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:
|
: CO + H 2 ––> CH 3 OH |
: C 2 H 5 OH ––> C 2 H 4 + H 2 O |
|
: CO + H 2 ––> CH 4 + H 2 O |
: C 2 H 5 OH ––> CH 3 CHO + H 2 |
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 ceases to change (Fig. 2.12) and the dependence of the reaction rate on the concentration of the substrate is described by the following equation:
(II.45)
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 = kvpP n 0 = k 0 e -Ea/RT vpP n 0 (4)
where v is the gas flow rate, p is the driving force of the process at P0.1 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, when oxidizing 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, finely 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, PaC; 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
acceleration of chemical reactions under the influence of small amounts of substances (catalysts), which themselves do not change during the reaction. Catalytic processes play a huge role in our lives. Biological catalysts, called enzymes, are involved in the regulation of biochemical processes. Without catalysts, many industrial processes could not take place. The most important property of catalysts is selectivity, i.e. the ability to increase the rate of only certain chemical reactions out of many possible ones. This allows reactions that are too slow to be practical under normal conditions and ensures the formation of the desired products. The use of catalysts contributed to the rapid development of the chemical industry. They are widely used in oil refining, obtaining various products, and creating new materials (for example, plastics), often cheaper than those used before. Approximately 90% of modern chemical production is based on catalytic processes. Catalytic processes play a special role in environmental protection. In 1835, the Swedish chemist J. Berzelius found that in the presence of certain substances, the rate of some chemical reactions increases significantly. For such substances, he introduced the term “catalyst” (from the Greek katalysis - relaxation). According to Berzelius, catalysts have a special ability to weaken the bonds between atoms in the molecules involved in the reaction, thus facilitating their interaction. A great contribution to the development of ideas about the work of catalysts was made by the German physical chemist W. Ostwald, who in 1880 defined a catalyst as a substance that changes the rate of a reaction. According to modern concepts, the catalyst forms a complex with reacting molecules, stabilized by chemical bonds. After rearrangement, this complex dissociates, releasing the products and catalyst. For a monomolecular reaction of converting a molecule X into Y, this entire process can be represented as X + Cat. -> X-Cat. -> Y-Cat. -> Y + Cat. The released catalyst again binds to X, and the entire cycle is repeated many times, ensuring the formation of large quantities of the product - substance Y. Many substances under normal conditions do not react chemically with each other. Thus, hydrogen and carbon monoxide at room temperature do not interact with each other, since the bond between the atoms in the H2 molecule is strong enough and is not broken when attacked by a CO molecule. The catalyst brings H2 and CO molecules together, forming bonds with them. After rearrangement, the catalyst-reagent complex dissociates to form a product containing C, H and O atoms. Often, when the same substances interact, different products are formed. A catalyst can direct a process along the path most favorable for the formation of a particular product. Consider the reaction between CO and H2. In the presence of a copper-containing catalyst, practically the only reaction product is methanol:
First, CO and H2 molecules are adsorbed on the surface of the catalyst. Then CO molecules form chemical bonds with the catalyst (chemisorption occurs), remaining in an undissociated form. Hydrogen molecules are also chemisorbed on the surface of the catalyst, but at the same time dissociate. As a result of the rearrangement, a transition complex H-Cat.-CH2OH is formed. After the addition of an H atom, the complex decomposes to release CH3OH and the catalyst. In the presence of a nickel catalyst, both CO and H2 are chemisorbed on the surface in dissociated form, and the Cat-CH3 complex is formed. The final products of the reaction are CH4 and H2O:
Most catalytic reactions are carried out at a certain pressure and temperature by passing the reaction mixture, which is in a gaseous or liquid state, through a reactor filled with catalyst particles. The following concepts are used to describe reaction conditions and product characteristics. Space velocity is the volume of gas or liquid passing through a unit volume of catalyst per unit time. Catalytic activity is the amount of reactants converted by a catalyst into products per unit time. Conversion is the fraction of a substance converted in a given reaction. Selectivity is the ratio of the amount of a particular product to the total amount of products (usually expressed as a percentage). Yield is the ratio of the amount of a given product to the amount of starting material (usually expressed as a percentage). Productivity is the number of reaction products formed per unit volume per unit time.
TYPES OF CATALYST
Catalysts are classified based on the nature of the reaction they accelerate, their chemical composition, or their physical properties. Almost all chemical elements and substances have catalytic properties to one degree or another - on their own or, more often, in various combinations. Based on their physical properties, catalysts are divided into homogeneous and heterogeneous. Heterogeneous catalysts are solid substances that are homogeneous dispersed in the same gas or liquid medium as the reacting substances. Many heterogeneous catalysts contain metals. Some metals, especially those belonging to group VIII of the periodic table of elements, have catalytic activity on their own; a typical example is platinum. But most metals exhibit catalytic properties when present in compounds; an example is alumina (aluminum oxide Al2O3). An unusual property of many heterogeneous catalysts is their large surface area. They are penetrated by numerous pores, the total area of which sometimes reaches 500 m2 per 1 g of catalyst. In many cases, oxides with a large surface area serve as a substrate on which metal catalyst particles are deposited in the form of small clusters. This ensures effective interaction of reagents in the gas or liquid phase with the catalytically active metal. A special class of heterogeneous catalysts are zeolites - crystalline minerals of the group of aluminosilicates (compounds of silicon and aluminum). Although many heterogeneous catalysts have a large surface area, they usually have only a small number of active sites, which account for a small portion of the total surface area. Catalysts may lose their activity in the presence of small amounts of chemical compounds called catalyst poisons. These substances bind to active centers, blocking them. Determining the structure of active sites is the subject of intensive research. Homogeneous catalysts have a different chemical nature - acids (H2SO4 or H3PO4), bases (NaOH), organic amines, metals, most often transition metals (Fe or Rh), in the form of salts, organometallic compounds or carbonyls. Catalysts also include enzymes - protein molecules that regulate biochemical reactions. The active site of some enzymes contains a metal atom (Zn, Cu, Fe or Mo). Metal-containing enzymes catalyze reactions involving small molecules (O2, CO2 or N2). Enzymes have very high activity and selectivity, but they work only under certain conditions, such as those under which reactions occur in living organisms. In industry, the so-called is often used. immobilized enzymes.
HOW CATALYSTS WORK
Energy. Any chemical reaction can occur only if the reactants overcome the energy barrier, and for this they must acquire a certain energy. As we have already said, the catalytic reaction X (r) Y consists of a number of successive stages. For each of them to occur, energy E is required, called activation energy. The change in energy along the reaction coordinate is shown in Fig. 1.
Rice. 1. CHANGE IN THE ENERGY OF REAGENTS during the catalytic and “thermal” pathways of the reaction.
Let us first consider the non-catalytic, “thermal” path. For the reaction to take place, the potential energy of the X molecules must exceed the energy barrier Et. The catalytic reaction consists of three stages. The first is the formation of the X-Cat complex. (chemisorption), the activation energy of which is Eads. The second stage is the regrouping of X-Cat. (r) Y-Cat. with activation energy Ekat, and finally, the third - desorption with activation energy Edes; Eads, Ekat and Edes are much smaller than Et. Since the reaction rate depends exponentially on the activation energy, the catalytic reaction proceeds much faster than the thermal reaction at a given temperature. A catalyst can be likened to a guide who guides climbers (reacting molecules) across a mountain range. He leads one group through the pass and then returns for the next. The path through the pass lies significantly lower than that through the peak (thermal channel of the reaction), and the group makes the transition faster than without a conductor (catalyst). It is even possible that the group would not have been able to overcome the ridge on its own.
Theories of catalysis. To explain the mechanism of catalytic reactions, three groups of theories have been proposed: geometric, electronic and chemical. In geometric theories, the main attention is paid to the correspondence between the geometric configuration of the atoms of the active centers of the catalyst and the atoms of that part of the reacting molecules that is responsible for binding to the catalyst. Electronic theories are based on the idea that chemisorption is caused by electronic interaction associated with charge transfer, i.e. these theories relate catalytic activity to the electronic properties of the catalyst. Chemical theory views a catalyst as a chemical compound with characteristic properties that forms chemical bonds with reagents, resulting in the formation of an unstable transition complex. After the decomposition of the complex with the release of products, the catalyst returns to its original state. The latter theory is now considered the most adequate. At the molecular level, a catalytic gas-phase reaction can be represented as follows. One reacting molecule binds to the active site of the catalyst, and the other interacts with it, being directly in the gas phase. An alternative mechanism is also possible: the reacting molecules are adsorbed on neighboring active centers of the catalyst and then interact with each other. Apparently, this is how most catalytic reactions proceed. Another concept suggests that there is a relationship between the spatial arrangement of atoms on the surface of a catalyst and its catalytic activity. The rate of some catalytic processes, including many hydrogenation reactions, does not depend on the relative position of the catalytically active atoms on the surface; the speed of others, on the contrary, changes significantly with changes in the spatial configuration of surface atoms. An example is the isomerization of neopentane into isopentane and the simultaneous cracking of the latter to isobutane and methane on the surface of a Pt-Al2O3 catalyst.
APPLICATION OF CATALYSIS IN INDUSTRY
The rapid industrial growth that we are now experiencing would not have been possible without the development of new chemical technologies. To a large extent, this progress is determined by the widespread use of catalysts, with the help of which low-grade raw materials are converted into high-value products. Figuratively speaking, a catalyst is the philosopher’s stone of a modern alchemist, only it transforms not lead into gold, but raw materials into medicines, plastics, chemicals, fuel, fertilizers and other useful products. Perhaps the very first catalytic process that man learned to use was fermentation. Recipes for preparing alcoholic beverages were known to the Sumerians as early as 3500 BC.
See WINE; BEER . A significant milestone in the practical application of catalysis was the production of margarine by the catalytic hydrogenation of vegetable oil. This reaction was first carried out on an industrial scale around 1900. And since the 1920s, catalytic methods for producing new organic materials, primarily plastics, have been developed one after another. The key point was the catalytic production of olefins, nitriles, esters, acids, etc. - “bricks” for the chemical “construction” of plastics. The third wave of industrial use of catalytic processes occurred in the 1930s and was associated with petroleum refining. In terms of volume, this production soon left all others far behind. Petroleum refining consists of several catalytic processes: cracking, reforming, hydrosulfonation, hydrocracking, isomerization, polymerization and alkylation. Finally, the fourth wave in the use of catalysis is related to environmental protection. The most famous achievement in this area is the creation of a catalytic converter for automobile exhaust gases. Catalytic converters, which have been installed in cars since 1975, have played a big role in improving air quality and thus saving many lives. About a dozen Nobel Prizes have been awarded for work in catalysis and related fields. The practical importance of catalytic processes is evidenced by the fact that the share of nitrogen included in industrially produced nitrogen-containing compounds accounts for about half of all nitrogen included in food products. The amount of nitrogen compounds produced naturally is limited, so the production of dietary protein depends on the amount of nitrogen added to the soil through fertilizers. It would be impossible to feed half of humanity without synthetic ammonia, which is produced almost exclusively through the Haber-Bosch catalytic process. The scope of application of catalysts is constantly expanding. It is also important that catalysis can significantly increase the efficiency of previously developed technologies. An example is the improvement in catalytic cracking through the use of zeolites.
Hydrogenation. A large number of catalytic reactions are associated with the activation of a hydrogen atom and some other molecule, leading to their chemical interaction. This process is called hydrogenation and underlies many stages of oil refining and the production of liquid fuels from coal (Bergius process). The production of aviation gasoline and motor fuel from coal was developed in Germany during World War II because the country had no oil fields. The Bergius process involves the direct addition of hydrogen to coal. Coal is heated under pressure in the presence of hydrogen to produce a liquid product, which is then processed into aviation gasoline and motor fuel. Iron oxide is used as a catalyst, as well as catalysts based on tin and molybdenum. During the war, 12 factories in Germany produced approximately 1,400 tons of liquid fuel per day using the Bergius process. Another process, Fischer-Tropsch, consists of two stages. First, the coal is gasified, i.e. They react it with water vapor and oxygen and obtain a mixture of hydrogen and carbon oxides. This mixture is converted into liquid fuel using catalysts containing iron or cobalt. With the end of the war, the production of synthetic fuel from coal in Germany was discontinued. As a result of the rise in oil prices that followed the oil embargo of 1973–1974, vigorous efforts were made to develop a cost-effective method of producing gasoline from coal. Thus, direct liquefaction of coal can be carried out more efficiently using a two-stage process in which the coal is first contacted with an aluminum-cobalt-molybdenum catalyst at a relatively low temperature and then at a higher temperature. The cost of such synthetic gasoline is higher than that obtained from oil.
Ammonia. One of the simplest hydrogenation processes from a chemical point of view is the synthesis of ammonia from hydrogen and nitrogen. Nitrogen is a very inert substance. To break the N-N bond in its molecule, an energy of about 200 kcal/mol is required. However, nitrogen binds to the surface of the iron catalyst in the atomic state, and this requires only 20 kcal/mol. Hydrogen binds to iron even more readily. Ammonia synthesis proceeds as follows:
This example illustrates the ability of a catalyst to accelerate both forward and reverse reactions equally, i.e. the fact that the catalyst does not change the equilibrium position of a chemical reaction. Hydrogenation of vegetable oil. One of the most important hydrogenation reactions in practical terms is the incomplete hydrogenation of vegetable oils to margarine, cooking oil and other food products. Vegetable oils are obtained from soybeans, cotton seeds and other crops. They contain esters, namely triglycerides of fatty acids with varying degrees of unsaturation. Oleic acid CH3(CH2)7CH=CH(CH2)7COOH has one C=C double bond, linoleic acid has two and linolenic acid has three. The addition of hydrogen to break this bond prevents oils from oxidizing (rancidity). This increases their melting point. The hardness of most resulting products depends on the degree of hydrogenation. Hydrogenation is carried out in the presence of fine nickel powder deposited on a substrate or a Raney nickel catalyst in an atmosphere of highly purified hydrogen.
Dehydrogenation. Dehydrogenation is also an industrially important catalytic reaction, although the scale of its application is incomparably smaller. With its help, for example, styrene, an important monomer, is obtained. To do this, ethylbenzene is dehydrogenated in the presence of a catalyst containing iron oxide; The reaction is also facilitated by potassium and some kind of structural stabilizer. The dehydrogenation of propane, butane and other alkanes is carried out on an industrial scale. Dehydrogenation of butane in the presence of a chromium-alumina catalyst produces butenes and butadiene.
Acid catalysis. The catalytic activity of a large class of catalysts is determined by their acidic properties. According to I.
Bronsted and T. Lowry, an acid is a compound capable of donating a proton. Strong acids easily donate their protons to bases. The concept of acidity was further developed in the works of G. Lewis, who defined acid as a substance capable of accepting an electron pair from a donor substance with the formation of a covalent bond due to the socialization of this electron pair. These ideas, together with ideas about reactions that produce carbenium ions, helped to understand the mechanism of a variety of catalytic reactions, especially those involving hydrocarbons. The strength of an acid can be determined by using a set of bases that change color when a proton is added. It turns out that some industrially important catalysts behave like very strong acids. These include a Friedel-Crafts process catalyst, such as HCl-AlCl2O3 (or HAlCl4), and aluminosilicates. Acid strength is a very important characteristic because it determines the rate of protonation, a key step in the acid catalysis process. The activity of catalysts such as aluminosilicates, used in oil cracking, is determined by the presence of Bronsted and Lewis acids on their surface. Their structure is similar to the structure of silica (silicon dioxide), in which some of the Si4+ atoms are replaced by Al3+ atoms. The excess negative charge that arises in this case can be neutralized by the corresponding cations. If the cations are protons, then the aluminosilicate behaves like a Bronsted acid:
The activity of acid catalysts is determined by their ability to react with hydrocarbons to form a carbenium ion as an intermediate product. Alkylcarbenium ions contain a positively charged carbon atom bonded to three alkyl groups and/or hydrogen atoms. They play an important role as intermediates formed in many reactions involving organic compounds. The mechanism of action of acid catalysts can be illustrated by the example of the isomerization reaction of n-butane to isobutane in the presence of HCl-AlCl3 or Pt-Cl-Al2O3. First, a small amount of C4H8 olefin attaches to the positively charged hydrogen ion of the acid catalyst to form a tertiary carbenium ion. The negatively charged hydride ion H- is then split off from n-butane to form isobutane and a secondary butylcarbenium ion. The latter, as a result of rearrangement, turns into a tertiary carbenium ion. This chain can continue with the elimination of a hydride ion from the next n-butane molecule, etc.:
It is significant that tertiary carbenium ions are more stable than primary or secondary ones. As a result, they are mainly present on the surface of the catalyst, and therefore the main product of butane isomerization is isobutane. Acid catalysts are widely used in oil refining - cracking, alkylation, polymerization and isomerization of hydrocarbons
(see also CHEMISTRY AND METHODS OF OIL PROCESSING).
The mechanism of action of carbenium ions, which play the role of catalysts in these processes, has been established. In doing so, they participate in a number of reactions, including the formation of small molecules by cleavage of large molecules, the combination of molecules (olefin to olefin or olefin to isoparaffin), structural rearrangement by isomerization, and the formation of paraffins and aromatic hydrocarbons by hydrogen transfer. One of the latest applications of acid catalysis in industry is the production of leaded fuels by adding alcohols to isobutylene or isoamylene. Adding oxygen-containing compounds to gasoline reduces the concentration of carbon monoxide in exhaust gases. Methyl tert-butyl ether (MTBE) with an octane mixing number of 109 also makes it possible to obtain the high-octane fuel necessary for running a high-compression automobile engine without introducing tetraethyl lead into gasoline. The production of fuels with octane numbers 102 and 111 has also been organized.
Basic catalysis. The activity of catalysts is determined by their basic properties. A long-standing and well-known example of such catalysts is sodium hydroxide, used to hydrolyze or saponify fats to make soap, and one recent example is catalysts used in the production of polyurethane plastics and foams. Urethane is formed by the reaction of alcohol with isocyanate, and this reaction is accelerated in the presence of basic amines. During the reaction, a base attaches to the carbon atom in the isocyanate molecule, as a result of which a negative charge appears on the nitrogen atom and its activity towards alcohol increases. Triethylenediamine is a particularly effective catalyst. Polyurethane plastics are produced by reacting diisocyanates with polyols (polyalcohols). When isocyanate reacts with water, the previously formed urethane decomposes, releasing CO2. When a mixture of polyalcohols and water interacts with diisocyanates, the resulting polyurethane foam foams with CO2 gas.
Double acting catalysts. These catalysts speed up two types of reactions and produce better results than passing the reactants in series through two reactors, each containing only one type of catalyst. This is due to the fact that the active sites of a double-acting catalyst are very close to each other, and the intermediate product formed at one of them is immediately converted into the final product at the other. A good result is obtained by combining a catalyst that activates hydrogen with a catalyst that promotes the isomerization of hydrocarbons. The activation of hydrogen is carried out by some metals, and the isomerization of hydrocarbons is carried out by acids. An effective dual-acting catalyst used in petroleum refining to convert naphtha into gasoline is finely divided platinum supported on acidic alumina. Converting naphtha constituents such as methylcyclopentane (MCP) to benzene increases the octane number of gasoline. First, MCP is dehydrogenated on the platinum part of the catalyst into an olefin with the same carbon skeleton; the olefin then passes to the acid portion of the catalyst, where it isomerizes to cyclohexene. The latter passes to the platinum part and is dehydrogenated to benzene and hydrogen. Double-action catalysts significantly accelerate oil reforming. They are used to isomerize normal paraffins into isoparaffins. The latter, boiling at the same temperatures as gasoline fractions, are valuable because they have a higher octane number compared to straight hydrocarbons. In addition, the conversion of n-butane to isobutane is accompanied by dehydrogenation, facilitating the production of MTBE.
Stereospecific polymerization. An important milestone in the history of catalysis was the discovery of the catalytic polymerization of a-olefins to form stereoregular polymers. Stereospecific polymerization catalysts were discovered by K. Ziegler when he was trying to explain the unusual properties of the polymers he obtained. Another chemist, J. Natta, suggested that the uniqueness of Ziegler polymers is determined by their stereoregularity. X-ray diffraction experiments have shown that polymers prepared from propylene in the presence of Ziegler catalysts are highly crystalline and indeed have a stereoregular structure. To describe such ordered structures, Natta introduced the terms “isotactic” and “syndiotactic”. In the case where there is no order, the term “atactic” is used:
A stereospecific reaction occurs on the surface of solid catalysts containing transition metals of groups IVA-VIII (such as Ti, V, Cr, Zr), which are in a partially oxidized state, and any compound containing carbon or hydrogen, which is bonded to the metal from groups I-III. A classic example of such a catalyst is the precipitate formed by the interaction of TiCl4 and Al(C2H5)3 in heptane, where titanium is reduced to the trivalent state. This exceptionally active system catalyzes the polymerization of propylene at normal temperatures and pressures.
Catalytic oxidation. The use of catalysts to control the chemistry of oxidation processes is of great scientific and practical importance. In some cases, oxidation must be complete, for example when neutralizing CO and hydrocarbon contaminants in automobile exhaust gases. However, more often it is necessary for the oxidation to be incomplete, for example, in many widely used industrial processes for converting hydrocarbons into valuable intermediate products containing functional groups such as -CHO, -COOH, -C-CO, -CN. In this case, both homogeneous and heterogeneous catalysts are used. An example of a homogeneous catalyst is a transition metal complex, which is used to oxidize para-xylene to terephthalic acid, the esters of which form the basis for the production of polyester fibers.
Catalysts for heterogeneous oxidation. These catalysts are usually complex solid oxides. Catalytic oxidation occurs in two stages. First, the oxygen in the oxide is captured by a hydrocarbon molecule adsorbed on the surface of the oxide. In this case, the hydrocarbon is oxidized, and the oxide is reduced. The reduced oxide reacts with oxygen and returns to its original state. Using a vanadium catalyst, phthalic anhydride is obtained by incomplete oxidation of naphthalene or butane.
Production of ethylene by dehydrodimerization of methane. Ethylene synthesis through dehydrodimerization converts natural gas into more easily transportable hydrocarbons. The reaction 2CH4 + 2O2 -> C2H4 + 2H2O is carried out at 850° C using various catalysts; the best results were obtained with the Li-MgO catalyst. Presumably the reaction proceeds through the formation of a methyl radical by the abstraction of a hydrogen atom from a methane molecule. The elimination is carried out by incompletely reduced oxygen, for example O22-. Methyl radicals in the gas phase recombine to form an ethane molecule and, during subsequent dehydrogenation, are converted to ethylene. Another example of incomplete oxidation is the conversion of methanol to formaldehyde in the presence of a silver or iron-molybdenum catalyst.
Zeolites. Zeolites constitute a special class of heterogeneous catalysts. These are aluminosilicates with an ordered honeycomb structure, the cell size of which is comparable to the size of many organic molecules. They are also called molecular sieves. Of greatest interest are zeolites, the pores of which are formed by rings consisting of 8-12 oxygen ions (Fig. 2). Sometimes the pores overlap, as in the ZSM-5 zeolite (Fig. 3), which is used for the highly specific conversion of methanol into gasoline fraction hydrocarbons. Gasoline contains significant amounts of aromatic hydrocarbons and therefore has a high octane number. In New Zealand, for example, a third of all gasoline consumed is produced using this technology. Methanol is produced from imported methane.
The catalysts that make up the group of Y-zeolites significantly increase the efficiency of catalytic cracking due primarily to their unusual acidic properties. Replacing aluminosilicates with zeolites makes it possible to increase gasoline yield by more than 20%. In addition, zeolites have selectivity regarding the size of the reacting molecules. Their selectivity is determined by the size of the pores through which molecules of only certain sizes and shapes can pass. This applies to both starting materials and reaction products. For example, due to steric restrictions, para-xylene is formed more easily than the bulkier ortho and meta isomers. The latter find themselves “locked” in the pores of the zeolite (Fig. 4).
The use of zeolites has made a real revolution in some industrial technologies - dewaxing of gas oil and engine oil, obtaining chemical intermediates for the production of plastics by alkylation of aromatic compounds, isomerization of xylene, disproportionation of toluene and catalytic cracking of oil. ZSM-5 zeolite is especially effective here.
Catalysts and environmental protection. The use of catalysts to reduce air pollution began in the late 1940s. In 1952, A. Hagen-Smith found that hydrocarbons and nitrogen oxides contained in exhaust gases react in light to form oxidants (in particular, ozone), which irritate the eyes and give other undesirable effects. Around the same time, Y. Khoudri developed a method for catalytic exhaust gas purification by oxidizing CO and hydrocarbons to CO2 and H2O. In 1970, the Clean Air Declaration (refined in 1977, expanded in 1990) was formulated, according to which all new cars, starting with 1975 models, must be equipped with catalytic converters. Standards for the composition of exhaust gases were established. Because lead compounds added to gasoline poison catalysts, a phase-out program has been adopted. Attention was also drawn to the need to reduce the content of nitrogen oxides. Catalysts have been created specifically for automobile neutralizers, in which active components are applied to a ceramic substrate with a honeycomb structure, through the cells of which exhaust gases pass. The substrate is coated with a thin layer of metal oxide, for example Al2O3, onto which a catalyst - platinum, palladium or rhodium - is applied. The content of nitrogen oxides formed during the combustion of natural fuels in thermal power plants can be reduced by adding small amounts of ammonia to the flue gases and passing them through a titanium vanadium catalyst.
Enzymes. Enzymes are natural catalysts that regulate biochemical processes in a living cell. They participate in energy exchange processes, breakdown of nutrients, and biosynthesis reactions. Without them, many complex organic reactions cannot occur. Enzymes function at ordinary temperatures and pressures, have very high selectivity, and are capable of increasing reaction rates by eight orders of magnitude. Despite these advantages, only approx. 20 of the 15,000 known enzymes are used on a large scale. Man has used enzymes for thousands of years to bake bread, produce alcoholic beverages, cheese and vinegar. Now enzymes are also used in industry: in the processing of sugar, in the production of synthetic antibiotics, amino acids and proteins. Proteolytic enzymes that accelerate hydrolysis processes are added to detergents. With the help of the bacteria Clostridium acetobutylicum, H. Weizmann carried out the enzymatic conversion of starch into acetone and butyl alcohol. This method of producing acetone was widely used in England during the First World War, and during the Second World War it was used to produce butadiene rubber in the USSR. An extremely important role was played by the use of enzymes produced by microorganisms for the synthesis of penicillin, as well as streptomycin and vitamin B12. Ethyl alcohol, produced by enzymatic processes, is widely used as automobile fuel. In Brazil, more than a third of about 10 million cars run on 96% ethyl alcohol derived from sugar cane, while the rest run on a mixture of gasoline and ethyl alcohol (20%). The technology for producing fuel, which is a mixture of gasoline and alcohol, has been well developed in the United States. In 1987, approx. 4 billion liters of alcohol, of which approximately 3.2 billion liters were used as fuel. The so-called also find various applications. immobilized enzymes. These enzymes are bound to a solid support, such as silica gel, over which the reagents are passed. The advantage of this method is that it ensures efficient contact of substrates with the enzyme, separation of products and preservation of the enzyme. One example of the industrial use of immobilized enzymes is the isomerization of D-glucose to fructose.
TECHNOLOGICAL ASPECTS
Modern technologies cannot be imagined without the use of catalysts. Catalytic reactions can occur at temperatures up to 650° C and pressures of 100 atm or more. This forces new solutions to problems associated with contact between gaseous and solid substances and with the transfer of catalyst particles. For the process to be effective, its modeling must take into account kinetic, thermodynamic and hydrodynamic aspects. Computer modeling is widely used here, as well as new instruments and methods for monitoring technological processes. In 1960, significant progress was made in ammonia production. The use of a more active catalyst made it possible to lower the temperature of hydrogen production during the decomposition of water vapor, which made it possible to lower the pressure and, therefore, reduce production costs, for example, through the use of cheaper centrifugal compressors. As a result, the cost of ammonia fell by more than half, there was a colossal increase in its production, and in connection with this, an increase in food production, since ammonia is a valuable fertilizer.
Methods. Research in the field of catalysis is carried out using both traditional and special methods. Radioactive tracers, X-ray, infrared and Raman (Raman) spectroscopy, electron microscopic methods are used; Kinetic measurements are carried out, the influence of methods for preparing catalysts on their activity is studied. Of great importance is the determination of the surface area of the catalyst using the Brunauer-Emmett-Teller method (BET method), based on measuring the physical adsorption of nitrogen at different pressures. To do this, determine the amount of nitrogen required to form a monolayer on the surface of the catalyst, and, knowing the diameter of the N2 molecule, calculate the total area. In addition to determining the total surface area, chemisorption of different molecules is carried out, which makes it possible to estimate the number of active centers and obtain information about their properties. Researchers have various methods at their disposal to study the surface structure of catalysts at the atomic level. The EXAFS method allows you to obtain unique information. Among spectroscopic methods, UV, X-ray and Auger photoelectron spectroscopy are increasingly used. Secondary ion mass spectrometry and ion scattering spectroscopy are of great interest. NMR measurements are used to study the nature of catalytic complexes. A scanning tunneling microscope allows you to see the arrangement of atoms on the surface of the catalyst.
PROSPECTS
The scale of catalytic processes in industry is increasing every year. Catalysts are increasingly used to neutralize substances that pollute the environment. The role of catalysts in the production of hydrocarbons and oxygen-containing synthetic fuels from gas and coal is increasing. The creation of fuel cells for the economical conversion of fuel energy into electrical energy seems very promising. New concepts of catalysis will make it possible to obtain polymeric materials and other products with many valuable properties, improve methods of obtaining energy, and increase food production, in particular by synthesizing proteins from alkanes and ammonia with the help of microorganisms. It may be possible to develop genetically engineered methods for producing enzymes and organometallic compounds that approach natural biological catalysts in their catalytic activity and selectivity.
LITERATURE
Gates B.K. Chemistry of catalytic processes. M., 1981 Boreskov G.K. Catalysis. Questions of theory and practice. Novosibirsk, 1987 Gankin V.Yu., Gankin Yu.V. New general theory of catalysis. L., 1991 Tokabe K. Catalysts and catalytic processes. M., 1993
Collier's Encyclopedia. - Open Society. 2000 .
Catalysis is a change in the rate of chemical reactions under the influence of substances - catalysts, which participate in the process, entering into intermediate chemical interactions with reagents, but remain chemically unchanged after the end of the catalytic act.
Catalysts can be substances in any of three states of aggregation - gases, liquids and solids.
Catalytic processes can be divided into two groups: homogeneous and heterogeneous. If a reaction accelerates in the presence of a catalyst, this phenomenon is called positive catalysis or simply catalysis. If the reaction slows down - anticatalysts or inhibitors.
The essence of catalysis is the same for all its types - homogeneous, heterogeneous, but each of these types has its own distinctive features. In general, the accelerating effect of catalysts is fundamentally different from the action of other factors that intensify chemical processes - temperature, pressure. Thus, as the temperature rises, the temperature of the reacting molecules increases due to heat introduced from outside.
When a catalyst is added, the energy level of the reacting molecules does not change. The action of a catalyst does not shift the equilibrium of a simple reaction, but only accelerates the achievement of equilibrium at a given temperature.
For processes occurring in the kinetic region, the reaction rate
Since ΔС does not change for catalytic and non-catalytic reactions, the effect of the catalyst is to increase the reaction rate constant.
The most common theory that serves as the basis for modern ideas about catalysis is the theory of intermediate compounds. According to this theory, the slow stage between the starting materials can be replaced by two or more faster stages involving a catalyst, which forms weak compounds with the starting materials. The reaction rate is greater, the lower the activation energy due to the exponential dependence
Change in the energy of the reacting system during non-catalytic (1) and catalytic (2) reactions
Let us consider the energy picture of the reaction system, for example for a bimolecular reaction
,
passing in the absence of a catalyst according to the scheme
through the formation of the active complex AB *. In the presence of a catalyst, the reaction follows a different path through several elementary stages:
A + = A
A + B = AB *
AB * = R + . . .
E – activation energy of non-catalytic reaction;
E cat – catalytic reaction;
e 1 and e 2 – activation energies of intermediate stages.
Catalyst activity
The most important characteristic of catalysts is their activity, since it is a measure of the accelerating effect of the catalyst in relation to a given reaction
Let us consider the example of the oxidation of sulfur dioxide
2SO 2 + O 2 = 2SO 3 + Q
The activation energy at 420 o C (693 K) is 420,000 J/mol. On a vanadium catalyst V 2 O 5 E k = 268 kJ/mol K. R = 8.3 J/mol K.
Selectivity (catalyst selectivity)
Selectivity of action is the most important feature of catalysts, which determined the success of their widespread use in a number of industries. It is especially important in the production of organic products, when selectivity allows one to greatly accelerate one useful reaction, carry out the process at a reduced temperature, suppressing other reactions.
The selectivity of the action of catalyst I cat can be expressed by the ratio of the rate of formation of the target product to the total rate of conversion of the main starting material.
,
where G p is the amount of product;
υ p υ out – the ratio of stoichiometric coefficients in the formation of products from the main starting material.
The overall integral selectivity of the action of catalysts can be expressed by the relation
.
where G is the total amount of the starting substance, mol;
Grem - the amount of the starting substance that entered into side reactions;
G p is the amount of the starting substance converted into the target product.
Selectivity is especially pronounced in complex organic reactions. For example, ethyl alcohol, depending on the type of catalyst, can be converted into ethylene
Consequently, it is possible to obtain different target products from the same raw material.
Introduction.
CATALYSIS is a process that involves changing the rate of chemical reactions in the presence of substances called catalysts.
Catalysts are substances that change the rate of a chemical reaction, which can participate in a reaction, be part of intermediate products, but are not part of the final reaction products and remain unchanged after the end of the reaction.
Catalytic reactions are reactions that occur in the presence of catalysts.
Catalysis is called positive, in which the rate of the reaction increases, and negative (inhibition), in which it decreases. An example of positive catalysis is the oxidation of ammonia on platinum to produce nitric acid. An example of a negative is the reduction in corrosion rate when sodium nitrite, potassium chromate and dichromate are introduced into the liquid in which the metal is used.
Catalysts that slow down a chemical reaction are called inhibitors.
Depending on whether the catalyst is in the same phase as the reactants or forms an independent phase, we speak of homogeneous or heterogeneous 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 O + I = H O + IO
H O + IO = H O + O + 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.
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 reacting substances 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 to SO over a VO catalyst in the production of sulfuric acid (contact method).
Promoters (or activators) are substances that increase the activity of the catalyst. In this case, promoters themselves may not have catalytic properties.
Catalytic poisons are foreign impurities in the reaction mixture, leading to partial or complete loss of catalyst activity. Thus, traces of arsenic and phosphorus cause a rapid loss of VO activity by the catalyst (contact method for the production of HSO).
Many important chemical productions, such as the production of sulfuric acid, ammonia, nitric acid, synthetic rubber, a number of polymers, etc., are carried out in the presence of catalysts.
Biochemical reactions in plant and animal organisms are accelerated by biochemical catalysts - enzymes.
Process speed is an extremely important factor determining the productivity of chemical production equipment. Therefore, one of the main tasks set for chemistry by the scientific and technological revolution is the search for ways to increase the rate of reactions. Another important task of modern chemistry, due to the sharply increasing scale of production of chemical products, is increasing the selectivity of chemical transformations into useful products, reducing the amount of emissions and waste. This is also related to environmental protection and more rational use of unfortunately depleting natural resources.
To achieve all these goals, the right means are needed, and such means are primarily catalysts. However, finding them is not so easy. In the process of understanding the internal structure of the things around us, scientists have established a certain gradation, a hierarchy of levels of the microworld. The world described in our book is the world of molecules, the mutual transformations of which constitute the subject of chemistry. We will not be interested in all of chemistry, but only in part of it, devoted to the study of the dynamics of changes in the chemical structure of molecules. Apparently there is no need to say that molecules are built from atoms, and the latter are made from a nucleus and the electron shell surrounding it; that the properties of molecules depend on the nature of their constituent atoms and the sequence of their connection with each other; that the chemical and physical properties of substances depend on the properties of molecules and the nature of their interconnection. We will assume that all this is generally known to the reader, and therefore the main emphasis will be on issues related to the idea of \u200b\u200bthe rate of chemical reactions.
Mutual transformations of molecules occur at very different rates. The rate can be changed by heating or cooling the mixture of reacting molecules. When heated, the reaction rate usually increases, but this is not the only means of accelerating chemical transformations. There is another, more effective method - catalytic, widely used in our time in the production of a wide variety of products.
The first scientific ideas about catalysis arose simultaneously with the development of the atomic theory of the structure of matter. In 1806, a year after Dalton, one of the founders of modern atomic theory, formulated the law of multiple ratios in the Proceedings of the Manchester Literary and Philosophical Society, Clément and Desormes published detailed data on the acceleration of the oxidation of sulfur dioxide in the presence of nitrogen oxides at room temperature. production of sulfuric acid. Six years later, in the Technology Journal, Kirchhoff presented the results of his observations on the accelerating effect of dilute mineral acids on the hydrolysis of starch to glucose. These two observations opened the era of experimental study of chemical phenomena unusual for that time, to which the Swedish chemist Berzelius gave the general name “catalysis” in 1835 from the Greek word “kataloo” - to destroy. This, in a nutshell, is the history of the discovery of catalysis, which with good reason should be classified as one of the fundamental phenomena of nature.
Now we should give the modern and most generally accepted definition of catalysis, and then some general classification of catalytic processes, since this is where any exact science begins. As you know, “physics is what physicists do (the same can be said about chemistry).” Following this instruction from Bergman, one could confine oneself to the statement that “catalysis is something that both chemists and physicists do.” But, naturally, such a humorous explanation is not enough, and since the time of Berzelius, many scientific definitions have been given to the concept of “catalysis.” In our opinion, the best definition is formulated by G.K. Vereskov: “Phenomenologically, catalysis can be defined as the excitation of chemical reactions or a change in their speed under the influence of substances - catalysts that repeatedly enter into intermediate chemical interactions with reaction participants and restore their chemical composition after each cycle of intermediate interactions "
The strangest thing about this definition is its final part - the substance that accelerates the chemical process is not consumed. If it is necessary to accelerate the movement of a heavy body, it is pushed and, therefore, energy is expended on this. The more energy spent, the greater the speed the body acquires. Ideally, the amount of energy expended will be exactly equal to the kinetic energy acquired by the body. This reveals the fundamental law of nature - conservation of energy.
Prominent figures in chemistry on catalysis
I. Berzelius (1837):
“Known substances, when in contact with other substances, have such an influence on the latter that a chemical effect occurs - some substances are destroyed, others are formed again, without the body, whose presence causes these transformations, taking any part in them. We call the cause that produces these phenomena the catalytic force.”
M. Faraday (1840).
“Catalytic phenomena can be explained by the known properties of matter, without supplying it with any new force.”
P. Rashig (1906):
“Catalysis is a change in the structure of a molecule caused by external factors, resulting in a change in chemical properties.”
E. Abel (1913):
“I came to the conclusion that catalysis occurs as a result of a reaction, rather than the mere presence of a substance.”
L. Gurvich (1916):
“Catalytically acting bodies, attracting moving molecules to themselves much more strongly than bodies lacking catalytic action, thereby increasing the impact force of molecules hitting their surface.”
G. K. Boreskov (1968):
“Once upon a time, catalysis was considered as a special, slightly mysterious phenomenon, with specific laws, the disclosure of which was supposed to immediately solve the selection problem in a general form. Now we know that this is not so. Catalysis in its essence is a chemical phenomenon. The change in the reaction rate during catalytic action is due to the intermediate chemical interaction of the reactants with the catalyst.”
If we do not take into account Berzelius's unsuccessful attempt to connect the observed phenomena with the action of a hidden "catalytic force", then, as can be seen from the above speeches, the discussion was mainly around the physical and chemical aspects of catalysis. For a long time, the energy theory of catalysis was especially popular, linking the process of excitation of molecules with resonant migration of energy.
The catalyst interacts with reacting molecules, forming unstable intermediate compounds that decompose to release the reaction product and the chemically unchanged catalyst. Our modern knowledge is best reflected in Boreskov's statement.
Here, however, the question arises: could the catalyst, since it itself chemically participates in the reaction, create a new equilibrium state? If this were so, then the idea of the chemical participation of the catalyst would immediately conflict with the law of conservation of energy. To avoid this, scientists were forced to accept and then experimentally prove that the catalyst accelerates the reaction not only in the forward but also in the reverse direction. Those compounds that change both the rate and equilibrium of a reaction are not catalysts in the strict sense of the word.
It remains for us to add that usually in the presence of a catalyst there is an acceleration of chemical reactions, and this phenomenon is called “positive” catalysis in contrast to “negative”, in which the introduction of a catalyst into the reaction system causes a decrease in the rate. Strictly speaking, catalysis always increases the rate of a reaction, but sometimes the acceleration of one of the stages (for example, the emergence of a new chain termination pathway) leads to the observed inhibition of a chemical reaction.
We will consider only positive catalysis, which is accepted
divided into the following types:
a) homogeneous, when the reaction mixture and catalyst are either in a liquid or gaseous state;
b) heterogeneous - the catalyst is in the form of a solid substance, and the reacting compounds are in the form of a solution or gaseous mixture; (This is the most common type of catalysis, thus carried out at the interface between two phases.)
c) enzymatic - catalysts are complex protein formations that accelerate the course of biologically important reactions in plant and animal organisms. (Enzyme catalysis can be either homogeneous or heterogeneous, but due to the specific features of the action of enzymes, it is advisable to separate this type of catalysis into an independent area.)
A little about industrial catalysis
I will remember for the rest of my life the distillation of the resulting condensate, carried out according to Engler, in which already at the beginning of the experiment the gasoline fraction was 67%. We stayed late into the night waiting for enough quantity to be tested on a race car, but we thought that due to the high gasoline yield the engine would detonate. I'll never forget my excitement the next morning when the car climbed the hill without detonating!
Y. Goodry, 1957
These words belong to Goodry, an outstanding researcher in the field of practical use of catalysis. They were said by him at the International Congress on Catalysis in 1957, twenty years after, as a result of a long routine search, a fundamentally new method was finally developed for converting heavy oil residues into high-octane motor fuel - catalytic cracking of oil. According to Goodry, the idea of using catalysis to break down petroleum hydrocarbons into low-molecular-weight products with a lower boiling point came to his mind back in 1927. But only ten years later in Paulsboro (USA) at the Socony-Mobil oil refinery was The world's first industrial catalytic cracking unit was built using compounds of silicon oxide and aluminum oxide (aluminosilicate) as a catalyst. After 1937, catalytic methods of oil refining, which included many different chemical processes, became firmly established in the oil industry. The main ones include: cleavage of carbon-carbon bonds and isomerization of primary cleavage products (cracking); dehydrogenation and isomerization of hydrocarbons with the formation of branched and aromatic molecules (min); hydrogenation of unsaturated hydrocarbons with time-based removal of sulfur and nitrogen in the form of hydrogen sulfide and ammonia (hydrotreating); introduction of hydrocarbon fragments into the benzene ring of aromatic compounds (alkylation).
Let us recall that until 1937, oil cracking was carried out exclusively by thermal methods: oil fractions were processed at a temperature of about 500 ° C and a pressure of 50-60 atm. Catalytic cracking is carried out at ~50-500°C and atmospheric pressure in the presence of bentonite clays or artificially prepared aluminosilicates. This produces higher octane fuel and aromatic hydrocarbons, which can be used for further chemical processing. Approximately one third of the world's motor fuels are produced by cracking. It should be noted that more than a quarter of all global chemical products are produced from various types of chemical petroleum products.
An important component of industrial catalysts are promoters - substances, the addition of which to the catalyst in small quantities (percents or fractions of a percent) increases its activity, selectivity or stability. If the promoter is added to the catalyst in large quantities or is itself catalytically active, the catalyst is called mixed. Substances, the effect of which on a catalyst leads to a decrease in its activity or complete cessation of the catalytic action, are called catalytic poisons. There are cases when the same additive to the catalyst is a promoter at some concentrations, and a poison at others. in heterogeneous catalysis (see below), substance carriers that are themselves catalytically inactive or have little activity are widely used.
The “geography” of catalysis is unusually wide and diverse - from large-scale production of organic substances to the control of vital biochemical processes in a living cell (and, possibly, also to “controlled” nuclear fusion) - and covers the field of activity of researchers of many profiles and directions. Of course, we do not intend to list all the main areas of use of catalysis and will give only some examples from the chemical industry.
You can start, for example, with the problem of “fixing” air nitrogen - an extremely inert substance that even reacts with oxygen only at 3500-4000 ° C. Natural resources of fixed nitrogen are limited, while huge quantities of nitrogen compounds are needed for the production of agricultural products. The resources of free nitrogen are practically unlimited. Chemists convert it to a bound (and more reactive) state using a reaction
Na + 3H 2NH.
In order for the rate of this reaction to be acceptable from a practical point of view, high temperatures and pressures are needed. However, with increasing temperature, the equilibrium of the reaction gradually shifts towards the formation of the starting substances. On the other hand, the lower the temperature and the more complete the reaction of ammonia formation occurs, the more noticeably the rate of the process decreases. The search for a compromise between factors acting in different directions led Haber (1907) to the creation of an industrial method for converting a nitrogen-hydrogen mixture into ammonia at 500° C and 300 atm. Now this is the main method of producing ammonia, which is widely used in the production of fertilizers, nitric acid (catalytic oxidation of ammonia over platinum), ammonium salts, soda, hydrocyanic acid, etc.
Catalysis is used to hydrogenate unsaturated chemical compounds. Thus, by treating carbon monoxide with hydrogen in the presence of zinc-chromium catalysts at 400 ° C and a pressure of about 300 atm, methanol CO + 2H CH3OH is obtained, which is widely used as a solvent for the starting product for the production of other valuable substances. In particular, by oxidizing it on silver or copper catalyst, you can get formaldehyde
CH3OH + O HСОН + НО
an equally important substance, consumed in large quantities for the synthesis of plastics.
Methanol can also be used to produce hydrogen CH3OH + H O 3H + CO.
As a result of processing vegetable oils with hydrogen in the presence of nickel catalysts, solid fats (in particular, margarine) are formed. Catalysis is used to accelerate the hydrolysis processes of polyatomic organic compounds, mainly plant carbohydrate-containing compounds. Here mineral acids serve as catalysts. When plant raw materials (wood waste, sunflower husks, straw, etc.) are treated with acid, polysaccharide chains (cellulose, pentosans) are split to form food and feed products, glucose, xylose, furfural and a number of other oxygen-containing derivatives. By combining the processes of acid hydrolysis and catalytic hydrogenation (the so-called hydrogenolysis), carried out under more severe conditions (200°C, 50 atm), products of deep cleavage of molecular chains, glycerin, ethylene glycol, and propylene glycol, are obtained. These substances are used in the production of explosives, glyphthalic resins, as well as plasticizers and solvents.
The production of polymers and synthetic fibers cannot be ignored. Here, the pride of domestic science is the process developed by S.V. Lebedev (1932) for producing synthetic rubber according to the scheme: ethyl alcohol - butadiene - polybutadiene. Catalytic reactions in this process are carried out at the first stage - dehydrogenation and simultaneous dehydration of ethyl alcohol. Nowadays, butadiene and isoprene are also produced by dehydrogenation of hydrocarbons of normal structure on chromium-aluminum catalysts, in particular from butane. This made it possible to involve natural gas resources and waste gases during oil refining into the production of synthetic rubber.
A major event in the production of polymers was the discovery of stereospecific polymerization of unsaturated compounds in the presence of mixed Ziegler-Natta catalysts (1952). An example of this type of catalyst is a mixture of triethylaluminum and titanium tetrachloride. The use of these catalysts made it possible to obtain macromolecules with a specific spatial configuration of monomer units. Products made from such polymers have excellent performance properties. Worth mentioning is the exceptionally active catalytic system developed by Morton (1947), codenamed "alfine" and consisting of a mixture of allyl sodium, sodium isopropylate and sodium chloride. In the presence of alfin, butadiene polymerizes in a few minutes to form chains containing tens and hundreds of thousands of monomer units.
The role of catalysis in ecology
Catalysis is expected to play a huge role in solving the most pressing problem - environmental protection. According to Cousteau, the globe resembles “a car rushing lonely through outer space without an exhaust pipe.” Indeed, we have nowhere to dump waste except into the same environment in which we live. This is a rather sad topic, but it is worth talking about, since a person is already beginning to feel the negative sides of his hectic or largely uncontrolled activity. Catalytic chemists are persistently working on this pro-. problem and have already achieved some results. Special devices have been developed for afterburning vehicle exhaust gases, operating on the basis of catalytic oxidation of harmful gas components. Catalysts and conditions for the neutralization of waste gases from chemical plants have been selected. Catalytic filters are designed as cartridges filled with metal mesh or ceramic materials coated with catalytic agents; These filters operate at 250-350° C.
We have given the temperature and pressure at which reactions are catalyzed under industrial conditions, partly in order to compare them with the conditions of similar chemical reactions occurring in organisms of the plant and animal world. The latter have a much higher speed at ordinary temperature and pressure. This is achieved with the help of biological catalysts - products of the long evolution of life on Earth, inevitably accompanied by millions of errors and dead ends. We will probably not soon recognize the tortuous path that nature followed in search of effective organic structures with their fantastic ability to accelerate processes in living organisms under mild conditions.
Energy barrier
All catalytic reactions are a spontaneous process, i.e. flow in the direction of a decrease in the Gibbs energy - a decrease in the energy of the system.
It has long been known that nonionic molecules react much less often than they collide with each other. Arrhenius explained this fact by suggesting that molecules can react only if at the moment of collision they have an energy reserve not lower than a certain critical value. In this case they are called "active molecules".
A. Rezchik, 1945
Such a theory exists, it is the theory of absolute reaction rates, which began with the theoretical studies of Polyani in 1931. Below we will get acquainted with it, but for now we will pay attention to another law of chemical kinetics, known as the Arrhenius law (1889). The law connects the reaction rate constant with a certain energy characteristic characteristic of a given reaction, called activation energy E.
(26)
where k0 is a constant, or pre-exponential factor; R is the gas constant equal to 1.987 cal/degree*mol", T is the temperature in degrees Kelvin; e is the base of natural logarithms.
To find the activation energy E, study the reaction rate at different temperatures and find the rate constant for each value of T. Since equation (26) contains two unknown quantities - k0 and E, they proceed as follows. Logarithm (26)
+ (27)
plot the dependence of Ln(k) on 1/T and determine the slope, which is equal to E/R. Usually, not natural, but decimal logarithms are used.
(28)
(The last number is the module for converting natural logarithms into decimals, multiplied by the value R = l.987.)
The Arrhenius law is associated with a symbolic representation of the reaction path that is widespread in chemistry in the form of an energy diagram shown in Fig. 1. The meaning of this image is as follows: in order for molecules to move from one state H1 to another H2, they must have a reserve of internal energy not less than a certain critical value E. The states Hl and H2 are thus separated by some energy barrier with a height equal to activation energy E, and the lower the barrier height, the greater the reaction rate in accordance with the Arrhenius equation. This increase is not infinite: even in the absence of a barrier (E = O), the reaction will proceed at a certain finite (and not infinitely large) speed, since at E = 0 at the same time
.
An important property of the energy diagram is that the initial H1 and final H2 levels do not depend on the barrier height. We can arbitrarily change the height of the barrier E (if, of course, we know how to do this practically), but at the same time the levels H1 and H2 will remain unchanged if certain external conditions are given - temperature, pressure, etc. In other words, they exist in in principle, different paths along which molecules can move from one fixed state to another, including those on which the energy barrier is zero; there cannot be such a case when E< О.
Passing through the energy barrier
Arrhenius's law is an experimentally established fact. He states that the rate of reaction increases with temperature for the vast majority of reactions, but he says nothing about exactly how the reaction system overcomes the energy threshold. It makes sense to understand this in more detail by introducing certain model representations.
Let us imagine a simple exchange interaction reaction
A - B+C - D => A - D+B - C. (a)
If we could divide the interaction of molecules into separate elementary acts - the breaking of old bonds A - B and C - D and the formation of new bonds B - C and A - D, we would observe the following picture: first, the reaction system absorbed the energy from the outside necessary for breaking initial chemical bonds, and then it was released due to the formation of new bonds. This would be reflected in the energy diagram by a certain curve, the maximum of which corresponded to the dissociation energy of old bonds (Fig. 2).
In reality, the activation energy is always lower than the dissociation energy. Consequently, the reaction proceeds in such a way that the energy of breaking bonds is partially compensated by the energy released during the formation of new bonds. Physically this could happen, for example, as follows. At the moment of atoms B and C approaching each other, a B...C bond is formed and, at the same time, bonds A - B and C - D are loosened. In this case, energy partially “flows” from one compartment to another. It is not difficult to imagine that as reacting atoms come closer together, sooner or later a moment comes when all bonds are in an equally loose state.
Polyany and Eyring call this state transitional and attribute to it all the properties of ordinary molecules, with the exception that the vibrations of atoms along the line along which the approach and breaking of bonds lead to the formation of final products.
For these reasons, it is reasonable to introduce into the reaction scheme some transition state A - B+C - D =>* => A - D+B - C (b)
corresponding to the top of the energy barrier. Only those molecules that have a certain amount of internal energy can climb to the top. They acquire this energy as a result of collisions with other molecules. Those of them that have not gained the required amount of energy roll back down to replenish their reserves. The climb to the top is the most difficult part of the journey, but, having reached the pass, the molecules roll down uncontrollably. There is no turning back for them. The more molecules on top, the faster the reaction rate. These simple arguments allow us to represent the rate constant as the product of two quantities:
k = a* K*,
one of which a* is the constant of the monomolecular transformation of the activated complex into reaction products, which has the dimension of frequency, and the second K* is the equilibrium constant of the formation of the transition complex.
In general, a reaction can proceed from starting materials to final products in different ways, i.e. through various passes (activation energy values). However, the reaction, as a rule, proceeds along one of the paths, the one where the energy costs are the least.
Accelerating chemical reactions using acids and bases is the most common technique used by chemists in their daily work. We will consider only catalysis with “protic” acids. In this case, the catalytically active principle is the hydronium ion formed during the dissociation of acid in aqueous solutions
(b)
At moderate concentrations, hydrochloric acid completely disintegrates into ions. Carboxylic acids, in particular acetic acid (b), do not dissociate completely: a certain equilibrium is established between the ions and undissociated forms of the acid. As a measure of the dissociation of weak acids, the dissociation constant is chosen, which for acetic acid is equal to 1.75 * 10^-15 mol/liter;
The H+ proton does not exist in its pure form in solution, since it is more advantageous for it to combine with a water molecule. (However, for brevity, simply write H+, meaning hydronium ion by this symbol.)
it is convenient to express the concentration of hydrogen ions in units of pH = -lg, i.e. in units of exponent (this unit of measurement was first introduced by Sørensen).
The higher the acid concentration (or acidity of the medium), the faster the reaction rate, but only up to a certain pH value. Based on this, let's try to understand two questions:
1) why does the rate increase with increasing concentration of H+ ions (i.e. with decreasing pH);
2) why the reaction slows down when acid is added above a certain norm;
There is a strong belief that it all starts with the attack of a nitrogen atom on the carbon of the carbonyl group. Nitrogen has two unpaired electrons, and carbon not only does not have them, but even has some deficiency in electron density. The carbonyl group is said to be polarized - part of the outer electron cloud is shifted towards the oxygen atom.
When we introduce an acid into the reaction mixture, the resulting hydrogen ions begin to attack the molecules of both partners, but only one type of attack will contribute to their chemical interaction - an attack on carbonyl oxygen. Why is this so? Yes, because the coordination of a proton with an oxygen atom will lead to a shift in the electron density from the carbon atom towards the proton. The carbon will be exposed and it will be able to easily accept the electrons of the nitrogen atom. This, in fact, is the nature of acid catalysis. It is not difficult to understand that the greater the concentration of H+ ions (we repeat - the lower the pH), the greater the concentration of protonated (oxygen) aldehyde molecules, the higher the reaction rate should be.
We have considered, of course, a simplified scheme of acid catalysis, but it is a good illustration of how the phenomenon is studied and what conclusions can be reached as a result of knowing the dependence of the reaction rate on the concentration of hydrogen ions. The analysis of the phenomena of catalysis under the influence of OH- ions (basic catalysis) differs fundamentally little from the analysis of acid catalysis just considered.
When studying the catalysis of organic reactions in strongly acidic media, difficulties are encountered that are usually easily overcome when working with dilute acids. But we will not focus on this; we will only pay attention to what kind of information is obtained by studying concentration dependences.
Homogeneous catalysis
Among the numerous catalytic reactions, catalysis occupies a special place in chain reactions.
“Chain reactions, as is known, are those chemical and physical processes in which the formation of some active particles (active centers) in a substance or in a mixture of substances leads to the fact that each of the active particles causes a whole series (chain) of sequential transformations of the substance” ( Emanuel, 1957).
This mechanism for the development of the process is possible due to the fact that the active particle interacts with the substance, forming not only reaction products, but also a new active particle (one, two or more), capable of a new reaction of converting the substance, etc. The resulting chain transformations of the substance continue until the active particle disappears from the system (the “death” of the active particle and the chain break occur). The most difficult stage in this case is the nucleation of active particles (for example, free radicals), but after nucleation the chain of transformations occurs easily.
Chain reactions are widespread in nature. Polymerization, chlorination, oxidation and many other chemical processes follow a chain, or rather, a radical-chain (with the participation of radicals) mechanism.
The mechanism of oxidation of organic compounds (in early stages) has now been established quite thoroughly. If we denote the oxidizing substance R-H (where H is the hydrogen atom that has the lowest bond strength with the rest of the R molecule), then this mechanism can be written in the following form:
Catalysts, such as compounds of variable valence metals, can influence any of the considered stages of the process.
Let us now dwell on the role of catalysts in processes of degenerate chain branching. The interaction of hydroperoxide with a metal can lead to both acceleration and inhibition of the oxidation reaction of organic substances by metal compounds of variable valency, depending on the nature of the products formed during the decomposition of hydroperoxide. Metal compounds form a complex with hydroperoxides, which disintegrates in the “cage” of the solvent medium; if the radicals formed during the decomposition of the complex manage to leave the cell, then they initiate the process (positive catalysis). If these radicals do not have time to leave and recombine in the cell into molecular inactive products, then this will lead to a slowdown of the radical chain process (negative catalysis), since in this case hydroperoxide, a potential supplier of new radicals, is wasted.
So far we have considered only shallow stages of oxidation processes; at deeper stages, for example in the case of the oxidation of hydrocarbons, acids, alcohols, ketones, and aldehydes are formed, which can also react with the catalyst and serve as an additional source of free radicals in the reaction, i.e. in this case there will be additional degenerate chain branching.
Heterogeneous catalysis
Unfortunately, until now, despite a fairly large number of theories and hypotheses in the field of catalysis, many fundamental discoveries have been made by chance or as a result of a simple empirical approach. As you know, a mercury catalyst for the sulfonation of aromatic hydrocarbons was accidentally discovered by M.A. Ilyinsky, who accidentally broke a mercury thermometer: mercury entered the reactor, and the reaction began. In a similar way, the now well-known Ziegler catalysts, which at one time opened a new era in the polymerization process, were discovered.
Naturally, this path of development of the doctrine of catalysis does not correspond to the modern level of science, and this explains the increased interest in the study of the elementary stages of processes in heterogeneous catalytic reactions. These studies are a prelude to creating a strictly scientific basis for the selection of highly efficient catalysts.
In many cases, the role of heterogeneous catalysts in the oxidation process is reduced to the adsorption of an organic compound and oxygen with the formation of an adsorbed complex of these substances on the surface of the catalyst. This complex loosens the bonds of the components and makes them more reactive. In some cases, the catalyst adsorbs only one component, which dissociates into radicals. For example, propylene on cuprous oxide dissociates to form an allylic radical, which then easily reacts with oxygen.
It turned out that the catalytic activity of metals of variable valence largely depends on the filling of d-orbitals in cations of metal oxides.
According to the catalytic activity in the decomposition reaction of many hydroperoxides, metal compounds are ranked as follows:
We considered one of the possible ways to initiate the process - the interaction of hydroperoxide with the catalyst. However, in the case of oxidation, the reaction of heterogeneous chain initiation can occur both through decomposition into hydroperoxide radicals and through the interaction of the hydrocarbon with oxygen activated by the catalyst surface. The initiation of chains may be due to the participation of the charged form of the organic compound RH+, formed during the interaction of RH with the catalyst. This is the case with catalysis in chain initiation (nucleation and branching) reactions. The role of heterogeneous catalysts in chain continuation reactions is especially clearly emphasized by changes in the rate and direction of isomerization of peroxide radicals.
Catalysis in biochemistry
Enzymatic catalysis is inextricably linked with the life activity of plant and animal organisms. Many of the vital chemical reactions that occur in a cell (something like ten thousand) are controlled by special organic catalysts called enzymes or enzymes. The term “special” should not be given close attention, since it is already known what these enzymes are made of. Nature chose for this purpose one single building material - amino acids and connected them into polypeptide chains of various lengths and in different sequences
This is the so-called primary structure of the enzyme, where R are side residues, or the most important functional groups of proteins, possibly acting as active centers of enzymes. These side groups bear the main load during the operation of the enzyme, while the peptide chain plays the role of a supporting skeleton. According to the Pauling-Corey structural model, it is folded into a helix, which in the normal state is stabilized by hydrogen bonds between acidic and basic centers:
For some enzymes, the complete amino acid composition and sequence of their location in the chain, as well as a complex spatial structure, have been established. But this still very often cannot help us answer two main questions:
1) why enzymes are so selective and accelerate the chemical transformations of molecules only of a well-defined structure (which we also know);
2) how the enzyme reduces the energy barrier, i.e., chooses an energetically more favorable path, so that reactions can proceed at normal temperatures.
Strict selectivity and high speed are two main features of enzymatic catalysis that distinguish it from laboratory and industrial catalysis. None of the man-made catalysts (with the possible exception of 2-hydroxypyridine) can compare with enzymes in the strength and selectivity of their action on organic molecules.
The activity of an enzyme, like any other catalyst, also depends on temperature: with increasing temperature, the rate of the enzymatic reaction also increases. At the same time, attention is drawn to the sharp decrease in activation energy E compared to the non-catalytic reaction. True, this does not always happen. There are many cases where the speed increases due to an increase in the temperature-independent pre-exponential factor in the Arrhenius equation.
To illustrate the unusually high efficiency of enzymes, we will give two examples and compare the effect of a conventional acid catalyst with enzymatic ones. As a measure of activity, we present all three parameters of the Arrhenius equation - the rate constant (k, l/mol*sec), the pre-exponential factor A and the activation energy (E, kcal/mol).
These examples are especially interesting in that in the first case, the increase in the rate constant in the presence of urease is mainly due to a decrease in the activation energy (by 17-18 kcal/mol), while in the second, the effect of myosin on the rate constant is due to an increase in the pre-exponential factor .
The activity of enzymes also depends on the acidity of the environment in which the chemical reaction takes place. It is noteworthy that the curve of this dependence on the pH of the medium resembles the bell-shaped curves of acid-base catalysis. (See Fig. 3)
It seems that enzymes are given the right to decide what is beneficial for them in this particular case - to organize a stronger connection of the active center with the substrate molecule or to disorder their structure.
It is difficult to say what considerations guide the enzyme when choosing the path of substrate activation. In any case, the study of the kinetics of the enzymatic reaction and the thermodynamics of the formation of intermediate complexes, although it provides valuable quantitative information, does not allow us to fully reveal the molecular and electronic mechanism of the enzyme. Here, as in the study of ordinary chemical reactions, one has to follow the path of modeling - roughly speaking, inventing molecular mechanisms that at least would not contradict experimental data and the elementary logic of chemical reactions. The trouble is that with a sufficiently developed imagination, you can come up with quite a lot of such “good” mechanisms. Below we will get acquainted with some of these model concepts, and now we will look at how researchers establish the nature of the active centers of enzymes.
An increase in the acidity of the medium will be beneficial for some elementary stages and unfavorable for others. Given these competing facts, it is easy to guess that there must be some optimal acidity of the environment at which the enzyme can work with maximum efficiency.
So, analysis of the dependence of the rate on pH is a very effective means of identifying the functional groups of the protein molecule of the enzyme involved in the process of activation of substrate molecules. Knowing the nature of active centers, one can imagine how they work. Of course, in this case we have to use the same ideas about the mechanism of elementary acts that were developed in the study of ordinary reactions of organic chemistry. There is no need to introduce any special mechanisms. There is a firm belief that the work of an enzyme is ultimately reduced to a set of simple operations similar to those performed during the interaction of organic molecules under ordinary test tube conditions.
So we know:
1) At least two functional groups take part in enzymatic catalysis, and the mechanism of the enzymatic reaction includes a certain sequence of elementary acts, which provides an energetically more favorable route than the non-enzymatic reaction;
2) active centers on the polypeptide chain are located so that at a certain moment and in a certain place they can interact with the substrate molecule and carry out a series of coordinated chemical acts.
List of references for the examination essay in chemistry
Gromov Sergei in publishing support from
Klochkova Yuri.
“In the world of catalysis.”, M, Nauka, 1977
“Big Chemical Encyclopedia”, vol. 2, M., Soviet Encyclopedia, 1990
“Schoolchildren's Guide to Chemistry”, M., Slovo, 1995
“Chemistry 11”, M., Education, 1992
“Organic chemistry”., Enlightenment, 1991
“General Chemistry”, Minsk, University, 1995