%D, %d %M %y
Time: %h~:~%m

Home  / GENERAL CHEMISTRY Textbook / Chapter 9.** Chemical Reactions

Chapter 9.** Chemical Reactions

In the course of chemical transformation, old bonds between atoms break and new ones form. Millions of chemical reactions take place simultaneously in live organisms at temperatures of 20°C to 40°C.

As we have seen in the previous section, we need energy of about 200 kJ/mol (i.e., about 2,000°C) to break a chemical bond. Thus, the main question concerning chemical reactions is:

How do chemical reactions proceed at temperatures of 20°C to 30°C if, in order to break a bond, a temperature of more than 2,000°C is actually required?

Before we go on to study complex reactions in organisms, let's take up some simple examples, which can be studied in any school lab. One of the best-known and most instructive experiments is as follows:

The teacher prepares a mixture of two gases: oxygen and hydrogen in a test-tube. He demonstrates that these gases will never start a reaction by themselves. Then he lets the gas mixture go into a tin can through an opening at the bottom and he passes an electric current (spark) into the gas. The current causes an explosion, which is the result of a momentary interaction of the hydrogen and oxygen via the reaction: 

 2H2+ O2 → 2H2

That is, in the process of this reaction, the bonds in molecules H2 and O2 were broken, and new bonds were formed between the hydrogen and the oxygen. It is worth mentioning here again that in order to break the bonds in molecules of hydrogen and oxygen a temperature of more than 2,000°C is required. 

Of course we can suppose that the spark had caused the momentary heating of the mixture to this temperature. However, a test made by passing a spark separately into either of the two gases has shown that the temperature of these gases practically does not change. 

The motor of a car receives air (oxygen) and fumes of gasoline (mixture of hydrocarbons) but the motor will not work because the reaction takes place only when the motor is switched on (when the spark appears). The process that takes place in a motor is well described by the following reaction: 

2C8H18 + 25O2 → 16CO2 + 18H2

We can also observe in a similar experiment, that if we mix water solutions, say, of barium chloride (BaCl2) and sodium sulfate (Na2SO4), we will not get a transparent mixture as with other combinations; we will get some sediment of barium sulfate that is formed via the reaction: 

BaCl2 + Na2SO4 →  2NaCl + BaSO4          

In this case, the bonds between barium (Ba) and chloride (Cl), and between sodium (Na+) and sulfate (SO42-) will break, though they actually require a temperature of more than 5,000°C for their rupture.

It is of interest to note that we had dried both these salts at a temperature over 200° C and then mixed them; there was no reaction between them. 

Other such phenomena are observed in interactions of chloride (Cl2) and hydrogen (H2). If we mix these gases in darkness, we will see that there is no reaction between them. If, however, we momentarily irradiate the mixture with light, then the following reaction takes place: 

      Cl2 + H 2  → 2HCl 

What conclusions can we make on the basis of these examples?

  1. Reactions between molecules for which energy of 200 kJ/mol is required for bond breaking (i.e., heating to over 2,000° C) in reality do not proceed at normal temperatures — with mixed substances. 
  2. To proceed with the reaction, we need a momentary energetic stimulant (spark or flash) or the introduction of a third substance (water).  

In order to understand the essence of these momentary energetic stimulants, scientists have studied the composition of gas after such coercions. They have found that gas (oxygen + hydrogen) contains separate atoms of oxygen and hydrogen though the initial gases did not contain separate atoms. The concentration of separate atoms amounted to nothing (less than 0.01%) of the number of molecules in the mixture.  

Thus, there can be no suggestion that the momentary action of the spark or light leads to bond breaking in the molecules and to dissociation into atoms which then unite to form new molecules.

At the same time, we must understand how the presence of species (separate atoms) in tiny concentrations leads to an explosive interaction of the whole mass of hitherto passive molecules.

During the development of lab experiments, scientists invariably studied the interaction of separate atoms and ions with molecules. They have found that an atom of chloride rapidly reacts with a hydrogen molecule along the following scheme:

      Cl + H2 → HCl + H, 

and that an atom of hydrogen also rapidly reacts with a chloride molecule even at room temperature according to the following scheme:

      H + Cl2 →HCl + Cl 

These results help us to answer the question: Why and how do reactions cause chemical transformation of molecules? 

If we take the transformation of the mixture hydrogen + chloride, the reaction can be described as follows. When this mixture is irradiated with light, a small portion of the chloride molecules dissociate into atoms along the scheme:                         

      Cl2  + hv→ 2Cl.

The cited scheme also explains why small concentrations of free atoms cause the transformation of many substances. Every atom performs a large number of cyclic transformations, i.e., this scheme shows how the chemical reaction between molecules proceeds, and the role of short-term energetic action upon reaction mixtures.

However, this scheme does not answer the main question, which arose during the studies of chemical reactions. Indeed, it has been found that the interaction of molecules is possible via the chain route where the carriers of the chain are radicals or ions. The mechanism scheme of the cited above reactions, has two more steps: 

Cl2 + hv → 2Cl and Cl + H2 → HCl + H

 Cl2 + H →HCl + Cl 

In the first of these reactions, all was logical. The molecule of chlorine under the influence of a portion of light (strong, energetic action) broke up into atoms. But in the second reaction, in the presence of the chlorine atom, the bond in the hydrogen molecule dissociated at room temperature without any additional energetic influence. That is, this scheme for chemical transformations allows specifying the main question which is now formulated as follows:

Why is it that the interactions of H2 + Cl  HCl + H and Cl2 + H  HCl + Cl in the course of which bonds between hydrogen atoms in molecule H2 break; and bonds between chlorine atoms in molecule Cl2 (that require heating up to over 2000° C and 700° C respectively) break during reaction with chlorine and hydrogen atoms at room temperature? 

It is of interest to note that the answer to this question was first received theoretically (via discourse) and only then was it proved experimentally. To be more precise, the basis for the theory was a more strictly limited experimental material than was the experimental material that appeared after the elaboration of the theory.

Then what is the answer to the above-mentioned question? 

The authors of this theory have suggested that the reaction of the active species (atom, ion, radical, etc.) proceeds not in one step, but in three: 1) association, 2) electronic isomerization,  3) dissociation.

For example, in the case of interaction between a chlorine atom (Cl) and hydrogen (H2) the reaction mechanism is described thus: 

      Cl + H : H → Cl...H : H         (1)

      Cl...H : H ⇆ Cl : H...H          (2)

      Cl : H...H → Cl : H + H        (3) 

Thereon, the colon (:) represents electrons on the atoms' outermost shell. Three dots (Cl...) indicate Van der Waals bonds (VWB).

      Then how does the above reaction scheme actually answer the above-mentioned question?

Each of the three steps have been studied experimentally, and it was found that each of the reactions proceeds at room temperature, and that the reaction rate of step #3 is close to that of the interaction of a chlorine atom with hydrogen. Then it was found that steps #1 and #2 proceed with rates by two orders of magnitude higher than that of step #3. That is, the rate of the whole reaction can be defined by the rate of step #3, i.e., step #3 is the rate-determining step.

Now let's make out what the rate determining step is, or figure out why the slow reaction stage controls the speed of the whole process that proceeds subsequently via several stages. The speed of the whole chemical process is measured by the amount of final product, in the given case, by the amount of HCl molecules formed, say, in one minute.

As already indicated, the reaction speed of #1 and #2 are more than 100 times faster than that of #3. That is, reactions #1 and #2 reach equilibrium. Analogous processes, consisting of successive stages, are found in everyday life.

For example, let's take a poorly organized car assembly line at a factory. The body of a car is moving on a conveyor. During the movement of the conveyor, the body gradually gets its wheels, its motor, its interior, etc. One man is working on each stage of this conveyor.

In order to complete each operation (of #1, #2, and #3) this would require, say, 10, 60, and 5 minutes respectively, while for all the previous and consequent operations less than 5 minutes is spent. Such a conveyor will produce only one car every 60 minutes; that is, the time required for installing a motor; that is, the time that corresponds to the slowest stage of the sequential process.

Relatively, stage 2 (the motor installation stage) is the stage that defines the speed of the whole process. Since the speed of the slowest stage controls the speed of the whole process, this stage is called the limiting stage, since the speed of the whole process (assembly of one car) cannot be less than 60 minutes. That is, this stage limits the amount of cars produced in 60 minutes — to one car.

If we want to increase the production of cars, we should increase the speed of the defining stage, increase the number of workers on the motor-installation stage. The increase of the number of workers on stages 1 and 3 will not have any effect on the production speed.

Relative to such consequential multistage reaction, the defined HCl formation speed of stage #3 (the slowest) in the process is equal to the speed of HCl formation during the process, since stage #3 is the limiting stage. The influence of parameter change (like the increase of temperature) is obvious (just as the change of the number of workers on the conveyor belt) mainly on stage #3, the defining stage for increasing the speed of the whole process.

The increase of temperature, just as the increase of the number of workers on the conveyor belt, increases the speed of all the stages in the process. But the temperature increase will cause the speed of the whole process to be close to that of the slowest stage.

We have noted that the speed increase of the fast stages does not effect the whole process. That is, its speed is always defined by the defining stage— the speed of the slowest stage.

When we study the influence of temperature on the speed of final product formation in a process proceeding via some sequential stages, we define the influence of temperature (temperature coefficient or activation energy) as the very slowest of the sequential reactions.   

Here is a well-known problem: A person has to travel from point A to point B in the shortest time possible. He is offered two possibilities: 1) he can drive half the way at 30 miles per hour and walk the second half of the way at 3 miles per hour; or 2) he can drive at 90 miles per hour half the way and walk the second half of the way at 1.5 miles per hour.

      Which of these two methods should the traveler choose to get to point B sooner?

Spend a while solving this problem; then try to realize the answer if all the conditions remain unchanged, but the cars travel at the rate of 120 and 150 miles per hour. Your reflection in this respect should lead you to the conclusion that the question which is the quickest way? depends on the correlation of the walking rates, but not on that of the driving rates. That is, the time of transition from point A to point B is actually defined by the rate of speed of the traveler's transition while walking, i.e., the rate of the slower transition step. 

Likewise, the speed rate of the whole multi-step reaction is defined by the rate of the slowest reaction step. It is this rate that is regarded as the rate-determining step, and in the above-mentioned scheme step #3 is the rate-determining step

      Cl : H...H → Cl : H + H

The VWB, with energy of less than 20 kJ/mol, breaks up at this step.

In the case of a thermal break-up of a hydrogen molecule, the reaction proceeds via a single step according to the following scheme: 

      H : H  → H˙+ H˙

In this case, the covalent bond, whose energy is equal to more than 200 kJ/mol, will break up on the rate-determining step. The ten-fold difference in the energies of bonds, which break on the rate-determining steps, explains why in the presence of radicals the bond breaking reaction proceeds at room temperature, while for breaking such bonds without active species, a temperature of more than 2,000°C is required.

According to the scheme for the interaction of a chlorine atom with hydrogen, a VWB is formed between the hydrogen atoms as a result of electronic isomerization. As a result of this isomerization  (transition of one electron), the covalent bond (H : H) is changed into a VWB.

Positive and negative ions react like radicals with saturated molecules do. Thus, for example, it has been experimentally confirmed that the interaction of a positively charged ion of potassium (K+) with sodium chloride proceeds via the mechanism as follows: 

K+  + Cl:Na → K+← :Cl:Na  ⇆  K:Cl:→Na→ K:Cl + Na+

In this case, we can suppose that the formation of associate occurs via the DAB between the calcium (K+) cation and the sodium chlorine (NaCl) at the expense of the lone electronic pair of the chlorine atom. During DAB formation, according to the theory of chemical bonding, the bond between the chlorine and sodium weakens. The breaking of the weakened covalent-polar bond in the gas phase is accelerated at the expense of the use of part of the energy that is discharged during DAB formation between the calcium and the chlorine.  

As you see, 2 electrons transit from bond Na:Cl to bond K:Cl.

Thus, the presence of the electronic isomerization step in reaction schemes explains why and how chemical bond breaking reactions take place at room temperature.

In school textbooks, the interaction of bromide anion with gaseous chlorine is demonstrated by the oxidation-reduction reaction. When we add a water solution of chlorine to the colorless solution of sodium bromide, the NaBr solution becomes orange-red because of the formation of bromine. Thus, we get the reaction: 

 Cl2 + 2 Br‾ → 2 Cl‾ + Br2  

However, the direct interaction of bromide anion (Br¯) with chlorine (Cl2) accompanied by the transition of the electrons from the bromide anion (Br¯) to the chlorine molecule is impossible since the ionization energy of bromide anion (Br¯) is equal to 3.36 eV, while the affinity of the gaseous chlorine to the electrons is equal to 2,38 eV.

In the framework of the offered approach, the interaction of bromide anion with gaseous chlorine is described analogously to that of the interaction of the anion with a molecule. This process proceeds in two phases: 

Br¯ + Cl : Cl → Br¯ ... :Cl : Cl ⇆ Br : Cl ... Cl → Br : Cl + Cl¯

Br¯ + Br : Cl → Br ... Br : Cl ⇆ Br : Br + Cl¯         

The enthalpy (electrostatic component) of bonds BrCl,  Cl2, and Br2 are equal to 1.14, 1.18, and 0.93 eV/mol, correspondingly.

The affinities of the chlorine (Cl) and bromide (Br) atoms to the electrons are equal to 3.6 eV and 3.36 eV correspondingly.

That is, there is not loss of energy in reaction:  

Br¯ + Cl2 → Br Cl + Cl¯  

which is equal to the difference between the electrostatic components of bonds Cl2 (1.18 eV) and BrCl (1.14 eV) equal to 0.04 eV (1.18 - 1,14) and there is an energy gain equal to the difference in the affinities to the electrons of the bromide (3.36 eV) and chlorine (3.6 eV) atoms equal to 0.24 eV. That is, the energy gain in this reaction is 0.2 eV (0.24 - 0.04).

The energy gain in the course of the second reaction is:

 Br¯ + Br - Cl → Br2 + Cl¯ 

The electrostatic component in bonding energy Br2 is equal to 0.93eV. Relatively, the energy loss at the expense of the formation of bond Br - Br instead of bond Br - Cl is equal to 0.14 eV (1.07 - 0.93). The energy gain at this stage, which is reached at the expense of the difference in the affinity of the chlorine and bromide atoms to the electron, just as in the first stage, that is, 0.24 eV; that  increases the energy loss (0.14 eV). Thus, the reaction proceeds with an energy gain.

The intermediate products offered by this reaction— BrClCl and BrBrCl — are the same as those of the well-known stable compound I3. Likewise the interaction of anion bromide (Br¯) with halogens and other oxidation-reduction reactions are described in textbooks.

The analogy is in the fact that during the reaction, the bonding electrons (one or two) transit from one energetic minimum to another. For example, the interaction of metals of group I (Li, Na, etc.) with oxygen proceeds thus:           

        ..         ..
Na:Na + :О: :О Na:Na ... :О: :О

       ..        ..
Na:Na :О: :О Na: ... Na :О: :О

Analogous to the oxidation-reduction reactions, there are interactions between acids and bases of proton transition reactions. Just as in the case of oxidation-reduction reactions, the weakening of the old (strong) bond and the formation of a new one occurs as a result of electronic isomerization. And just as in the case of interaction of anions with molecules, electronic pairs take part in the isomerization.

The new electronic pair sort of 'squeezes out' another electronic pair from the atom's shell during the isomerization process.

Unlike the interaction of anion with molecules, in this case the associate forms not at the expense of the Van der Waals bond (VWB), but at the expense of the formation of the hydrogen bond. That is, the reaction proceeds along the following scheme: 

A : H + : B → A : H ... : B ⇆ A : ... H : B ⇆ A : + HB, 

where B is the molecule with a lone pair on one of the atoms (that does not partake in bond formation).

Since the transition speed of the proton in the liquid phase is much greater than that of the other atoms, in this case, besides the electronic isomerization, the transition of the proton can play a role here thereby increasing the transition speed of the system from one minimum to another.

With the formation of the intermediate compound at the expense of the hydrogen bond, acid dissociation proceeds in the elementary fashion as indicated in the following scheme: 

H2O : + H : Cl ⇆ H2O : ... H : Cl ⇆ H2O : H ... Cl ⇆ H3O+  + Cl- 

and the neutralization of acids by the bases.  

HO- + HCl ⇆ HO- : ... H : Cl ⇆ H : O : H ... Cl- ⇆ H2O + Cl-     

The dissociation of salts proceeds with DAB formation (O : Na) and hydrogen bond formation (Cl : H). 

This example illustrates exactly how, under the influence of water, the bond in molecule NaCl, for whose eruption in the gas phase or in a carbon-hydrogen solvent requires heating up to 5,000°C, actually falls apart into ions at room temperature.

As seen from the above, the formation of the intermediate compound occurs at the expense of the formation of all the types of dynamic bonds described in the section devoted to chemical bonding.

Generally, the chemical reaction can be illustrated thus:

The molecules  that enter the reaction break up into ions, radicals, etc., thereby forming intermediate compounds with the initial molecules. These compounds differ from the initial molecules by the fact that the intermediate compound has more electronic isomers, which have practically the same potential energy, as compared to that of the initial molecules.

In the course of electronic isomerization, the electrons transit from one minimum to another. Since the transition speed of the nuclei is by more than one order of magnitude smaller than that of the electrons, the nuclei occupy an intermediate position in the molecules.

The existence of each electronic isomer is defined by the electronic energy of the given isomer. The greater is the attraction energy of the electrons to the nuclei, the greater is the isomer concentration in the mixture (and the greater is the bonding energy). Proportionally it is as follows:  E / 2; where E is the energy of the electronic isomer. The energy of the electrons is measured in values of about 1,500 kJ/mol.

The electrostatic component of chemical bonding is equal to about half the sum of the electrostatic components of bonds in the initial (non-isomerized) associate; that is, half the sum of the electrostatic components of a covalent and hydrogen bonds, or a VWB bond; i.e., equal to about 100 kJ/mol, which comprises less than 7% of the average energy value of the electrons in the molecule. The latter presupposes the existence of electron-nuclear isomers in measurable amounts in the mixture.

As already indicated, the concentration of isomers exponentially depends on the potential electronic energy of the given isomer eE/2. On the other hand, the reaction speed of the isomer's transition to the final product exponentially depends on the energy of the breaking bond: 

      W = k e-EcB

The greater is the bonding energy, the smaller is the reaction speed. Or,  the greater the bonding energy, the more electronic isomer there is in the mixture, and the smaller is its final product transformation speed.   That is, in the first approach, the transformation speed and the activation energy can be evaluated via the breaking speed and by the energy of the weak bond in the associate (VWB or hydrogen bond).

The bonding energy in the Van der Waals complex (VWC) is equal to about 15 kJ/mol; that is, this complex is more stable from the thermodynamic viewpoint than the products formed of it. 

      Why doesn't the reaction cease on stage 2 (stage for the formation of the VWC)?

Indeed, the VWC is more thermodynamically stable than the equilibrium in the system

      Cl : H...H ⇆Cl : H + H 

is shifted to the left with a normal temperature (300 K). The speeds of the direct and reverse reactions according to experimental data, are very large even at 300 K. The reaction reaches equilibrium in less than 1 second. The amount of HCl in the beginning of the reaction is much smaller than that of initial products in Cl2 and H2. Interaction of H with HCl leads to the formation of initial H2 and Cl2 products via the scheme:

 HCl + H → H2 + Cl ;  That is:  HCl + Cl → Cl2 + H 

This scheme is accomplished only after producing concentrations of HCl in the system that is commensurable with the concentration of initial substances (H2 and Cl2).

Previously, the formed hydrogen atom reacted with Cl2 of which there was more than HCl along the scheme: 

      Cl2 + H →  HCl + Cl ; that is, 

—together with the formation of the final product and the Cl atom. 

However, the following question still remains:

      Why are radicals and ions considered as active species? Why can't saturated molecules react between themselves?

We have thus returned to the question given at the beginning of this chapter, only now we can answer it on a higher level. In accordance with the following chemical reaction scheme:     

1 2 3
Cl+H:H ⇆  Cl...H:H ⇆  Cl:H...H →  Cl:H+H.

—the reaction rate is equal to the rate of the third step which is directly proportional to the concentration of the isomerized associate (IA) (i.e., to the concentration of compound Cl:H...H) and which exponentially depends on the energy of the VWB (H...H) of this compound.

Concentration IA depends on the concentration of the nonisomerized associate and on the rate of the isomerization step. The greater the concentration of the nonisomerized associate (in our example: Cl...H:H) and the greater the isomerization rate, the greater is the concentration of isomerized associate in the compound, and the greater is the rate of the whole reaction. The concentration of non-isomerized associates is defined by the concentration of initial products  (in our case, Cl, H2 and bonding energy Cl...H in compound Cl...H:H).

In order to answer the above offered question: Why don't saturated molecules react among themselves along the mechanism: association -electronic isomerization - dissociation? we should compare the probable interaction mechanism for saturated molecules and its interaction along the radical mechanism.

The molecular interaction mechanism for hydrogen (H2) and chlorine (Cl2), according to the theory, can be shown as follows:   

1 H:H 2 Н...Н 3
H:Н+Cl:Cl  : :   .. ..  2НCl

  Сl:Cl Cl...Cl

Now let's compare the rates of each of the steps of the radical and molecular mechanisms. We will begin with step 1.

The origin of the four-part complex, as a result of the association reaction (step 1), is not only a rare phenomenon as compared to the origin of the associate between radical and saturated molecules; it is hardly possible. The obstacle in its formation is the mutual repulsion of the two electronic circles between the hydrogen and chlorine atoms.

There is a much greater possibility in the formation of a linear complex of the H:H...Cl:Cl type. Out of this complex, as a result of electronic isomerization via the scheme we get:  

 H:H...Cl:Cl ⇆ H....H:Cl...Cl

An associate is formed containing two VWBs and one covalent bond. The concentration of this associate is much smaller than the initial, non-isomerized associate. The initial associate has two covalent bonds (H:H...Cl:Cl) and one VWB (H:H...Cl:Cl). The final associate has one covalent bond (H:Cl) and two VWBs, i.e., the electronic energy in this associate is much greater than in the initial one - over 100 kJ/mol.

The concentration of a non-isomerized associate is defined by the bonding energy between the molecules. The greater the energy,— the greater is the associate concentration. The VWB energy between the saturated molecules comprises less than 5 kJ/mol; to break the bond of a radical-saturated molecule, energy of more than 20 kJ/mol is required.   

When studying electronic isomerization reaction, it was found that the isomerization rate depends on the distance between the atoms and on the number of electrons transiting in the course of isomerization. Thus, for example, isomerization reaction:  

      ClŸ... H:H → H:Cl ... H

where one electron transits a distance of 1.5Å in the course of the isomerization, the time is equal to 10-13 sec.

The transition of two electrons in the isomerization reaction: 

      K+... Cl:Na → K:Cl ... Na+ 

—is possible at a time equal to about 10-12 and by changing two electrons in the following reaction:

      F: ... Cl:Cl → F:Cl... Cl: 

The isomerization rate becomes smaller than that of the association reaction, i.e., it takes place during more than 10-11 sec., that is, the isomerization rate in the case of molecular interaction is by two orders of magnitude smaller than in the case of the radical route.

And finally, what is most important, during the breaking of the VWB in an isomerized associate, according to the scheme:                

      H....H:Cl...Cl˙→ H˙..H:Cl + Cl˙ 

—a radical is formed that enters reaction with hydrogen molecules along the chain mechanism described above.

The formation of radicals during the interaction of hydrogen and chlorine is not due to the interaction of hydrogen and chlorine, but it is due to the more rapid reaction along the following route: 

Cl : Cl + Cl : Cl → Cl : Cl...Cl : Cl → Cl˙...Cl : Cl...Cl˙ → 2Cl +Cl2.              

As a result of this reaction, the final product is not produced. The greater radical formation rate in this route is conditioned by the fact that the bonding energy in Cl2 is much smaller than in H2.

When studying the interaction mechanism for chlorine with hydrogen at high temperatures, it was found that as a result of direct interaction of chlorine molecules and hydrogen, less than 0.001% of the final product is formed. That is, after the formation of about 0.05% of the radicals from Cl2 the reaction proceeds along the chain radical route.

Thus the molecules can start an interaction between themselves. This interaction proceeds along the same mechanism: 

 association - electronic isomerization - dissociation. 

However, this interaction proceeds much slower than does the interaction of active species (first of all — radicals and ions) with saturated molecules. This is conditioned by: a much smaller energy gain during the formation of the initial associate: saturated molecule -- saturated molecule as compared to the associate: radical - saturated molecule; a great loss of energy as a result of the isomerization reaction of the initial associate; a smaller electronic isomerization speed.

The greater activity of the radicals and ions (as compared to that of saturated molecules) is caused by: 1) a greater bonding energy of the radical - saturated molecule, 2) a greater associate isomerization rate, and 3) a greater energy gain during isomerization (as a result of which the number of covalent and VWBs does not change). Interactions between saturated molecules proceed along the chain mechanism; radicals or ions play the role of the active intermediate species.

There is almost no reaction between saturated molecules at normal conditions (T = 20°). The reaction speed between saturated molecules increases abruptly when there are active species like ions, radicals, etc. in the system.

In cases where the initial saturated molecules form stronger bonds with each other, than do the common VWBs, and where the formed associates turn into electromers, which have electronic energies identical to those of the initial associates, as a result of electronic isomerization, the interaction between these molecules proceeds with a great speed, usually limited only by diffusion. Examples of such reactions are the dissociation of salts in water, the neutralization of acids with alkalis, and the Lewis bases. These reactions have been described in detail above.

These reaction speeds practically do not depend on the temperature, which implies that the temperature does not have any influence on the electronic isomerization speed, which is much greater than that of the diffusion.

The paradigm, regarding the fact that reactions occur as a result of the collision of molecules with high kinetic energies, has existed for more than 150 years.

In 1868 Lotar Mayer announced: "Chemical phenomena must be treated as if they were problems of mechanics." This announcement serves as an epigraph to the chapter Rates and Mechanisms of Chemical Reactions by Richard E. Dickerson, Harry B. Grey, and Gilbert P. Haight, Jr.  issued by Chemical Principles in 1979.In textbooks already issued at the beginning of the twenty first century, there are visual aids for students that illustrate drawings where the interacting molecules approach each other at a great speed thus then they collide and break up:



Relative to the model for the observation of the influence of the parameters on the reaction speed, this is limited by the observance of the influence of the temperature and of the initial substance concentration on the reaction speed.

To illustrate the model, precipitation reactions are given, then there are reactions for the neutralization of acids via bases, and there are oxidation-reduction reactions proceeding in compounds.

All these reactions proceed at room temperature. The kinetic energy of molecules at room temperature is close to 2 - 4 kJ/mol, while the bonding energy of the breaking bonds during the reaction amounts to 200 - 500 kJ/mol. That is, those molecules that have enough kinetic energy to break the initial bonds are measured by the value of e-100.

In the framework of the formerly, and still existing, paradigms, saturated molecules enter the reaction.

During the 150 years of the existence of the mechanical approach to chemical phenomena, the electron was discovered, atom structure was elucidated, it was resolved that chemistry was actually the change of potential energy of the electrons in the outermost shells of the atoms.

During the existence of the paradigm that was confirmed only by the calculations of three or four reactions of the H2 + I2 ⇆ 2HI type, it was found that these reactions proceed along the radical mechanism.

The mechanisms of millions of reaction had been studied — mainly in organic chemistry.

Already, for more then 50 years, reactions in organic chemistry are divided into

  1. nucleophyl reactions (with active anion particle),
  2. electrophyl reactions (active cation particle) and
  3. radical reactions.

It was found that ions and radicals enter reaction with saturated molecules. It was also found that the speed of their interaction with the molecules was by 10 orders of magnitude greater than the speed of the separate molecules' interaction.

In 1982 we offered the theory of elementary interactions (TEI), which was a chemical theory, but not a mechanical theory of chemical reactions. This theory proved that the main role in breaking the old chemical bond is played by the transition of the bonding electrons from the old bond to the new one, formed in the course of the reaction.

As a rule, in the course of this process, the absolute value of the potential energy of the electrons increases. That is, the electrons in the products of the reaction are closer to the nuclei than they are in the initial products. The general models of chemical reactions for the interactions of ions and radicals with saturated molecules can be illustrated thus: 

For cations:

A+ + B : C → A+... B : C⇆  A : B ... C+ → A : B + C+

For radicals:

A˙ + B : C →  A˙... B : C ⇆ A : B...C˙→ A : B + C˙

For anions:

A-: + B : C → A- : ... B : C ⇆ A : B...C- → (A : B) +  C- : 

In the case of cations and radicals, the breaking of the old bond and the scatter of the atoms formerly bonded by this bond, is conditioned by the repulsion of the nuclei of atoms B and C because of the absence or decrease of the bonding electrons that were situated between them.

In the case of anionic reactions, the scattering of the atoms is conditioned by the repulsion between the electrons of the outermost shell of atom B and anion C-.  

      What is the physical nature of the phenomenon called electron-nuclear isomerization?   

Answer: The electrons in the molecules are situated in energy minimums of 1-2 eV. 

But first let's take the question which arises in this connection:  

      At the expense of what energy does the electron transit from one minimum to another?   

The difficulty in answering this question was caused by one of the main principles of quantum mechanics, which presupposed that the electron could not change its energy in such small portions as 2 or 3 eV.

Traditionally, the energy exchange between electrons was never presupposed in quantum chemistry, while in the framework of our theory, the energy exchange between electrons is not only presupposed, but is even proven by the exchange mechanism via the energy exchange between the electrons.  

The electronic energy in molecules is defined by the energy of the nuclei. The Maxwell-Boltzmann law describes the distribution of the vibrational, rotational, and translational energies of the nuclei. Relatively, we can suppose that this law describes the distribution of the electrons' energies as well.

Besides this energy-transmitting mechanism, which can be regarded as a mechanism of translational movement, the atomic systems (electrons and nuclei) should also have a mechanism for changing the vibrational energy.

If, in a mechanism of translational movement, the energy transits from particles with a greater energy to particles with a smaller energy, then, in the case of vibrational energy, the energy transition is defined by the vibrational phase and not by the energy. 

Recall the energy exchange between pendulums vibrating on one axis. Both mechanisms allow us to presuppose an accumulation of sufficient energy on the minor portions of electrons, so that the latter could overcome the barrier that divides the minimums, since the average energy of only the outermost electrons in the system is equal to 12 eV.

The evaluation of the barrier energy that the electrons have to overcome, is possible judging by the following consideration:

According to the chemical bonding theory, the energy gain during chemical bond formation is conditioned by the approach of the nuclei to the bonding electrons located between them. In the system D ... H : T, the transition of the electrons from covalent bonding  (H : T) to VWB (D ... H) with the formation of D : H ... T, is caused by the receding of the electron from atom T and its approach to atom H. That is, the topmost barrier limit, during the transition of one electron from bond H : T onto atom H, amounts to half of the energy gain (enthalpy) during bond formation and is equal to 1 eV.

The receding from one nucleus of atom T with a simultaneous approach to the nucleus of atom (H) does not significantly change the general potential energy of the system, and therefore, the height of the barrier is much lower (close to zero) and not 1 eV.

However, if we consider that the electron recedes from atom T and approaches two D ... H atoms during the transition, then the height of the barrier is equal to 0  —  0, 2 eV; that is, a height that is greater than what the electrons with an average energy of 12 eV can overcome.

The electron in the atom has kinetic energy of about 10 eV.  The following equation for kinetic energy  

  E =  mv2  / 2, 

—helps to  evaluate the speed of the electrons: 

      v = (2E / m)0.5   =  (2 · 10 · 96 · 10/ 9.1 · 10-31  ·  6.02 · 1023)0.5

      ≈10m/ sec = 1016Å / sec.


This speed can be compared with the transitional speed of an electron measured experimentally and calculated as follows: 

 v = 1017 e-2R/L

—where R is the distance Å  (in our case 1Å); and L is the coefficient equal to 6.5 in conjugated systems, and — 0.7  in  non-conjugated ones.

According to this equation, the transitional speed of the electron for conjugated systems can be estimated at 1017 Å/sec, while for non-conjugated ones, 1016 Å/sec, i.e., values close to those already calculated. The above-mentioned explanation and the data obtained from the experiment, coincide semi-quantitatively.

As mentioned above, the vibrational system can contain another mechanism, which imitates the transition of electrons.  In this case, it does not really transit but simply passes its energy to another electron, which happens to be rotating, for example, in a parallel circle.

As a result of this exchange, the bond that had hitherto been covalent, for example, in the system D ... H : H, becomes a  VWB, and vice versa — the system becomes D : H ... H. The time of these transitions can be evaluated by the electron's rotation frequency, which is about 10-16 sec.  The simultaneous existence of both mechanisms is possible.

The mechanism, being characteristic of translational movement, should bring the system to equilibrium as far as the nuclei are concerned.  The system is in dynamic equilibrium defined by thermodynamics, i.e., the speed of the mutual transitions is defined by the free energy change.

In the vibrational mechanism, the system transits from one state to another with an identical speed, irrespective of the change in the free energy.  Indeed, according to the experiment, there are cases of electronic transition, when the transition speed depends on the free energy change, and cases when it does not.

The mechanism, which is usually present in the textbooks describing electron transition in molecular systems supposes tunneling transition of electrons. At the present time the most popular is the mechanism proposed by R. Marcus (Nobel Laureate in 1993) according to which the tunnel transition is a rare phenomenon.

Our explanation of isomerization consists in the following. In the course of electron isomerization electrons shift between the three atoms A, B and C.

In an extreme position the electrons are in a small potential well 0.2 eV, at the same time they are affected by two charges of atom B and C. The electrons begin to move toward these atoms. They "skip" the atom B inertially. The situation repeats itself in the extreme position of the electrons in case of the atom C but now the electron is affected by another pair A and B. The reverse motion between atoms A and C occurs and, as the result of it, bonds between atoms A, B and C become equal along the length and energy. On the one hand vibrational motions of nuclei initiate isomerization, on the other hand the motion speed of the nuclei is several orders of magnitude lower than the electron velocities, so the nuclei in a molecule take up an intermediate position. As it will be shown later, as a result of the phenomenon of isomerization the depth of potential well for electrons decreases (see Appendix C: Semiquantitative simulation of electrical conductivity in metals and nonmetals, p. 117). For the proceeding of electron isomerization is necessary that the particles approached each other at a distance of 3 Å and remained at that distance for no less than 10-13 s. The speed of isomerization is estimated to be 1,013 Å/s, and the speed of substitution reactions is 109, so isomerization does not restrict the proceeding of a chemical reaction anyhow (for the calculations, see "How chemical bond forms and chemical reactions proceed").

Chapter 6. Molecule structure >> 
Chapter 7.** Chemical Bonds in Solid Bodies >>
Chapter 8. Three-dimensional structures of chemical compounds  >>
Chapter 9.** Chemical Reactions 
Chapter 10.** Catalysis  >>
Chapter 11. Physical and chemical properties of substances  >>