What is a catalytic reaction? Basic principles and types

Most of the processes underlying chemical technology are catalytic reactions. This is due to the fact that with the introduction of the catalyst, the rate of interaction of substances increases significantly. At the same time, manufacturers manage to reduce costs or to get a larger number of reaction products for the same period of time. That is why the study of catalysis is given a lot of attention in the preparation of technologists. However, this phenomenon plays an important role in nature. So, special substances regulate the flow of biochemical reactions in living organisms, thereby affecting the metabolism.

The concept of catalysis

The essence of this chemical phenomenon is to control the rate of conversion of substances using special reagents that can slow down or accelerate this process. At the same time, they speak of positive or negative catalysis. There is also the phenomenon of autocatalysis, when one of the intermediate products of a chemical reaction affects the reaction rate. Catalytic processes are diverse, they differ in mechanisms, the state of aggregation of compounds and direction.

Substances that slow down chemical interactions are called inhibitors, and catalytic accelerating reactions are called catalysts. Both those and others change the reaction rate by multiple intermediate interactions with one or more of its participants. However, they are not part of the products and are restored after the end of the cycle of transformation of substances. Therefore, the participation of the catalyst does not reflect stoichiometrically in the reaction equation, but merely indicate as a condition for the interaction of substances.

Types of catalytic reactions

According to the state of aggregation of substances involved in a chemical reaction, they distinguish:

• homogeneous reactions - reacting substances, products and the catalyst are in the same state of aggregation (phase), while the molecules of the substances are uniformly distributed throughout the volume;
• interphase catalytic reactions - occur at the interface of immiscible liquids, and the role of the catalyst is reduced to the transfer of reagents through it;
• heterogeneous catalytic reactions - in them the catalyst has an aggregate state different from the reactants, and it itself occurs at the interface;
• heterogeneous homogeneous reactions - are initiated at the interface with the catalyst, and continue in the reaction volume;
• microheterogeneous reactions - small particles of solid catalyst form micelles throughout the entire volume of the liquid phase.

There is also redox catalysis, accompanied by a change in the degree of oxidation of the catalyst when interacting with reagents. Such transformations are called catalytic oxidation and reduction reactions. The most common in the chemical industry is the oxidation of sulfur dioxide to trioxide in the production of sulfuric acid.

Types of Catalysts

According to the state of aggregation, the catalysts are liquid (H ₂ SO ₄ , H ₃ PO ₄ ), solid (Pt, V ₂ O ₅ , Al ₂ O ₃ ) and gaseous (BF ₃ ).

By type of substance, the catalysts are classified into:

• metals - can be pure, alloys, whole or deposited on a porous base (Fe, Pt, Ni, Cu);
• metal compounds of the type M _m E _n - the most common oxides are MgO, Al ₂ O ₃ , MoO _3, and others;
• acids and bases - are used for acid-base catalytic reactions, these can be Lewis, Bronsted, and others;
• metal complexes — salts of transition metals, for example PdCl ₂ , Ni (CO) ₄ , are also included in this group;
• enzymes (they are enzymes) are biocatalysts that accelerate the reactions that occur in living organisms.

By the specifics of the electronic structure, d-catalysts with d-electrons and d-orbitals, as well as s, p-catalysts, the center of which is an element with valence s and p-electrons, are distinguished.

Catalyst Properties

For effective use, a fairly extensive list of requirements applies to them, changing for a specific process. But the following two properties of the catalysts are most significant:

• Specificity lies in the ability of catalysts to influence only one reaction or a series of transformations of the same type and not affect the speed of others. So, platinum is most often used in organic hydrogenation reactions.
• Selectivity is characterized by the ability to accelerate one of several possible parallel reactions, thereby increasing the yield of the most important product.

Catalytic reaction rate

The reason for accelerating the interaction of substances is the formation of an active complex with a catalyst, which leads to a decrease in activation energy.

According to the basic postulate of chemical kinetics, the rate of any chemical reaction is directly proportional to the product of the concentrations of the starting materials, which are taken in degrees corresponding to their stoichiometric coefficients:

v = k ∙ C _A ^x ∙ C _B ^y ∙ C _D ^z ,

where k is the rate constant of a chemical reaction, numerically equal to the rate of the same reaction, provided that the concentrations of the starting compounds are 1 mol / L.

According to the Arrhenius equation, k depends on the activation energy:

k = A ∙ exp ^ (- E _A / RT).

The indicated regularities are also valid for catalytic reactions. This also confirms the equation for the ratio of the rate constants:

k _K / k = A _K / A ∙ exp ^ (( - ) / RT),

where the variables with index K refer to catalytic reactions.

Catalytic Reaction Stages

For homogeneous catalytic reactions, two main stages are sufficient:

1. The formation of the activated complex: A + K -> AK.
2. The interaction of the activated complex with other starting materials: AK + B -> C + K.

In general, an equation of the form A + B -> C is written.

The mechanism of heterogeneous catalytic reactions is complex. The following six stages are distinguished:

1. Bringing the starting compounds to the catalyst surface.
2. Adsorption of the starting reagents by the surface of the catalyst and the formation of an intermediate complex: A + B + K -> AVK.
3. Activation of the resulting complex: Α -> Α ^* .
4. The decomposition of the complex compound, while the resulting products are adsorbed by the catalyst: Α ^* -> CDK.
5. Desorption of the obtained products by the surface of the catalyst: CDK -> C + D + K.
6. The removal of products from the catalyst.

Examples of catalytic reactions

Catalysts are used not only in the chemical industry. Any person in his daily life is faced with various catalytic reactions. This, for example, the use of hydrogen peroxide in the treatment of wounds. Hydrogen peroxide, when interacting with blood, begins to decompose under the influence of the enzyme catalase :

2H ₂ O ₂ -> O ₂ + 2H ₂ O.

In modern cars , the exhaust system is equipped with special catalytic chambers that contribute to the decomposition of harmful gaseous substances. For example, platinum or rhodium helps to reduce the level of environmental pollution by nitrogen oxides, which are destroyed with the formation of harmless O ₂ and N ₂ .

Some toothpastes contain enzymes that trigger decomposition of plaque and food debris.