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THEORY OF CHEMICAL REACTIONS

Summary

The main questions that were raised in the course of the development of the theory of chemical reactions were as follows:  

1) Why don't all chemical reactions proceed if they are thermo-dynamically possible?

2) Why does the reaction speed increase along the exponent with the increase of temperature?

3) Why is it that in reactions proceeding with bond breaking, the activation (additional) energy is usually much smaller than that necessary to break the bond thermally?  

Indeed, why is it that reactions proceed with the breaking of the chemical bond in normal conditions, while we need a temperature of more than 4,000° C to break such bonds thermally?  Examples of such reactions are interactions of radicals and ions with molecules, catalytic and photochemical reactions. This question was not even touched upon before the beginning of our work. 

In the course of our studies we managed to learn that unlike the opinions accepted in 1980, the interactions between molecules take place not via the scheme: 

AB + CD → AC + BD

but mainly along the chain reaction scheme:

AB ⇆ A + B

A + CD → AC + D

D + AB → DB + A

where A, B, and D are active species (radicals, ions, conences).

That is, unlike the ideology accepted in 1980, the active elements in chemical reactions (species which cause chemical transformations) are not at all species or molecules with a high kinetic energy. The active elements are actually specific chemical species like radicals, ions, conences, etc.

It has been found that the interaction of these species with saturated molecules proceeds via three elementary stages (steps): association — electronic isomerization —  dissociation..

Dissociation is the step that usually limits the process. The given scheme for the procedure of the chemical reactions answered the main questions, which arose during the development of chemical kinetics, mentioned above.

Active species are in thermodynamic equilibrium with the initial molecules. When the temperature decreases, the concentration of active species exponentially drops, causing an abrupt decrease in the speed of the chemical reactions. This dependence of the chemical reaction rate on the temperature explains why all the possible thermodynamic chemical reactions do not proceed at normal temperature (1st question) and why the reaction rate between the molecules exponentially depends on the temperature (2nd question).

The presence of the electronic isomerization step in the mechanism of the chemical reaction answers the 3rd question. The electronic isomerization speed is by many orders of magnitude higher than that of the dissociation step. That is why the kinetic parameters of the electronic isomerization step, and first of all — the activation energy, do not effect the activation energy of the whole reaction process which defines the energy consumption. That is, the energy necessary to break the chemical bond in the initial molecule is equal to the activation energy of the slowest step (dissociation) in the reaction of an active species with a saturated molecule.  

As a result of electronic isomerization, the initial covalent bond, which requires about 400 kJ/mol in order to break, transforms into a Van der Waals bond (VWB), which requires less than 20 kJ/mol for its rupture. Thus it becomes comprehensible why, in order to break a chemical bond in a molecule in the presence of an active species, we need an energy one order smaller than for the thermal rupture of the bond in this molecule (3rd question).

In the framework of the offered scheme, the driving forces of chemical reactions have become comprehensible.

In the light of the above, such notions as catalysis become understandable. A catalyst is a chemical compound which forms (produces) a greater number of active species in the system than initial molecules do at the same temperature.

The second method for accelerating the interaction of molecules is by increasing the concentration of the associates in the system in the presence of catalysts. Thus, for example, a catalyst substance unites other substances that react with each other along the scheme:

AB + K →ABK

ABK + CD → ABKCD

ABKCD → AC + BD + K

where AB and CD are reacting substances.  

The acceleration of the reaction (catalytic action of the substance) is explained thus: Due to the absence of the catalyst, the intermediate compound in the reaction is: AB —CD; while in the presence of the catalyst it is:  AB— K — CD. The speed of the whole reaction in both cases is proportional to the concentration of the intermediate compound. In correlation with concentration AB—CD (without the catalyst) and AB — K —CD (with the catalyst), the concentration of these compounds in the system is defined by bonding energies AB— CD and AB —K — CD.

The bonding energies of both molecules with catalysts are much higher than the bonding energy between themselves, so the concentration of the intermediate compounds with catalysts is much greater, and correspondingly, much greater is the reaction speed. 

This mechanism for reaction acceleration is typical for biological systems in which the catalysts are actually ferments (enzymes), which unite the reacting molecules in their centers.

Besides the above-mentioned mechanism for accelerating reactions, the catalyst can accelerate the reaction by other mechanisms:

1) it can improve the role of chemically activated routes;

2) it can serve as a bridge for the transition of electrons.

As expected, the substances that accelerate reactions in the above-mentioned mechanisms are catalysts; and the essence of catalysis, in the general approach, becomes better understood if compared with such general definitions as:

"A catalyst is a substance that accelerates a reaction. A catalyst carries the reaction along the route requiring less activation energy."