Molecule Structure

It has been proven experimentally that atoms are bonded into molecules, therefore, in order to break up a molecule into atoms, energy should be applied. For example, to break a hydrogen molecule (H₂) into two atoms, it is necessary to heat the hydrogen to a temperature over 3,000 K or spend an energy of more than 400 kJ/mol. At 3,000 K the nuclei of hydrogen atoms have an energy of about 20 kJ/mol. As the atoms contain no other particles besides electrons, we can conclude that the energy which is spent on the breaking of the molecule into atoms, is actually used to change the electrons' energies.

The energy gain during bond formation is conditioned by the following: During molecule formation the electrons start moving in a field with a unified positive charge of the bonding nuclei. The electrons' energies are proportional to the square of the nuclear charges and the inter-nuclear repulsion energy is proportional to the first power of the charges. That is why when the atoms approach each other to a certain point the system's energy decreases. When the molecule is formed out of the atoms the kinetic and potential energies of the electrons increase. The energy gain is conditioned by the fact that the increase of potential energy is twice greater than the increase of the electrons’ kinetic energy.

In the dual-atomic molecule, the positively charged nuclei are united with a circle of electrons rotating on a plane perpendicular to the axis connecting the nuclei. All the main parameters have been calculated for this system. (See section 2.4)

As a result of the calculations, it was found that:

1) When the bonding atoms have identical first ionization potentials (FIPs), the electrons' rotation plane is at the same distance from the bonding atoms. If the atoms' FIPs are different, the rotation plane will be shifted towards the atom with a greater FIP, and the molecule gets a dipole moment.

2) Two electrons partake in the formation of a chemical bond; there is a greater energy gain when two electrons combine (not 1 or 3).

3) Both bonding electrons enter the outer shells of the atoms being bonded. That is, during chemical bond formation, the number of electrons in the outer shell of the atoms being bonded, is increases by one.

During covalent bond formation, the bonding pair of electrons is composed of the electrons of the atoms being bonded: one from each atom. The number of covalent bonds which one atom can form with other atoms (atom’s valence) is limited, due to inter-electronic repulsion, by the number of electrons which the given atom can connect to its outer shell with an energy gain.

The maximal number of electrons which the outer electronic layer can contain is equal to the number of electrons in the shells of the noble gas nearest to it in the table. Noble gases cannot bond electrons in the outer electronic layer (the electron affinity values for noble gas atoms are smaller than zero) therefore, they cannot form covalent bonds.

Therefore, the number of covalent bonds which an atom can form (valence) is defined by the number of electrons in the outer layer of the given atom (one outer-layer electron is spent on the formation of one bond) and by the maximal number of electrons which can exist in the outer layer of the given atom (the number of outer-layer electrons is increased by 1).

Besides covalent bonds, there exist donor-acceptor bonds (DAB) and Van der Waals bonds (VWB).

During the formation of DABs, both bonding electrons come from one atom. This bond is formed between molecules, one of which contains atoms whose shells are not filled up to the nearest inert gas (acceptor atoms) while the other molecules contain atoms whose shells have nonbonding electron pairs. An example of such a bond is what we find in coordinate (complex) compounds. The energy of the DAB is about two times weaker than that in a covalent bond.

The VWB is possible between atoms whose shells are filled up to the shell of the inert gas. Examples of such bonds are those between atoms of inert (noble) gases and intermolecular bonds, say, between hydrocarbons. The energies of these bonds are about 10 times smaller than those of covalent bonds.

Chemical Reactions and Catalysis

In the course of chemical transformation, the old bonds between atoms break and new ones form. Millions of chemical reactions take place simultaneously in live organisms at temperatures of 20° to 30° C.

As we have seen in the previous section, we need an 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° to 30° C if, in order to break the 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 goes 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 a brief 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 ® 2H₂O

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 (i.e., the spark appears). The process taking place in a motor is well described by the reaction:

2C₈H₁₈ + 500₂ ® 16CO₂ + 18H₂O

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 a sediment of barium sulfate which 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 (SO₄^2-) will break though they actually require a temperature of more than 5,000° C for their rupture.

It is of interest to note that if we dry both these salts at a temperature over 200° C and then mix them, there will be 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 with bond breaking for which an energy of 200 kJ/mol is required (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 coercion’s. They have found that gas (oxygen + hydrogen) contained separate atoms of oxygen and hydrogen though the initial gases did not contain separate atoms. The concentration of separate atoms amounted to next 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 hydrogen rapidly reacts with a chloride molecule even at room temperature according to the following scheme:

H + Cl₂ ® HCl + Cl

and that an atom of chloride also rapidly reacts with a hydrogen molecule along the following scheme:

Cl + H₂ ® HCl + H

These results allow us to answer the question: Why and how do reactions lead to chemical transformations of molecules? If we take the transformation of the mixture hydrogen + chloride the reaction can be described by the following equations. 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 schemes of the cited above reactions, besides step:

Cl₂ + hv® 2Cl

was also the following step:

Cl + H₂ ® HCl + H

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 to specify the main question which is now formulated as follows: Why is it that the reaction for breaking the chemical bond in a molecule of Cl₂ requires an energy of more than 100 kJ/mol (the energy of a light quantum) while the breaking of a Cl₂ bond in a molecule where a hydrogen atom (H) and a chlorine molecule (Cl₂) interact without any light influence at room temperature.

It is of interest to note that the answer to this question was first received theoretically (i.e., 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 that experimental material which appeared after the elaboration of the theory.

Then what is the answer to the above mentioned question?

The authors of this theory have made a suggestion to the effect that the reaction of the active spicies (ion, radical, etc.) proceeds not in one step, but in three: 1) association, 2) electronic isomerization, 3) dissociation.

For example, in the case of an 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)

Here and thereon, the dots represent electrons on the outer orbit of the atoms. Three dots at the bottom (Cl...) indicate Van der Waals bonds (VWB).

How does such a reaction scheme 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 speed of reaction 3 is close to that of the interaction of the chlorine atom with hydrogen. Then it was found that steps 1 and 2 proceed with speeds by two orders higher than that of step 3. That is, the speed of the whole reaction can be defined by the speed of step 3, i.e., step 3 is the rate-determining step.

Now let’s try to explain as simply as possible what the rate determining step is. 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 use 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 will 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 speeds, but not on that of the driving speeds. That is, the time of transition from point A to point B is actually defined by the speed of the traveler’s transition while walking, i.e., the speed of the slower transition step.

Likewise, the speed of the whole multi-step reaction is defined by the speed of the slowest reaction step. It is this speed 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 an energy of less than 20 kJ/mol, breaks at this step.

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

H : H ® H& + H&     

That is, in this case, the covalent bond, whose energy is equal to more than 200 kJ/mol, will break 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 3,000° 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 ^U K:Cl...Na^{+ ®} K:Cl + Na⁺

In this case, as you see, 2 electrons transit from bond Na:Cl to bond K:Cl. In the case when a saturated molecule is attacked by a negatively charged ion, the following reaction takes place:

F:^- + Cl₂ ® F:^- ... Cl : Cl ® F : Cl... Cl ® F : Cl + Cl:^-

In this case, during electronic isomerization two fluorine electrons sort of substitute two chlorine electrons.

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

However, some questions still remain:

Why is it that saturated molecules cannot react between themselves? Why are radicals and ions considered to be active species? Indeed, why shouldn’t chlorine interact with hydrogen molecules according to this scheme?

That is, we have returned to the question given at the beginning of this chapter, only now we can answer it on a higher level.

.According to the following chemical reaction scheme:

the reaction speed is equal to the speed of the third step which is directly proportional to the concentration of the isomerized associate (IA) (in the given example: to the concentration of compound Cl:H...H) and exponentially depends on the energy of the VWB (H...H) of this compound.

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

In order to answer the above offered question: Why is it that saturated molecules do not react among themselves along the mechanism: association - electronic isomerization - dissociation? let’s compare the possible interaction mechanism for saturated molecules with their interaction along the radical mechanism.

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

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

The origin of the four-member 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 the radical and saturated molecule, but is hardly possible. The main obstacle in its formation is, first of all, the mutual repulsion of the two electronic rings (circles) between the hydrogen and chlorine atoms.

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

H:H...Cl:Cl ® H& ...H:Cl...Cl&

an associate is formed containing two VWBs and one covalent bond. The concentration of this associate should be 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 the non-isomerized associate is defined by the bonding energy between the molecules. The greater the energy - the greater the concentration of the associate. It is known that the VWB energy between the saturated molecules comprises less than 5 kJ/mol, while to break the bond radical - saturated molecule an energy of more than 20 kJ/mol is required.

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

Cl& ... H:H ® H:Cl ... H&   

where one electron transits a distance of 1.5A 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 speed becomes smaller than that of the association reaction, i.e., it takes place during more than 10^-11 sec., that is, the isomerization speed in the case of molecular interaction should be 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 which enters the reaction with saturated molecules, in this case, with hydrogen molecules along the chain mechanism described above. Actually, the formation of radicals in the system 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 speed 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: the greater bonding energy of the radical - saturated molecule, a greater associate isomerization speed, and a greater energy gain during the isomerization reaction (as a result of which the number of covalent and VWBs does not change). The interaction between saturated molecules proceeds along the chain mechanism where the role of the active intermediate species is played by radicals or ions. There is practically no reaction between saturated molecules at normal conditions (T=20° ). The reaction speed between saturated molecules increases abruptly when there are active species in the system like ions, radicals, etc. The concentration of active species can be created thermally (heating the reaction system), by radiation, or by electric charges.

Another way of creating the necessary concentration of active species in the system is by introducing available active species like ions, radicals, or substances causing the dissociation of initial substances into active species, or substances that dissociate into ions and radicals much easier than the initial saturated molecules. Examples of reaction acceleration at the expense of substances causing the dissociation of the initial substances into active species are reactions in water solutions where the initial substances (like salts) dissociate into ions. In all these cases the mechanism of the reactions does not change. The reaction proceeds along the chain ionic or radical mechanism.

Another way of accelerating the reaction is by means of introducing substances in the presence of which molecules can react directly.

As indicated above, the low mutual interaction speed of the saturated molecules is conditioned first of all by the following: a low concentration of the initially formed associate, a low concentration of an isomerized associate. Both of these reasons are conditioned by a small bonding energy between the molecules in the associate and by a small electronic isomerization speed. That is, in order to have the molecules react among themselves, we must introduce a substance which forms stronger bonds with both molecules.

A direct interaction of saturated molecules (AB and CD) in the presence of a catalyst (K) can be shown on a scheme as follows:

That is, at the end of the reaction we get the final products and the initial catalyst, i.e., the catalyst does not undergo any chemical changes as a result of this reaction. In order to have substance K as a catalyst, bonds AB-K and ABK-CD should be much stronger than VWB AB-CD. On the other hand, the rate-determining step is step 4. Saturated molecule bonds with catalysts break in the course of this reaction. That is, in order to accelerate the reaction in the presence of catalyst K it is important that the bonds catalyst - saturated molecule are not very strong.

Previously, in chapter Chemical Bonding we have described three types of bonds: VWB, DAB, and covalent bonds (including homo- and heteroatomic covalent bonds). That is, the catalyst should form bonds with saturated molecules, and these bonds must be stronger than VWBs and not so strong as covalent bonds. It can be supposed that if there is acceleration of the molecular interaction, the catalyst forms DABs with the reacting molecules.

In the chapter devoted to chemical bonding and chemical structure, we have elucidated the fact that DABs are most possible between saturated molecules which include atoms with non-bonding electron pairs and with molecules which include atoms with incompletely filled outer electronic shells. Such shells have surface atoms of hard substances and coordinately unsaturated compounds called conences. It is these atoms that catalyze molecular interaction which is accomplished in industrial processes when synthesizing ammonia out of hydrogen and nitrogen, in getting aldehydes out of hydrogen (H₂), carbon oxide (CO), and olefins in the process of hydroformylation, etc. The coordinately unsaturated atoms in the ferment are catalysts of biochemical transformations.

As compared to the first method for accelerating chemical reactions where the reaction accelerates at the expense of introducing active species (ions, radicals) to the system, the second method (acceleration of the direct molecular interaction) has priority from the point of view of selectivity; i.e., as a result of these processes, we generally get individual products.

Conence homogenic catalysis is, from the point of view of both methods, a mixed type of catalysts. Here the acceleration is gained at the expense of the increase of the conence concentration (i.e., active species). The reaction proceeds along the chain mechanism, but during this reaction the initial molecules do not disintegrate and in the associates they do not form VWBs (which is characteristic of the first method for accelerating reactions), they form DABs.

Examples of mechanisms for accelerating reactions via the first and second methods are described in this book to say nothing of world chemical literature. [See pages 167 and 272-280.]