of the catalyst, can be the degree of compensation for the energy of broken bonds Di:

 = (Di – E cat)/Di (1)

The lower the activation energy of the catalytic process, the higher the degree of compensation.

Simultaneously with the decrease in activation energy, in many cases there is a decrease in the order of the reaction. The decrease in the reaction order is explained by the fact that in the presence of a catalyst, the reaction proceeds through several elementary stages, the order of which may be less than the order of non-catalytic reactions.

Types of catalysis

Based on the phase state of the reagents and catalyst, catalytic processes are divided into homogeneous and heterogeneous. In homogeneous catalysis, the catalyst and reactants are in the same phase (gas or liquid); in heterogeneous catalysis, they are in different phases. Often, the reacting system of a heterogeneous catalytic process consists of three phases in various combinations, for example, the reactants can be in the gas and liquid phases, and the catalyst can be in the solid phase.

A special group includes enzymatic (biological) catalytic processes, common in nature and used in industry for the production of feed proteins, organic acids, alcohols, as well as for wastewater treatment.

Based on the types of reactions, catalysis is divided into redox and acid-base. In reactions proceeding according to the redox mechanism, intermediate interaction with the catalyst is accompanied by homolytic cleavage of two-electron bonds in the reacting substances and the formation of bonds with the catalyst at the site of the latter's unpaired electrons. Typical catalysts for redox reactions are metals or oxides of variable valence.

Acid-base catalytic reactions occur as a result of intermediate protolytic interaction of the reactants with the catalyst or interaction involving a lone pair of electrons (heterolytic) catalysis. Heterolytic catalysis proceeds with a rupture of the covalent bond in which, unlike homolytic reactions, the electron pair performing the bond remains in whole or in part with one of the atoms or a group of atoms. Catalytic activity depends on the ease of transfer of a proton to the reagent (acid catalysis) or abstraction of a proton from the reagent (base catalysis) in the first act of catalysis. According to the acid-base mechanism, catalytic reactions of hydrolysis, hydration and dehydration, polymerization, polycondensation, alkylation, isomerization, etc. occur. Active catalysts are compounds of boron, fluorine, silicon, aluminum, sulfur and other elements with acidic properties, or compounds of elements of the first and the second groups of the periodic table, which have basic properties. The hydration of ethylene by the acid-base mechanism with the participation of the acid catalyst NA is carried out as follows: in the first stage, the catalyst serves as a proton donor

CH 2 =CH 2 + HA  CH 3 -CH 2 + + A -

the second stage is the actual hydration

CH 3 -CH 2 + + HON  CH 3 CH 2 OH + H +

third stage – catalyst regeneration

N + + A -  NA.

Redox and acid-base reactions can be considered according to the radical mechanism, according to which the strong molecule-catalyst lattice bond formed during chemisorption promotes the dissociation of the reacting molecules into radicals. In heterogeneous catalysis, free radicals, migrating over the surface of the catalyst, form neutral product molecules that are desorbed.

There is also photocatalysis, where the process is initiated by exposure to light.

Since heterogeneous catalysis on solid catalysts is most common in inorganic chemistry, we will dwell on it in more detail. The process can be divided into several stages:

1) external diffusion of reacting substances from the core of the flow to the surface of the catalyst; in industrial devices, turbulent (convective) diffusion usually predominates over molecular;

2) internal diffusion in the pores of the catalyst grain, depending on the size of the catalyst pores and the size of the reagent molecules, diffusion can occur by the molecular mechanism or by the Knudsen mechanism (with constrained movement);

3) activated (chemical) adsorption of one or more reactants on the surface of the catalyst with the formation of a surface chemical compound;

4) rearrangement of atoms to form a surface product-catalyst complex;

5) desorption of the catalysis product and regeneration of the active center of the catalyst; for a number of catalysts, not its entire surface is active, but individual areas - active centers;

6) diffusion of the product in the pores of the catalyst;

7) diffusion of the product from the surface of the catalyst grain into the gas flow.

The overall rate of a heterogeneous catalytic process is determined by the rates of individual stages and is limited by the slowest of them. Speaking about the stage limiting the process, it is assumed that the remaining stages proceed so quickly that in each of them equilibrium is practically achieved. The speeds of individual stages are determined by the parameters of the technological process. Based on the mechanism of the process as a whole, including the catalytic reaction itself and the diffusion stages of substance transfer, processes occurring in the kinetic, external diffusion and internal diffusion regions are distinguished. The speed of the process in the general case is determined by the expression:

d/d = k c (2)

where c is the driving force of the process, equal to the product of the effective concentrations of the reactants; for a process occurring in the gas phase, the driving force is expressed in partial pressures of the reactants p; k is the rate constant.

In general, the rate constant depends on many factors:

k = f (k 1 , k 2 , k sub, …..D and, D and / , D p, ….) (3)

where k 1, k 2, k inc are the rate constants of the direct, reverse and side reactions; D and, D and /, D p are the diffusion coefficients of the starting substances and the product, which determine the value of k in the external or internal diffusion regions of the process.

IN kinetic region k does not depend on diffusion coefficients. The general kinetic equation for the rate of a gas catalytic process, taking into account the influence of the main parameters of the technological regime on the rate:

u = kvpP n  0 = k 0 e -Ea/RT vpP n  0 (4)

where v is the gas flow rate, p is the driving force of the process at P0.1 MPa (1 at), P is the ratio of operating pressure to normal atmospheric pressure, that is, a dimensionless quantity,  0 is the conversion factor to normal pressure and temperature, n - reaction order.

The mechanism of chemical stages is determined by the nature of the reactants and catalyst. The process can be limited by chemisorption of one of the reactants by the surface of the catalyst or desorption of reaction products. The rate of the reaction can be controlled by the formation of a charged activated complex. In these cases, charging the catalyst surface under the influence of some factors has a significant impact on the course of the reaction. In the kinetic region, processes occur mainly on low-activity, fine-grained catalysts with large pores in a turbulent flow of reagents, as well as at low temperatures close to the ignition temperatures of the catalyst. For reactions in liquids, the transition to the kinetic region can also occur with increasing temperature due to a decrease in the viscosity of the liquid and, consequently, an acceleration of diffusion. With increasing temperature, the degree of association, solvation, and hydration of reagent molecules in solutions decreases, which leads to an increase in diffusion coefficients and, accordingly, a transition from the diffusion region to the kinetic region. Reactions whose overall order is higher than unity are characterized by a transition from the diffusion region to the kinetic region with a significant decrease in the concentration of the initial reagents. The transition of the process from the kinetic region to the external diffusion region can occur with a decrease in the flow rate, an increase in concentration, and an increase in temperature.

In external diffusion region First of all, processes take place on highly active catalysts, which provide a fast reaction and sufficient product yield during the contact time of the reagents with the catalysts, measured in fractions of a second. The very fast reaction takes place almost entirely on the outer surface of the catalyst. In this case, it is not advisable to use porous grains with a highly developed internal surface, but one must strive to develop the outer surface of the catalyst. Thus, during the oxidation of ammonia on platinum, the latter is used in the form of extremely fine meshes containing thousands of interweavings of platinum wire. The most effective means of accelerating processes occurring in the region of external diffusion is mixing of reagents, which is often achieved by increasing the linear speed of the reagents. Strong turbulization of the flow leads to a transition of the process from the external diffusion region to the internal diffusion region (with coarse-grained, fine-porous catalysts) or to the kinetic region.

where G is the amount of substance transferred over time  in the x direction perpendicular to the surface of the catalyst grain at concentration c of the diffusing component in the core of the reagent flow, S is the free outer surface of the catalyst, dc/dx is the concentration gradient.

A large number of methods and equations have been proposed for determining the diffusion coefficients of substances in various media. For a binary mixture of substances A and B according to Arnold

where T is temperature, K; M A, M B - molar masses of substances A and B, g/mol; v A, v B - molar volumes of substances; P - total pressure (0.1 M Pa); C A+B is the Sutherland constant.

The Sutherland constant is:

C A+B = 1.47(T A / +T B /) 0.5 (7)

G
de T A /, T B / - boiling temperatures of components A and B, K.

For gases A and B with close values ​​of molar volumes, we can take =1, and if there is a significant difference between them, 1.

The diffusion coefficient in liquid media D g can be determined by the formula

where  is the viscosity of the solvent, PaC; M and v are the molar mass and molar volume of the diffusing substance; xa is a parameter that takes into account the association of molecules in the solvent.

In intradiffusion region, that is, when the overall rate of the process is limited by the diffusion of reagents in the pores of the catalyst grain, there are several ways to accelerate the process. It is possible to reduce the size of the catalyst grains and, accordingly, the path of the molecules to the middle of the grain; this is possible if they move simultaneously from the filter layer to the boiling one. It is possible to produce large-porous catalysts for a fixed layer without reducing the grain size to avoid an increase in hydraulic resistance, but this will inevitably reduce the internal surface and, accordingly, reduce the intensity of the catalyst compared to a fine-grained, large-porous catalyst. You can use a ring-shaped contact mass with a small wall thickness. Finally, bidisperse or polydisperse catalysts, in which large pores are transport routes to the highly developed surface created by thin pores. In all cases, they strive to reduce the depth of penetration of reagents into the pores (and products from the pores) so much as to eliminate intra-diffusion inhibition and move into the kinetic region, when the rate of the process is determined only by the rate of the actual chemical acts of catalysis, that is, the adsorption of reagents by active centers, the formation of products and its desorption. Most industrial processes occurring in the filter bed are inhibited by internal diffusion, for example large-scale catalytic processes of methane-steam reforming, carbon monoxide conversion, ammonia synthesis, etc.

The time  required for the diffusion of a component into the pores of the catalyst to a depth l can be determined using the Einstein formula:

 = l 2 /2D e (10)

The effective diffusion coefficient in pores is determined approximately depending on the ratio of pore sizes and the free path of molecules. In gaseous media, when the mean free path of a component molecule  is less than the equivalent pore diameter d=2r (2r), it is assumed that normal molecular diffusion occurs in the pores D e =D, which is calculated by the formula:

In a constrained mode of movement, when 2r, D e =D k is determined using the approximate Knudsen formula:

(
12)

where r is the transverse radius of the pore.

(
13)

Diffusion in the pores of a catalyst in liquid media is very difficult due to a strong increase in the viscosity of the solution in narrow channels (abnormal viscosity), therefore, dispersed catalysts, that is, small non-porous particles, are often used for catalysis in liquids. In many catalytic processes, with changes in the composition of the reaction mixture and other process parameters, the mechanism of catalysis, as well as the composition and activity of the catalyst, can change, so it is necessary to take into account the possibility of changing the nature and speed of the process even with a relatively small change in its parameters.

Catalysts can increase the reaction rate constant indefinitely, but unlike temperature, catalysts do not affect the rate of diffusion. Therefore, in many cases, with a significant increase in the reaction rate, the overall rate remains low due to the slow supply of components to the reaction zone.

Catalysis

The term "catalysis" was introduced in 1835 by the Swedish scientist Jons Jakob Berzelius.

The phenomenon of catalysis is widespread in nature (most processes occurring in living organisms are catalytic) and is widely used in technology (in oil refining and petrochemistry, in the production of sulfuric acid, ammonia, nitric acid, etc.). Most of all industrial reactions are catalytic.

Basic principles of catalysis

The catalyst changes the reaction mechanism to an energetically more favorable one, that is, it reduces the activation energy. The catalyst forms an intermediate compound with a molecule of one of the reagents, in which the chemical bonds are weakened. This makes it easier to react with the second reagent. It is important to note that catalysts accelerate reversible reactions, both forward and reverse.

Types of catalysis

Based on their effect on the reaction rate, many sources of catalysis are divided into positive (the reaction rate increases) and negative (the reaction rate decreases). In the latter case, an inhibition process occurs, which cannot be considered “negative catalysis”, since the inhibitor is consumed during the reaction.

Catalysis happens homogeneous And heterogeneous(contact). In homogeneous catalysis, the catalyst is in the same phase as the reaction reagents, while heterogeneous catalysts differ in phase.

Homogeneous catalysis

An example of homogeneous catalysis is the decomposition of hydrogen peroxide in the presence of iodine ions. The reaction occurs in two stages:

H 2 O 2 + I → H 2 O + IO H 2 O 2 + IO → H 2 O + O 2 + I

In homogeneous catalysis, the action of the catalyst is due to the fact that it interacts with reacting substances to form intermediate compounds, this leads to a decrease in activation energy.

Heterogeneous catalysis

In heterogeneous catalysis, the acceleration of the process usually occurs on the surface of a solid body - the catalyst, therefore the activity of the catalyst depends on the size and properties of its surface. In practice, the catalyst is usually supported on a solid porous support.

The mechanism of heterogeneous catalysis is more complex than that of homogeneous catalysis. The mechanism of heterogeneous catalysis includes five stages, all of which are reversible.

1. Diffusion of reactants to the surface of a solid
2. Physical adsorption on the active centers of the surface of a solid substance of reacting molecules and then their chemisorption
3. Chemical reaction between reacting molecules
4. Desorption of products from the catalyst surface
5. Diffusion of product from the catalyst surface into the general flow

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 (contact method).

Catalyst carrier

Metal platinum (shown by arrows), stabilized on a carrier - aluminum oxide

catalyst carrier, otherwise substrate (catalyst) (English carrier or support) - an inert or low-active material that serves to stabilize particles of the active catalytic phase on its surface.

The role of the support in heterogeneous catalysis is to prevent agglomeration or sintering of the active component, which allows maintaining a high contact area between the active substance (see active catalytic phase) and the reactants. The amount of carrier is usually much greater than the amount of active component applied to it. The main requirements for carriers are large surface area and porosity, thermal stability, chemical inertness, and high mechanical strength. In some cases, the carrier affects the properties of the active phase (the “strong metal–carrier interaction” effect). Both natural (clays, pumice, diatomaceous earth, asbestos, etc.) and synthetic materials (active carbons, silica gel, aluminosilicates, oxides of aluminum, magnesium, zirconium, etc.) are used as carriers.

Chemistry of catalysis

Chemistry of catalysis studies substances that change the rate of chemical reactions. Substances that slow down reactions are called inhibitors. Enzymes- These are biological catalysts. The catalyst is not in a stoichiometric relationship with the products and is regenerated after each cycle of converting reactants into products. Despite the emergence of new methods of activating molecules (plasma chemistry, radiation and laser effects, and others), catalysis is the basis of chemical production (the relative share of catalytic processes is 80-90%).

The reaction that fed humanity (the solution to the fixed nitrogen problem) is the Haber-Bosch cycle. Ammonia is produced with a catalyst - porous iron. Occurs at P = 30 MPa and T = 420-500 °C

3H 2 + N 2 = 2NH 3

Hydrogen for the synthesis of NH 3 is obtained by two sequential catalytic processes: the conversion of CH 4 (CH 4 + H 2 O → CO + 3H 2) on Ni − catalysts and the conversion of the resulting carbon monoxide (CO + H 2 O → CO 2 + H 2) . To achieve high degrees of conversion, the last reaction is carried out in two stages: high temperature (315-480 °C) - on Fe - Cr - oxide catalysts and low temperature (200-350 °C) - on Cu - Zn - oxide catalysts. Ammonia is used to produce nitric acid and other nitrogen compounds - from drugs and fertilizers to explosives.

There are different types of catalysis homogeneous, heterogeneous, interfacial, micellar, enzymatic.

The activation energy E of catalytic reactions is significantly lower than for the same reaction in the absence of a catalyst. For example, for the non-catalytic decomposition of NH 3 into N 2 + H 2 E ~ 320 kJ/mol, for the same decomposition in the presence of Pt E ~ 150 kJ/mol. Thanks to the decrease in E, catalytic reactions are accelerated compared to non-catalytic ones.

Literature

• Boreskov G.K. Catalysis. Questions of theory and practice. - Novosibirsk, 1987.
• Gates B. Chemistry of catalytic processes / B. Gates, J. Ketsir.
• Journal "Kinetics and Catalysis".
• Kolesnikov I. M. Catalysis and production of catalysts. - M.: Technology, 2004. - 399 p.
• Shuyt G.- M.: Mir, 1981. - 551 p.
• Yablonsky G. S., Bykov V. I., Gorban A. N. Kinetic models of catalytic reactions. - Novosibirsk: Science (Siberian Branch), 1983. - 255 p.

see also

Links

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See what “Catalysis” is in other dictionaries:

CATALYSIS- CATALYSIS, catalysts. Catalyst is a name introduced into science by Berzelius (Berzelius; 1835) to designate substances that cause or accelerate chemical reactions. processes without taking visible participation in them. Later Ostwald (SY a1s1) and his school... ... Great Medical Encyclopedia

- (from the Greek katalysis destruction) acceleration of a chemical reaction in the presence of catalyst substances that interact with reagents, but are not consumed in the reaction and are not part of the products. In homogeneous catalysis, the starting reagents and... ... Big Encyclopedic Dictionary

CATALYSIS, changing the rate of a chemical reaction by adding a CATALYST substance that does not participate in the reaction. Catalytic action makes it possible to clarify the reaction mechanism; used in many industrial processes... Scientific and technical encyclopedic dictionary

- (from the Greek katalysis destruction), acceleration of a chemical reaction in the presence of a catalyst substance that interacts with the reagents, but is not consumed in the reaction and is not part of the final products. Use of catalysts... ... Modern encyclopedia

CATALYSIS, catalysis, man. (from Greek katalysis dissolution) (chem.). Acceleration or slowdown of a chemical reaction under the influence of catalysts. Ushakov's explanatory dictionary. D.N. Ushakov. 1935 1940 ... Ushakov's Explanatory Dictionary

Noun, number of synonyms: 4 autocatalysis (2) biocatalysis (1) photocatalysis (1) ... Synonym dictionary

Acceleration or deceleration of chemical reactions with the help of some specifically active substances (catalysts) that can repeatedly enter into short-term interaction with the reacting compounds, facilitating the course of the reaction. The essence of action... ... Geological encyclopedia

catalysis- a, m. catalyse f. gr. catalysis termination. Changes in the rate of a chemical reaction under the influence of certain substances (catalysts). BAS 1. Borrowed from French. language in 1837. First recorded in the Mining Journal of 1837 (2 5 380) translated... ... Historical Dictionary of Gallicisms of the Russian Language

catalysis- - Topics of biotechnology EN catalysis ... Technical Translator's Guide

catalysis- - changing the speed or excitation of a chemical reaction by catalyst substances. General chemistry: textbook / A. V. Zholnin ... Chemical terms

Catalysis- [gr. katalysis destruction] the phenomenon of an increase in the rate of chemical reactions in the presence of a substance that does not undergo changes during the reaction. [Usherov Marshak A.V. Concrete science: lexicon. M.: RIF Building Materials. 2009. – 112 pp.]… … Encyclopedia of terms, definitions and explanations of building materials

Katamliz- selective acceleration of one of the possible thermodynamically allowed directions of a chemical reaction under the action of a catalyst(s), which repeatedly enters into intermediate chemical interactions with reaction participants and restores its chemical composition after each cycle of intermediate chemical interactions. The term "catalysis" was introduced in 1835 by the Swedish scientist Jons Jakob Berzelius.

The phenomenon of catalysis is widespread in nature (most processes occurring in living organisms are catalytic) and is widely used in technology (in oil refining and petrochemistry, in the production of sulfuric acid, ammonia, nitric acid, etc.). Most of all industrial reactions are catalytic.

Catalysts substances that change the rate of chemical reactions are called.

Some catalysts greatly accelerate the reaction - positive catalysis, or simply catalysis, while others slow down - negative catalysis. Examples of positive catalysis include the production of sulfuric acid, the oxidation of ammonia into nitric acid using a platinum catalyst, etc.

Based on their effect on the reaction rate, many sources of catalysis are divided into positive (the reaction rate increases) and negative (the reaction rate decreases). In the latter case, an inhibition process occurs, which cannot be considered “negative catalysis”, since the inhibitor is consumed during the reaction.

Catalysis can be homogeneous and heterogeneous (contact). In homogeneous catalysis, the catalyst is in the same phase as the reaction reagents, while heterogeneous catalysts differ in phase.

Homogeneous catalysis.

Example homogeneous catalysis is the decomposition of hydrogen peroxide in the presence of iodine ions. The reaction occurs in two stages:

H 2 O2+ I > H2O+IO, H2O2+IO> H2O + O2+I

In homogeneous catalysis, the action of the catalyst is due to the fact that it interacts with reacting substances to form intermediate compounds, this leads to a decrease in activation energy.

Heterogeneous catalysis.

In heterogeneous catalysis, the acceleration of the process usually occurs on the surface of a solid body—the catalyst; therefore, the activity of the catalyst depends on the size and properties of its surface. In practice, the catalyst is usually supported on a solid porous support.

The mechanism of heterogeneous catalysis is more complex than that of homogeneous catalysis. The mechanism of heterogeneous catalysis includes five stages, all of which are reversible.

• 1. Diffusion of reactants to the surface of a solid
• 2. Physical adsorption on the active centers of the surface of a solid substance of reacting molecules and then their chemisorption
• 3. Chemical reaction between reacting molecules
• 4. Desorption of products from the catalyst surface
• 5. Diffusion of the product from the surface of the catalyst into the general flow

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 (contact method).

Most catalytic reactions are carried out on porous catalysts, the inner surface of which consists of pores and channels of different sizes and lengths. These pores can be isolated or connected to each other. The main factor determining the speed and nature of the movement of gases in the pores of the catalyst is the pore size. The speed of free movement of molecules can reach 1000 m/s, and the inhibition of movement in pores is associated with collisions between gas molecules and with pore walls.

Most catalytic reactions are nonselective, which imposes certain limitations on kinetic analysis methods.

Most catalytic reactions involve several different types of atoms and molecules. Determining the reaction mechanism and the nature of the forces acting between these atoms and molecules and between them and the surface is naturally a complex problem, but it can be simplified by studying the adsorption behavior of one type of atom or molecule. Such studies have shown that when certain molecules are adsorbed on certain adsorbents, the bond in the molecule is broken and two bonds with the adsorbent are formed; in this case, the adsorbed molecule transforms into two adsorbed atoms. This process is a surface chemical reaction, and the resulting adsorbed atoms are usually called chemisorbed atoms. If at sufficiently low temperatures such a reaction does not occur and the adsorbed molecules do not disintegrate into two adsorbed atoms, then such molecules are called physically adsorbed.

Catalysis is one of the most common methods in chemistry for accelerating chemical reactions.

Catalysis are called selective changes in the rate of chemical reactions in the presence of substances (catalysts) that, taking part in intermediate processes, are regenerated during the reaction and are not part of the final products.

Positive catalysis or simply catalysis, - this is a significant increase in the rate of a reaction, for example, the production of sulfuric acid or the oxidation of ammonia in the presence of platinum. Negative catalysis, or inhibition, - this is a slowdown of a reaction, for example, the interaction of a sodium sulfite solution with atmospheric oxygen in the presence of ethyl alcohol or the decomposition of hydrogen peroxide at low concentrations of sulfuric acid (inhibitors are ethyl alcohol and sulfuric acid, respectively).

Reactions that occur under the influence of catalysts are called catalytic.

The action of a catalyst in the process of changing a chemical system can be not only accelerating but also orienting: if the initial chemical system can, under given conditions, develop in several thermodynamically possible directions, the catalyst preferentially accelerates one of them.

Catalysis changes the reaction mechanism. The catalyst and one of the starting materials form activated complex- an intermediate compound that reacts with another starting material to form reaction products and regenerate catalyst molecules.

Let some reaction A + B = AB have a very high activation energy E a and therefore proceeds slowly. Its energy diagram is shown in Fig. 4.4, A.

Rice. 4.4. Change in enthalpy during the reaction: a - without a catalyst: b- with catalyst

If this reaction is carried out in the presence of catalyst K (Fig. 4.4, b), then it enters into a chemical interaction with one of the starting substances (for example A), as a result of which, through the activated AKB* complex, a fragile chemical compound AK is formed according to the reaction A + K = AK. The activation energy of this process E" less than that in the absence of a catalyst (E a "therefore, the reaction proceeds quickly. Next, the intermediate compound AK, through another activated complex, ABC*, interacts with the second starting material B: AK + B = AB + K; in this case, the catalyst returns to its initial state. The activation energy of this process is also low (E" which causes it to proceed at a high speed. When summing up both sequential processes, the final equation for a rapidly occurring reaction is obtained: A + B (+ K) = AB (+ K). The catalyst is indicated in this equation only to emphasize the fact of its regeneration.

What all catalysts have in common is that they always change the activation energy, decreasing it during positive catalysis, i.e. reducing the height of the energy barrier. In the presence of a catalyst, an activated complex is formed with a lower energy level than without it, resulting in a significant increase in the reaction rate.

Based on phase characteristics, homogeneous (uniform) and heterogeneous (inhomogeneous) catalysis are distinguished; Enzymatic catalysis is considered separately.

At homogeneous catalysis the catalyst and reactants form one phase (gas or solution), in which there are no interfaces (phase boundaries). Gas and liquid phase catalytic processes are very numerous. An example of homogeneous catalysis in the gas phase is the catalytic oxidation of sulfur (IV) oxide in a chamber method for producing sulfuric acid. Oxidation of sulfur dioxide to trioxide by the reaction:

proceeds slowly. The introduction of NO catalyst changes the reaction mechanism:

and reduces the activation energy, and therefore increases the rate of reaction.

In homogeneous catalysis, the rate of a chemical reaction is proportional to the catalyst concentration. The disadvantages of homogeneous catalysis in solutions are the limited temperature range and, in some cases, the difficulty of separating the catalyst from the reaction products.

At heterogeneous catalysis the catalyst (usually a solid) is present in the system as an independent phase, i.e. There is an interface between the catalyst and the reactants (gases or liquids). Thus, the oxidation of ammonia (gaseous phase) is carried out in the presence of platinum (solid phase), and the decomposition of hydrogen peroxide (liquid phase) is accelerated by coal or manganese (IV) oxide present in the form of a solid phase:

In heterogeneous catalysis, all reactions occur at the phase boundary, i.e. on the surface of the catalyst, the activity of which depends on the properties of its surface - area size, chemical composition, defective structure and state. Features of the kinetics of processes are determined by diffusion and adsorption.

The surface of the catalyst (adsorbent) is physically heterogeneous and has so-called active centers, on which catalytic reactions mainly occur due to the adsorption of reactants (adsorbates) on these centers and an increase in their concentration on the surface of the catalyst. This partly leads to a faster reaction. However, the main reason for the increase in the reaction rate is a significant increase in the chemical activity of adsorbed molecules, in which, under the influence of a catalyst, the bonds between atoms are weakened, which makes these molecules more reactive. The acceleration of the reaction in this case also occurs as a result of a decrease in the activation energy, to which a certain contribution is made by the formation of surface intermediate compounds.

Substances that poison the solid catalyst, i.e. reducing or completely destroying its activity are called catalytic poisons. For example, compounds of arsenic, mercury, lead, and cyanide poison platinum catalysts, which in this case must be regenerated under production conditions.

Substances that enhance the effect of catalysts for a given reaction, but are not catalysts themselves, are called promoters. It is known, for example, to promote platinum catalysts with additives of iron, aluminum, etc.

Selectivity of action The effectiveness of catalysts is manifested, in particular, in the fact that with the help of different catalysts it is possible to obtain different products from the same substance. Thus, in the presence of catalyst A1 2 Oe at 300 °C, water and ethylene are obtained from ethyl alcohol:

But if copper powder is used as a catalyst at the same temperature, then ethyl alcohol decomposes into hydrogen and acetaldehyde:

Thus, each reaction has its own catalyst.

With the participation of biological catalysts, enzymes, Complex chemical processes occur in plant and animal organisms. For example, saliva contains the enzyme ptyalin, which catalyzes the conversion of starch into sugar, and pepsin, present in gastric juice, promotes the breakdown of proteins. There are about 3,000 different enzymes in the human body, each of which is an effective catalyst for a corresponding reaction.

Many catalysts, especially enzymes, have purely individual catalytic action, which is why they are called individually specific.

S. I. LEVCHENKOV

PHYSICAL AND COLLOIDAL CHEMISTRY

Lecture notes for students of the Faculty of Biology of Southern Federal University (RSU)

2.3 CATALYTIC PROCESSES

The rate of a chemical reaction at a given temperature is determined by the rate of formation of the activated complex, which, in turn, depends on the value of the activation energy. In many chemical reactions, the structure of the activated complex may include substances that are not stoichiometrically reagents; It is obvious that in this case the activation energy of the process also changes. In the case of the presence of several transition states, the reaction will proceed mainly along the path with the lowest activation barrier.

Catalysis is the phenomenon of changing the rate of a chemical reaction in the presence of substances, the state and quantity of which remain unchanged after the reaction.

Distinguish positive And negative catalysis (respectively, an increase and decrease in the rate of a reaction), although the term “catalysis” often means only positive catalysis; negative catalysis is called inhibition.

A substance that is part of the structure of the activated complex, but is not stoichiometrically a reagent, is called a catalyst. All catalysts are characterized by such common properties as specificity and selectivity of action.

Specificity A catalyst lies in its ability to accelerate only one reaction or a group of similar reactions and not affect the rate of other reactions. For example, many transition metals (platinum, copper, nickel, iron, etc.) are catalysts for hydrogenation processes; aluminum oxide catalyzes hydration reactions, etc.

Selectivity catalyst - the ability to accelerate one of the parallel reactions possible under given conditions. Thanks to this, it is possible, using different catalysts, to obtain different products from the same starting materials:

The reason for the increase in the reaction rate with positive catalysis is the decrease in activation energy when the reaction proceeds through an activated complex with the participation of a catalyst (Fig. 2.8).

Since, according to the Arrhenius equation, the rate constant of a chemical reaction is exponentially dependent on the activation energy, a decrease in the latter causes a significant increase in the rate constant. Indeed, if we assume that the pre-exponential factors in the Arrhenius equation (II.32) for catalytic and non-catalytic reactions are close, then for the ratio of rate constants we can write:

If ΔE A = –50 kJ/mol, then the ratio of the rate constants will be 2.7 10 6 times (indeed, in practice such a decrease in E A increases the reaction rate by approximately 10 5 times).

It should be noted that the presence of a catalyst does not affect the magnitude of the change in thermodynamic potential as a result of the process and, therefore, no catalyst can make possible the spontaneous occurrence of a thermodynamically impossible process (a process whose ΔG (ΔF) is greater than zero). The catalyst does not change the value of the equilibrium constant for reversible reactions; the influence of the catalyst in this case is only to accelerate the achievement of an equilibrium state.

Depending on the phase state of the reagents and the catalyst, homogeneous and heterogeneous catalysis is distinguished.

Rice. 2.8 Energy diagram of a chemical reaction without a catalyst (1)
and in the presence of a catalyst (2).

2.3.1 Homogeneous catalysis.

Homogeneous catalysis - catalytic reactions in which the reactants and catalyst are in the same phase. In the case of homogeneous catalytic processes, the catalyst forms intermediate reactive products with the reagents. Let's consider some reaction

A + B ––> C

In the presence of a catalyst, two rapidly occurring stages are carried out, as a result of which particles of the intermediate compound AA are formed and then (through the activated ABC complex #) the final reaction product with catalyst regeneration:

A + K ––> AK

AK + B ––> C + K

An example of such a process is the decomposition reaction of acetaldehyde, the activation energy of which is E A = 190 kJ/mol:

CH 3 CHO ––> CH 4 + CO

In the presence of iodine vapor, this process occurs in two stages:

CH 3 CHO + I 2 ––> CH 3 I + HI + CO

CH 3 I + HI ––> CH 4 + I 2

The decrease in activation energy of this reaction in the presence of a catalyst is 54 kJ/mol; the reaction rate constant increases approximately 105 times. The most common type of homogeneous catalysis is acid catalysis, in which hydrogen ions H + act as a catalyst.

2.3.2 Autocatalysis.

Autocatalysis– the process of catalytic acceleration of a chemical reaction by one of its products. An example is the hydrolysis of esters catalyzed by hydrogen ions. The acid formed during hydrolysis dissociates to form protons, which accelerate the hydrolysis reaction. The peculiarity of an autocatalytic reaction is that this reaction proceeds with a constant increase in the concentration of the catalyst. Therefore, in the initial period of the reaction, its speed increases, and at subsequent stages, as a result of a decrease in the concentration of reagents, the speed begins to decrease; the kinetic curve of the product of an autocatalytic reaction has a characteristic S-shaped appearance (Fig. 2.9).

Rice. 2.9 Kinetic curve of the product of an autocatalytic reaction

2.3.3 Heterogeneous catalysis.

Heterogeneous catalysis – catalytic reactions occurring at the interface between the phases formed by the catalyst and reactants. The mechanism of heterogeneous catalytic processes is much more complex than in the case of homogeneous catalysis. In each heterogeneous catalytic reaction, at least six stages can be distinguished:

1. Diffusion of starting substances to the catalyst surface.

2. Adsorption of starting substances on the surface with the formation of some intermediate compound:

A + B + K ––> АВК

3. Activation of the adsorbed state (the energy required for this is the true activation energy of the process):

AVK ––> AVK #

4. Decomposition of the activated complex with the formation of adsorbed reaction products:

АВК # ––> СДК

5. Desorption of reaction products from the catalyst surface.

СDК ––> С + D + К

6. Diffusion of reaction products from the catalyst surface.

A specific feature of heterocatalytic processes is the ability of the catalyst to promote and poison.

Promotion– an increase in the activity of the catalyst in the presence of substances that are not themselves catalysts for this process (promoters). For example, for the nickel metal catalyzed reaction

CO + H 2 ––> CH 4 + H 2 O

the introduction of a small cerium impurity into a nickel catalyst leads to a sharp increase in the activity of the catalyst.

Poisoning– a sharp decrease in catalyst activity in the presence of certain substances (so-called catalytic poisons). For example, for the reaction of ammonia synthesis (the catalyst is sponge iron), the presence of oxygen or sulfur compounds in the reaction mixture causes a sharp decrease in the activity of the iron catalyst; at the same time, the ability of the catalyst to adsorb starting materials decreases very slightly.

To explain these features of heterogeneous catalytic processes, G. Taylor made the following assumption: not the entire surface of the catalyst is catalytically active, but only some of its areas - the so-called. active centers , which can be various defects in the crystal structure of the catalyst (for example, protrusions or depressions on the surface of the catalyst). Currently, there is no unified theory of heterogeneous catalysis. For metal catalysts it was developed multiplet theory . The main provisions of the multiplet theory are as follows:

1. The active center of a catalyst is a set of a certain number of adsorption centers located on the surface of the catalyst in geometrical accordance with the structure of the molecule undergoing the transformation.

2. During the adsorption of reacting molecules on the active center, a multiplet complex is formed, resulting in a redistribution of bonds, leading to the formation of reaction products.

The theory of multiplets is sometimes called the theory of geometric similarity of the active center and reacting molecules. For different reactions, the number of adsorption centers (each of which is identified with a metal atom) in the active center is different - 2, 3, 4, etc. Such active centers are called doublet, triplet, quadruplet, etc., respectively. (in the general case, a multiplet, which is what the theory owes its name to).

For example, according to the theory of multiplets, the dehydrogenation of saturated monohydric alcohols occurs on a doublet, and the dehydrogenation of cyclohexane occurs on a sextet (Fig. 2.10 - 2.11); The theory of multiplets made it possible to relate the catalytic activity of metals to the value of their atomic radius.

Rice. 2.10 Dehydrogenation of alcohols on a doublet

Rice. 2.11 Dehydrogenation of cyclohexane on a sextet

2.3.4 Enzymatic catalysis.

Enzyme catalysis – catalytic reactions occurring with the participation of enzymes – biological catalysts of protein nature. Enzyme catalysis has two characteristic features:

1. High activity , is several orders of magnitude higher than the activity of inorganic catalysts, which is explained by a very significant decrease in the activation energy of the process by enzymes. Thus, the rate constant for the decomposition reaction of hydrogen peroxide catalyzed by Fe 2+ ions is 56 s -1 ; the rate constant of the same reaction catalyzed by the enzyme catalase is 3.5·10 7 , i.e. the reaction in the presence of the enzyme proceeds a million times faster (the activation energies of the processes are 42 and 7.1 kJ/mol, respectively). The rate constants for urea hydrolysis in the presence of acid and urease differ by thirteen orders of magnitude, amounting to 7.4·10 -7 and 5·10 6 s -1 (the activation energy is 103 and 28 kJ/mol, respectively).

2. High specificity . For example, amylase catalyzes the breakdown of starch, which is a chain of identical glucose units, but does not catalyze the hydrolysis of sucrose, the molecule of which is composed of glucose and fructose fragments.

According to generally accepted ideas about the mechanism of enzymatic catalysis, the substrate S and enzyme F are in equilibrium with the very quickly formed enzyme-substrate complex FS, which relatively slowly decomposes into the reaction product P with the release of free enzyme; Thus, the stage of decomposition of the enzyme-substrate complex into reaction products is rate-determining (limiting).

F+S<––>FS ––> F + P

A study of the dependence of the rate of an enzymatic reaction on the concentration of the substrate at a constant concentration of the enzyme showed that with increasing concentration of the substrate, the reaction rate first increases and then stops changing (Fig. 2.12) and the dependence of the reaction rate on the concentration of the substrate is described by the following equation:

(II.45)