There are numerous reactions in nature that do not obey the Arrhenius equation, the law of mass action. The rates of such reactions cannot be explained by any theory of kinetics or equations of formal kinetics. These are chain reactions.
Chain reactions called reactions that occur with the participation of chemically active particles (free atoms and radicals) and consist of a large number of repeating stages. Chain reactions include combustion reactions, slow oxidation, radioactive decay, nerve impulse transmission, nuclear reactions, etc.
Characteristic features of chain reactions include:
1) the reaction rate does not coincide with the rate calculated according to the theory of active collisions, i.e. W observed >> W calculated;
2) exceptional sensitivity to impurities of positive and negative catalysts that accelerate or slow down the reaction;
3) dependence of the reaction rate on the size, shape, and materials of the vessel. In a vessel with a larger volume, the reaction rate is greater. The reaction rate slows down if fragments of quartz, glass, porcelain, etc. are placed in the free space.
4) The presence of lower and upper limits of ignition or explosion for gas oxidation reactions, below and above which the reactions proceed slowly or do not occur at all.
The peculiarities of these reactions are explained by the chain reaction mechanism, to the development of which Academician N.N. made a significant contribution. Semenov (Fig. 5.6). The active particle “A”, formed by collision or any other way, can be deactivated, but can give an intermediate substance “Z”, which in turn can decompose without forming products or give reaction products and a new active particle. In this case, the activation of one molecule of the starting substances leads to the formation of a large number of molecules of reaction products.
Rice. 5.6. Chain reaction diagram
Basic principles of the theory of chain reactions:
1) atoms or free radicals (particles with an unpaired electron) participate in the reaction. The atom or free radical has exceptional activity towards the valence-rich molecule;
2) when a valence-saturated molecule interacts with an atom or free radical, “free valence” does not disappear, but gives rise to a new active particle (the principle of the indestructibility of free valence).
By “free valence” we mean particles that have an unpaired electron (atom or free radical).
The main stages of the chain reaction:
1) chain nucleation - the elementary stage of a chain reaction leading to the formation of free valence from a valence-saturated molecule;
2) continuation or development of the chain - an elementary stage that occurs with the preservation of free valency and leads to the consumption of starting substances and the formation of reaction products;
3) chain termination - an elementary stage leading to the disappearance of free valence.
An example of a chain reaction is the synthesis of HCl.
H2 + Cl2 = 2HCl
1. Origin of the chain (excitation, initiation):
Cl 2 + hν = 2Cl ˙ (E dissoc Cl 2 = 57 cal/mol),
H2 + hν = 2H ˙ (E dissoc H 2 = 103.3 cal/mol).
Since the binding energy of the chlorine molecule is lower, the nucleation of the chain occurs due to the dissociation of Cl 2 molecules.
2. Continuation or development of the chain:
Cl ˙ + H2 = HCl + H ˙,
H ˙ +Cl 2 = HCl + Cl ˙, those. free valence does not disappear.
3. Open circuit:
H ˙ +H ˙ + walls = H 2
Cl ˙ +Cl ˙ + walls = Cl 2
H ˙ +Cl ˙ + walls = HCl
Chain termination is possible due to the recombination of free radicals (homogeneous process) and a heterogeneous process is chain termination due to the walls of the vessel. In the considered example of a chain reaction, each active particle gives rise to one chain - this is a stationary unbranched chain reaction.
Branched chain reactions are reactions in which the chain development stage occurs with an increase in the number of “free valences”. Its diagram looks like this:
Rice. 7. Scheme of a branched reaction.
An example of a branched reaction is the synthesis of water at high temperature.
2H 2 + O 2 = 2H 2 O
1. Chain initiation: H 2 + O 2 = 2OH ˙,
OH ˙ + H 2 = H 2 O + H ˙
2. Chain development: H ˙ + O 2 = OH ˙ + Ö (biradical)
Ö +H2=OH ˙ +H ˙,
H ˙ + O 2 + H 2 = 2OH ˙ +H ˙ - stage of chain development.
From one active particle several are obtained, each of which gives rise to its own chain.
3. Open circuit: 2H ˙ =H2
A chain reaction is a self-sustaining chemical reaction in which initially appearing products take part in the formation of new products. Chain reactions usually occur at high speed and often have the character of an explosion.
Chain reactions go through three main stages: origin (initiation), development and chain termination.
Rice. 9.13. The energy profile of a reaction (a plot of potential energy versus reaction coordinate) showing a minimum that corresponds to the formation of a reaction intermediate.
Initiation stage. At this stage, the formation of intermediates (intermediate products) occurs. Intermediates can be atoms, ions or neutral molecules. Initiation can be accomplished by light, nuclear radiation, thermal (thermal) energy, anions, or catalysts.
Stage of development. At this stage, intermediates react with the original reactants to form new intermediates and final products. The development stage in chain reactions is repeated many times, which leads to the formation of a large number of final and intermediate products.
Circuit break stage. At this stage, the final consumption of intermediate products or their destruction occurs. As a result, the reaction stops. The chain reaction can break spontaneously or under the influence of special substances - inhibitors.
Chain reactions play an important role in many branches of chemistry, in particular in photochemistry, combustion chemistry, nuclear fission and nuclear fusion reactions (see Chapter 1), and organic chemistry (see Chapters 17-20).
Photochemistry
This branch of chemistry covers chemical processes associated with the effect of light on matter. An example of a photochemical process is photosynthesis.
Many chain reactions are initiated by light. The initiating particle in this case is a photon, which has energy (see Section 1.2). A classic example is the reaction between hydrogen and chlorine in the presence of light
This reaction proceeds explosively. It includes the following three stages.
Initiation. At this stage, the covalent bond in the chlorine molecule is broken, resulting in the formation of two atoms, each with an unpaired electron:
A reaction of this type is homolysis, or hemolytic division (see Section 17.3). It is also an example of photolysis. The term photolysis means photochemical decomposition. The two chlorine atoms formed are intermediates. They are radicals. A radical is an atom (or group of atoms) that has at least one unpaired electron. It should be noted that although the initiation stage is the slowest stage of the chain reaction, it does not determine the speed of the entire chain reaction.
Stage of development. At this stage, chlorine atoms react with hydrogen molecules, forming the final product - hydrogen chloride, as well as hydrogen radicals. Hydrogen radicals react with chlorine molecules; as a result, new portions of the product and new chlorine radicals are formed:
These two reactions, which together make up the developmental stage, are repeated millions of times.
Circuit break stage. The chain reaction finally stops as a result
reactions such as
To absorb the energy that is released during these chain termination reactions, it is necessary for some third body to take part in them. This third body is usually the walls of the vessel in which the reaction is carried out.
Quantum yield
The absorption of one photon of light by a chlorine molecule in the chain reaction described above can result in the formation of millions of hydrogen chloride molecules. The ratio of the number of product molecules to the number of light quanta (photons) initiating the reaction is called the quantum yield. The quantum yield of photochemical reactions can range from one to several millions. A high quantum yield indicates the chain nature of the reaction occurring.
Pulse photolysis
This is the name of the technique used to obtain radicals at a concentration high enough to detect them. In Fig. Figure 9.14 shows a simplified diagram of the setup used for flash photolysis. The reaction mixture is affected
Rice. 9.14. Pulsed photolysis.
with a powerful flash of light from a special pulsed source. Such a source makes it possible to create flashes of light with an energy of up to 105 J and with a duration of the order of s or less. Modern methods of pulsed photolysis use pulsed lasers with a flash duration of the order of a nanosecond (10-9 s). The reaction occurring as a result of such a flash of light can be monitored by recording a sequence of optical absorption spectra of the reaction mixture. The first flash is followed by a series of flashes from a low-power pulsed source. These flashes follow each other at intervals of the order of milliseconds or microseconds and make it possible to record the absorption spectra of the reaction mixture at such time intervals.
Combustion
The reaction with oxygen, resulting in the release of heat energy and light, is called combustion. Combustion usually occurs as a complex sequence of radical reactions.
Let's take hydrogen combustion as an example. Under certain conditions, this reaction occurs explosively. In Fig. Figure 9.15 presents experimental data for the reaction of a stoichiometric mixture of hydrogen and oxygen in a Pyrex reactor. The shaded area of the diagram corresponds to the explosive region of this reaction. For the hydrogen combustion reaction, this section of the diagram has the shape of an explosive peninsula. The explosion area is limited by the boundaries of the explosion.
Rice. 9.15. Conditions for the explosive occurrence of the hydrogen combustion reaction:
Academician N. SEMENOV.
Academician Nikolai Nikolaevich Semenov. He planted this lilac bush near the building of the Institute of Chemical Physics of the Russian Academy of Sciences himself.
Rice. 1. The rate of most chemical reactions W changes rapidly with temperature. The left scale of the ordinate axis refers to curve 1, and the right scale refers to curve 2, which is a continuation of curve 1.
Rice. 2. Chemical process occurring between H molecules 2 and about 2, can be likened to a sled standing on the top of a mountain, the profile of which is shown in the figure.
Rice. 3. The chain reaction can be compared to going down a mountain, the profile of which is shown in the figure, where there is a sled in each depression.
Science and life // Illustrations
Rice. 4. One primary center can cause an entire avalanche of chemical transformations. Two types of such avalanches are depicted, where each line represents one elementary reaction act.
The Nauka publishing house, with the support of the Russian Foundation for Basic Research, is completing the publication of selected works of N. N. Semenov in four volumes. The publication includes major works, starting with student publications completed in 1913.
In April 2006, the domestic and world scientific community celebrates 110 years since the birth of the great natural scientist of the 20th century - Academician Nikolai Nikolaevich Semenov, the first and so far only Russian scientist to receive the Nobel Prize for his work in chemistry.
He was interested in chemistry since childhood, carried out experiments that sometimes ended in explosions, read textbooks voraciously, and looked for answers to questions that arose. In the article “About Time and About Myself” (see “Science and Life” No. 6, 1966), Nikolai Nikolaevich recalls the following episode: “I could not understand why, for example, ordinary salt, consisting of the soft metal sodium and poisonous gas chlorine, is so different in properties from the components of which it consists. With a childish desire to check everything for myself, I burned a piece of sodium in chlorine at home and, having obtained a precipitate, salted a piece of bread with it and ate it. You can’t say anything: it was really salt!"
Even in his youth, he came to the conclusion that to understand chemistry you need to know physics, and in 1913 he entered the physics department of the Faculty of Physics and Mathematics of St. Petersburg University. The physical approach to chemical reactions turned out to be unusually fruitful: a new science was born - chemical physics, which examined chemical processes based on physical concepts of the structure of matter.
Nikolai Nikolaevich Semenov was for many years a member of the editorial board of the journal Science and Life and the author of many remarkable articles. On the pages of the magazine, he talked about his teachers and colleagues, about the joys and difficulties of searching for scientific truth, about the ways of development of science, about new directions in chemistry and, of course, about the theory of chain reactions, which brought him world fame, and in 1956, the Nobel Prize , together with the British chemist S. Hinshelwood. The mechanism of chain reactions has become the key to understanding many different phenomena - combustion, explosion, biochemical processes. In 1940, Academician N. N. Semenov published an article “Combustion Theory” in the journal “Science and Life”. We bring to the attention of readers the section of this article devoted to chain reactions.
The weak development of chemical kinetics and the unusually bright thermal and hydrodynamic effects of flames and explosions forced previous researchers to fix their attention precisely on the thermal and hydrodynamic side of combustion, obscuring questions about the rate of chemical transformation underlying the phenomenon itself. This was their mistake and the reason for the failures in all theoretical constructions about the nature of the flame (excluding the theory of steady-state detonation). The scientific school of the Institute of Chemical Physics posed the question differently from the very beginning. Since the root cause of the thermal and hydrodynamic phenomena of combustion lies in the chemical transformation itself, the kinetics of chemical transformation has become the main link of the issue for us. This is where we directed the main attack.
Considering, however, that the powerful thermal and hydrodynamic effects of the reaction have a strong inverse effect on the rate of chemical transformation, we directed an auxiliary effort to solve the problems of hydrodynamics and heat transfer in flames in their close interaction with kinetics. This interaction of several types of weapons has led us to significant progress towards creating the theory of combustion and explosions.
The rate of most chemical reactions changes rapidly with temperature (Figure 1).
In these simplest cases, as the temperature of the combustible gas increases, the slow reaction, accelerating, leads to self-ignition when a certain critical temperature is reached. The point comes down to the fact that when a certain (as can be shown, small) reaction rate is reached, the heat generated by it does not have time to be removed through the gas and the walls of the vessel to the outside. This causes progressive heating of the gas, which in turn leads to an even greater acceleration of the reaction, etc. As a result of such a thermal avalanche, a violent combustion process occurs, ending within a split second and perceived by us as an explosion. All this is completely similar to the spontaneous combustion of undried haystacks or sulfur in dumps. This interpretation of self-ignition was very briefly qualitatively formulated by Van't Hoff in 1883 and quantitatively developed by me in 1928 and tested experimentally. Employees of our institute Todes and Frank-Kamenetsky have detailed and refined this theory in recent years.
As a result of all this work, the self-ignition temperature, as a constant of matter, was completely discredited. It turned out to be a derived quantity from the constants that determine the rate of chemical transformation and from the conditions of heat transfer (size of the vessel, thermal conductivity of the mixture, etc.).
We have shown that if the kinetics of a chemical reaction is known, the auto-ignition temperature can be predicted with great accuracy.
However, such a simple picture is observed only for a few reactions, especially for those where the reaction is reduced to the simple disintegration of molecules into parts.
In the case of an oxidation reaction (and most technically interesting reactions belong to this class), the kinetics turns out to be much more complex and leads to new remarkable phenomena in the field of spontaneous ignition. A large number of new facts that we have discovered over the past 12 years, as well as an analysis of old, long-forgotten works, led us to the formulation of the chain theory of chemical reactions, set out in my book, published in 1934. I am forced to touch upon this old material here, since our new works are closely related to this theory.
The direct connection of fuel and oxygen molecules (for example, hydrogen H 2 + O 2) is a very difficult process, because although water, for example, is thermodynamically incomparably more stable than H 2 and O 2, the relative stability of H 2 and O 2 is still also very large. The chemical process occurring between them can be likened to a sled standing on top of the mountain profile shown in Fig. 2.
The position of the sled at the bottom of the mountain is much more stable, but in order to slide down, the sled has to be pulled onto the hill, having previously expended energy.
Therefore, the reaction usually chooses a different path, which leads to a high conversion rate.
It is known that free atoms, radicals and some unstable intermediate compounds react with molecules much more easily than molecules react with each other. In such a reaction, along with the product molecule, a new radical is usually formed, which in turn reacts with the molecule, etc. In this case, one primary radical creates a long chain of subsequent reactions. Using our analogy, we could compare this With mountain profile shown in Fig. 3, where there is a slide in each cavity.
After we have dragged the first of them onto the hill and let it down, they will hit the second and push them, the second will push the third, etc. Naturally, such a process turns out to be much more economical than if we drag each sled onto its own hill and go down.
Under certain conditions, during separate elementary reactions, two radicals appear at once, which leads to chain branching. In this case, one primary center can cause an entire avalanche of chemical transformations (see Fig. 4, which schematically depicts two types of such avalanches, where each line represents one elementary reaction act).
Under some external conditions of pressure, temperature, etc., this avalanche can develop, but under others it cannot. There are no conditions yet for the development of a chain avalanche, and with the very rare appearance of the initial radicals, the reaction practically does not occur at all. As soon as the conditions for the development of an avalanche are created, no matter how small the number of initial centers is created, the reaction, accelerating like an avalanche, will lead to complete combustion of the substance.
Formally, this picture is extremely similar to the proliferation of bacteria, in particular those that cause disease in the body. With an insignificant number of bacteria entering the body during infection, if the conditions for their reproduction are favorable, an avalanche of bacterial reproduction leads to disease. The active centers of the valuable theory are the same bacteria of the chemical process, the reproduction of which ends in the death of the original substance. Here, as with infection, the principle reigns - all or nothing. Either, given the appropriate state of the body, bacteria practically do not multiply, or they multiply in huge quantities. Either the reaction is very small or it occurs at a high rate.
Impurities that break chains and inhibit the process are similar to serums that kill or stimulate the body to kill bacteria.
If a chemical avalanche grows slowly, then we are dealing with an auto-accelerating process, leading to an explosion with a large delay - at the moment when the reaction rate reaches such a value that the heat it generates no longer has time to be removed through thermal conductivity. If a chain avalanche develops quickly, it leads to the phenomenon of self-ignition and burnout of the substance, even completely independently of thermal phenomena. The ignition of phosphorus, phosphine, and carbon disulfide vapors at a concentration of about hundredths of a percent in the air does not cause virtually any increase in temperature. However, ignition occurs under strictly defined conditions. This is a typical implementation of a chain isothermal avalanche in its pure form. However, even when rich mixtures are ignited, the nature of the process leading to an explosion is the same. Ignition occurs with the help of a chain avalanche; it is the primary cause, and violent heating and sound are secondary phenomena here.
An explosion of the first kind, when its primary cause is a thermal avalanche, occurs in such a way that at a temperature just below the explosive one, a small but still quite measurable reaction occurs. An explosion of the second kind, when its root cause is a chain avalanche, is distinguished by the fact that at a temperature slightly below the explosive one, the reaction may practically not occur at all.
As an example, we present a diagram of the hydrogen oxidation reaction (Fig. 5).
Already from this diagram it is clear that impurities in minute quantities can greatly inhibit reactions by combining with H, O or OH atoms and thereby breaking the chain.
Although the patterns of chain reactions can be very diverse, it has been possible to establish a number of general laws of chain reactions and explain and predict a number of surprising facts. Of the huge number of those discovered here and abroad, I will demonstrate here only one.
We are accustomed to thinking that the greater the pressure of a combustible mixture, the more easily it ignites and burns. In many chain avalanche cases this is not the case. Not only is there no ignition, but there are no traces of a reaction at high pressure. When the pressure decreases below a certain critical value, ignition occurs.
I have often been reproached that we talk very easily about radicals and intermediate products without establishing their presence in chain reactions. It seems to me that the experiments of the Institute of Chemical Physics over the past two years free us from this reproach. Prof. Kondratyev and his co-workers showed that in hydrogen flames at low pressures (several millimeters of mercury), where the flame temperature can optionally vary from 600 to 800°, relatively very large concentrations of OH radicals are present, reaching 0.1 mm of mercury, t i.e. several percent of the original mixture. The registration of radicals was carried out using the absorption spectra method. A flame was placed in the path of a beam of light emitted by a discharge tube filled with water vapor (such a tube emits OH lines). Passing through the flame, a beam of light fell on the slit of the spectrograph. As a result of the absorption of light by OH radicals, the intensity of the emission line weakened. From the drop in intensity, it was possible to calculate the OH concentration in the flame. It is interesting to note that the OH concentration is hundreds of thousands of times higher than its thermodynamic equilibrium values at Kondratieff flame temperatures. This proves that OH appears as a result of a chemical avalanche, and not as a result of thermal dissociation. Using a similar method, Kondratiev and his co-workers proved the presence of large quantities of the CS radical and SO molecules in cold carbon disulfide flames.
By studying the self-accelerating slow reaction of hydrogen sulfide oxidation H 2 S, Emanuel, Pavlov and I showed this year that the products of this long-known reaction are not only SO 2 (sulfur dioxide) and H 2 O, but also such “exotic " product like SO. In the first stages of the reaction, SO appears in very large quantities, reaching 7% of the initial substance and up to 40% of the H 2 S converted to this point, and only in the final stages does SO disappear, turning into SO 2. SO was recorded from absorption spectra during the course of the reaction, as well as by a second, new method, the details of which I cannot go into here. Thus, SO is a typical intermediate product. It can be shown that it is its formation that causes the reaction to auto-accelerate and facilitates the explosion.
Similar experiments are carried out by Neumann, Sokolik and co-workers on intermediate products of the oxidation of hydrocarbons and ethers. All these experiments lay the foundation for the chemistry of intermediate compounds, which should become the experimental basis for a new branch of chemistry - process chemistry, or chemical kinetics, just as experiments with a microscope became the basis for the development of bacteriology.
We believe that we have completed the theory of self-ignition in its main features. It has firmly entered the world scientific literature.
CHAPTER 22. CHAIN REACTIONS
22.1. Basic Concepts of Chain Reactions
In some cases, in chemical reactions, active particles such as free atoms and radicals, which have free valencies and therefore have high reactivity, act as intermediate products. These particles enter into reactions, as a result of which free atoms and radicals re-emerge. This sequence of periodically repeating reactions involving active particles (free atoms and radicals) is called chain reaction.
Although the process of formation of free atoms or radicals requires high activation energy, their high reactivity and the emergence of new active particles in reactions with saturated molecules lead to the fact that the rate of chain reactions is usually much higher than the rate of non-chain reactions. The nucleation of a small number of reactive particles at the beginning of the reaction leads to the conversion of a large number of starting substances. Since chain reactions occur cyclically, the active particle that appears at the end of the cycle gives rise to a new cycle, at the end of which the active particle is regenerated again.
Chain reactions include reactions of various classes. For example, combustion or slow oxidation reactions in the gas phase proceed through a chain mechanism:
2H 2 + O 2 2H 2 O
CH 4 + 2O 2 CO 2 + 2H 2 O
Chain reactions include many reactions involving hydrocarbons (polymerization reactions, decomposition reactions), photochemical reactions (formation of HCl, HBr, COCl 2, etc.), nuclear chain reactions - the decay of uranium-235 or plutonium in a nuclear reactor or bomb.
A characteristic feature of chain reactions is the great sensitivity of the rate of these reactions to the presence of certain impurities. For example, thoroughly dried hydrogen and oxygen react with each other very slowly, but the reaction proceeds at a normal rate in the presence of a small amount of water vapor. A mixture of hydrogen and chlorine does not react in the dark at room temperature, but reacts quickly when small amounts of sodium vapor are introduced into the system. In other cases, the presence of impurities leads to a sharp decrease in the reaction rate. For example, when photochemically initiating the reaction of hydrogen with chlorine, the rate of formation of hydrogen chloride decreases by about a thousand times in the presence of one percent oxygen.
The rate of many gas reactions is affected by the shape and material of the vessel in which the reaction occurs. Typically reactions slow down as the ratio increases S/V(S– surface area of the vessel, V– its volume). This ratio can practically be changed by introducing fragments of the vessel material - glass, quartz, etc. - into the vessel.
Many oxidation reactions in the gas phase are characterized by the fact that a fast reaction (spontaneous ignition) occurs only within certain limits of pressure and temperature. Figure 22.1 shows the dependence of the ignition limits on pressure and temperature, which is observed during the oxidation of hydrogen, phosphorus vapor, carbon disulfide, etc.
Rice. 22.1. Flammability limits for hydrogen oxidation reaction
Ignition of the mixture occurs only under conditions corresponding to the shaded area in the figure, which is called peninsula ignition. Outside the peninsula, ignition does not occur and the reaction proceeds at a low speed or practically does not occur at all. Based on the point A, ignition can be caused by heating the mixture or reducing the pressure of the mixture to values lying in the area between curves II and I.An explanation for these features is provided by the theory of chain reactions, the development of which dates back to 1913, when Bodenstein introduced the concept of a chain reaction.
There are two types of chain reactions: with unbranched And branched chains. An example of the first type of reaction is the reaction of the formation of hydrogen chloride from hydrogen and chlorine
H 2 + Cl 2 2HCl,
the mechanism diagram of which was proposed by Nernst.
Three groups of reactions can be distinguished in the scheme. The process begins with chain nucleation reactions:
Cl 2 Cl + Cl
This reaction of dissociation of a chlorine molecule into atoms can occur when light is absorbed
Cl2+ h Cl + Cl,
thermally - in the collision, for example, of two chlorine molecules with increased energy:
Cl 2 + Cl 2 Cl + Cl + Cl 2,
chemically - for example, during the interaction of a chlorine molecule with a sodium atom, the vapors of which are introduced into the system. The resulting chlorine atoms are highly reactive and enter into further interaction with the starting substances; the second group of reactions occurs - chain development:
Cl + H 2 HCl + H
H + Cl 2 HCl + Cl
As a result of the first reaction, a hydrogen atom appears, which easily interacts with a chlorine molecule, as a result of which hydrogen chloride is formed and the chlorine atom is regenerated, which gives rise to the next link:
Cl + H 2 HCl + H
Cl H Cl H Cl ...
Under favorable conditions, such a chain can consist of many thousands of links. As a result, for one initially activated chlorine molecule, not two HCl molecules are formed, as in a conventional bimolecular reaction, but thousands and tens of thousands of molecules.
For the above reaction, it is characteristic that for each active particle of Cl or H that enters the reaction, one active particle is again formed. Such chains are called unbranched.
In addition to the above reactions of chain nucleation and development, a third group of reactions occurs in the system - open circuit reactions, leading to the death of active particles upon collision with any third particle M or the wall of the vessel S:
N + N + M H 2 + M
H + H + S H 2 + S
Cl + Cl + M(S) Cl 2 + M(S)
H + Cl + M(S) HCl + M(S)
In the presence of, for example, oxygen, chain termination can occur as a result of the reaction
H + O 2 + M 
+ M
The resulting low-active radical 
dies on the walls of the vessel or by reaction
+ HH 2 + O 2
At low pressures, active centers die mainly on the walls of the vessel, and at high pressures, trimolecular termination occurs in the volume. Therefore, chain reactions are characterized by the features mentioned above - the dependence of the reaction rate on the specific surface area of the vessel, on the presence of any inert substance, on the pressure or concentration of the reacting substances.
The kinetic equation of an unbranched chain reaction can be obtained based on the reaction mechanism. For example, a detailed study of the reaction between hydrogen and bromine
H 2 +Br 2 2HBr
showed that the reaction occurs in several elementary stages with different rate constants k:
Br 2 Br+Br k 1
Br+H 2 HBr+H k 2
H+Br 2 HBr+Br k 3
H+HBrH 2 +Br k 4
Br+BrBr 2 k 5
Based on this scheme, the rate of formation of hydrogen bromide can be represented by the equation:
+
–
. (22.1)
Considering bromine and hydrogen atoms as intermediate products, we can apply Bodenstein’s principle of stationary concentrations to them (see Section 20.6):
From the sum of these equilibria we find the concentration of bromine atoms:
. (22.4)
After substituting this concentration into equation (22.3), we obtain the concentration of hydrogen atoms:
. (22.5)
Substituting the concentrations of bromine and hydrogen atoms into equation (22.1) gives the final equation for the reaction rate:
. (22.6)
This equation coincides with equation (20.6), obtained from experimental data.
In a number of reactions, as a result of one elementary act, not one, but two or more chemically active particles can arise, i.e. chain branching occurs. Such reactions are called branched chain reactions. In such reactions, in the initial period of time, the number of active particles, and therefore the reaction rate, increases like an avalanche until the moment when, due to the consumption of the starting substance, the reaction rate begins to decrease. An example of such processes is the reaction of hydrogen oxidation, the mechanism of which, according to modern concepts, can be represented as a set of sequentially occurring elementary chemical acts:
Chain initiation
(4)
+ H 2 H 2 O + H Continuation of the chain
Chain branching
Broken circuit on the wall
(9) H + O 2 + M 
+ M Open circuit in volume
Low-active radicals formed 
may disintegrate on the wall:
2
+ SH 2 O 2 + O 2 + S
At high pressures, reactions in volume are possible:
(10)
+ H 2 H 2 O 2 + H Continuation of the chain through
(11)
+ H 2 OH 2 O 2 + 
low-active radical
If chain branching occurs frequently, then even one initially formed chain can lead to the development of many chains. In the extreme case, one can imagine that branching occurs in each link, and then one speaks of a completely branched chain reaction. In other cases, branching may occur more rarely.
The existence of lower and upper flammability limits can be qualitatively explained as follows. At pressures below the lower limit, active particles easily diffuse to the walls of the vessel, where they die. Breakage of chains on the walls prevails over branching, and a fast reaction does not develop. As the pressure increases, diffusion to the walls becomes more difficult, and the number of double collisions of types (5) and (6) increases, which lead to chain branching; chain nucleation and branching begin to dominate over termination. As a result, the reaction self-accelerates and can end in spontaneous combustion or explosion, which occurs inside the ignition peninsula.
Rice. 22.2. Dependence of the speed of a branched chain reaction on time inside the ignition peninsula
With a further increase in pressure, triple collisions in the volume become more and more likely, leading to circuit breakage. If the pressure exceeds the value of the upper limit II (Figure 22.1), the termination begins to dominate the development of chains and the possibility of a rapid reaction disappears.The ignition of the combustible mixture inside the ignition peninsula is preceded by an induction period t ind (Fig. 22.2). It is explained by the fact that at first the number of chains can be very small and the reaction is practically invisible due to the insufficient sensitivity of the analysis methods. But after some time t Ind the number of chains increases very quickly due to their multiplication and self-ignition or explosion occurs. Reaction rate dependence v from time t can be represented by the equation:
, (22.7)
Where A And – constant for a given reaction and depending on a number of conditions. When deriving this dependence, the decrease in the concentration of reacting substances due to burnout was not taken into account, so the tendency of the speed to infinity over time has no physical meaning - the speed becomes high, but not infinite.
In some cases, a third ignition limit is also observed (Fig. 22.1), which lies at higher pressures. Its existence is associated with the occurrence of chain reactions due to low-active radicals or with the development of a thermal explosion.
22.2. Elementary theory of chain reactions
There are two versions of the theory of chain reactions - a more rigorous one, based on solving a system of differential equations, and a less strict, but more visual probabilistic version, which is discussed below.
An important characteristic of a chain reaction is the average chain length - the average number of elementary reactions caused by one active particle (atom or radical), which initially arose in some independent way. If n o is the number of such independently arising particles per unit time in a unit volume, then n o can be called rate of chain nucleation.
The reciprocal of the average chain length is probability of circuit breakage. This relationship can be understood using the circuit diagram in Figure 22.3. In the diagram, a dot means the appearance and regeneration of an active particle, and a cross means its death, i.e. circuit break.
Rice. 22.3. Schematic representation of the chain reaction:
A– unbranched chain; b– branched chain
There are only cases per favorable case - a cliff, therefore, =1/ . (22.8)
Let us also assume the possibility of chain branching - the appearance of two or more active particles in any link and characterize this possibility probability of chain branching .
Let us denote by the time during which, on average, one link of a chain reaction occurs. Then the product is equal to the average time of passage of the entire chain from the moment of initiation to breakage. The concentration of active particles, i.e. their number per unit volume, let it be n. The rate of change in the concentration of these particles will be equal to the difference in the rates of their formation n o and disappearance.
If the chain length = 1 (i.e., there is actually no chain), then the active particle dies in each link. Then for the average development time of one link everyone will react n particles, and the disappearance rate will be equal to n/ particles/cm 3 With. If the chains develop and their average length is equal to > 1, the particle will react on average times, and the average time of its life will be equal . Consequently, the rate of decrease in particle concentration will be expressed by the relation
. (22.9)
If branching of the chain is possible, i.e. > 0, then its influence can be taken into account, considering that the branching acts as if in the direction opposite to the break, lengthening the chain and reducing the probability of a break to the value ( – ). Then for the rate of change in the concentration of active particles we can write the expression:
. (22.10)
This differential equation can be solved as follows. For simplicity of notation, we introduce the notation a= ( – )/ . Then
. (22.11)
We assume at first that n o = 0, and after separating the variables we get:
, (22.12)
integration of which gives:
ln n = –at + ln Z(t), (22.13)
Where Z(t) – some conditional “constant” of integration. Then
n=Z(t)e –at . (22.14)
Let us differentiate this equation taking into account the fact that Z is not a constant:
From a comparison of this equation with equation (22.11) it follows that
(22.16)
. (22.17)
After integrating this equation we get
, (22.18)
Where I– integration constant. Substituting this quantity into equation (22.15) gives
. (22.19)
From the condition that at the initial moment of the reaction ( t= 0) value n= 0, it follows that
(22.20)
. (22.21)
After substituting the value a we get
. (22.22)
Speed reaction v can be defined as the rate of increase in the concentration of molecules of the reaction product. Since in one link per time one molecule appears, then the total number of molecules formed per unit volume per unit time is equal to n/ . Thus, we obtain the basic equation of the theory of chain reactions:
. (22.23)
Let's consider the use of this equation for some special cases.
When an unbranched chain reaction occurs = 0. Since the average chain length = 1/ , then the rate of such a reaction
. (22.24)
Rice. 22.4. Dependence of the speed of the chain reaction on time:
1 – = 0; 2 – 0 < < ; 3 – >
As follows from this equation, the reaction rate must increase over time and reach a limit equal to n o = n o/ (Fig. 22.4), i.e. the system must reach a steady state in which the rate of reaction is constant. This speed in times the rate of nucleation of primary reactive particles n o, i.e. reaction rates in the absence of chains ( =1).If a branched chain reaction is possible, the probability of branching may be less than the probability of termination, i.e. 0< < . In this case, according to equation (22.23), the system must also reach a stationary speed, but this speed is greater than in the first case:
If the probability of branching is greater than the probability of breakage, i.e. > , equation (22.23) takes the form:
, (22.25)
Where A andare positive constants. The resulting equation coincides with the previously given equation (22.7). The equation shows that the reaction rate can become infinitely large (Fig. 22.4), i.e. indicates the possibility of a chain fire or explosion.
It should, however, be noted that the obtained dependences refer to some idealized reaction conditions - it is assumed that the concentrations of the starting substances are maintained constant, and the reaction products are removed from the reaction zone. Under real conditions, for example, when carrying out a reaction in a closed vessel, the starting substances “burn out”, and the products remain in the reaction mixture. Therefore, for unbranched chain reactions or for branched reactions with < the speed passes through a maximum (dotted line in Fig. 22.4). It is possible that the steady state will not be reached at all, since the maximum speed may be less than the stationary one. In the event of a branched chain reaction with > taking into account the burnout of reagents should give, as already indicated earlier, a very high, but still finite, rate.
Condition > corresponds to the reaction occurring in the region of the ignition peninsula, and the condition > - outside of it. Thus, the theory of branched chain reactions quantitatively explains the existence of lower and upper flammability limits.
Self-heating of the reacting mixture can also lead to ignition or explosion, regardless of the reaction mechanism. According to the Arrhenius equation, the reaction rate increases exponentially with increasing temperature, while the rate of heat removal increases more slowly (proportional to the temperature difference). In the case of an exothermic reaction, if heat is not removed from the reaction zone at a sufficient rate, the reaction mixture will begin to self-heat and the reaction rate will increase more and more. The development of these processes can lead to ignition of the reaction mixture or explosion. In this case they talk about thermal self-ignition(thermal explosion). The kinetics of thermal self-ignition may not differ in appearance from the kinetics of chain ignition, which must be kept in mind when studying reactions that lead to ignition or explosion.
The content of the article
CHAIN REACTIONS– chemical reactions that occur through a sequence of the same elementary stages, at each of which one or more active particles (atoms, free radicals, ions, radical ions) appear. Cracking, combustion, polymerization and a number of other reactions proceed through a chain mechanism
Bodenstein–Nernst chains.
By the end of the 19th century. The most important chapter of physical chemistry was developed - the study of equilibria of chemical reactions (chemical thermodynamics). It has become possible to calculate to what maximum possible depth a specific reaction can proceed under given conditions. At the same time, the doctrine of the rates of chemical processes - chemical kinetics - was created. Accumulated by the second half of the 19th century. numerous experimental data could be explained based on the law of mass action and the Arrhenius equation. At the same time, facts appeared that could not be explained by any of the existing theories. One of the most mysterious was the seemingly simple reaction of hydrogen with chlorine: H 2 + Cl 2 → 2HCl.
In 1845, the English chemist John Draper discovered that when exposed to sunlight, chlorine becomes particularly active in its reaction with hydrogen ( cm. PHOTOCHEMISTRY). An even more surprising fact was discovered in 1857 by the German chemist Robert Bunsen and his student from England, Henry Roscoe. It turned out that some impurities, even in the smallest concentrations, can have a huge impact on the rate of this reaction. For example, small additions of oxygen slowed it down hundreds of times. This was a paradoxical result, since oxygen itself reacts perfectly with hydrogen. Other strange phenomena were also discovered. For example, the reaction rate depended on the material of the vessel wall and even on its size. A gap appeared in the seemingly harmonious doctrine of reaction rates, and no one knew how to deal with it.
And the reaction of hydrogen with chlorine presented scientists with new surprises. At the beginning of the 20th century. Albert Einstein formulated a law according to which each absorbed quantum of light (photon) causes changes in only one molecule. It is easy to experimentally measure the number of reacted (or formed) molecules and the number of light quanta absorbed in the reaction. The ratio of these quantities is called the quantum yield of the reaction. Thus, if for each quantum of light absorbed by the reagents one molecule of the product is formed, then the quantum yield of such a reaction is equal to unity. However, the experimentally measured quantum yields of many reactions did not correspond to the law of quantum equivalence. In 1913, one of the founders of chemical kinetics, German chemist Max Bodenstein, measured the quantum yield of the photochemical reaction of hydrogen with chlorine H 2 + Cl 2 2HCl. The result turned out to be incredible: the number of HCl molecules formed when the mixture absorbed one quantum of light, under some conditions reached a million! Bodenstein explained this amazing result with the only reasonable method: each absorbed quantum of light “triggers” a long chain of transformations in which hundreds of thousands of molecules of the starting substances (H 2 and Cl 2) react, turning into molecules of the reaction product (HCl). This is similar to how dominoes lined up in a row quickly, as if on cue, fall one after another if you successfully push the first one.
Bodenstein also formulated the basic principles of a new type of chemical transformation - chain reactions. These reactions necessarily have three stages: 1) chain nucleation, when active particles are formed; 2) continuation (development) of the chain; 3) open circuit. The nucleation of chains in a thermal reaction occurs as a result of the dissociation of molecules upon heating. In a photochemical reaction, chain nucleation occurs upon absorption of a light quantum. At the stage of chain continuation, molecules of reaction products are formed and at the same time a new active particle appears that is capable of continuing the chain. At the termination stage, the active particle disappears (deactivates).
When exposed to high heat or intense exposure to ultraviolet light, the chain reaction of hydrogen with chlorine occurs explosively. But if the temperature is not very high or the light intensity is low, the reaction proceeds calmly. Based on this fact, Bodenstein put forward a very important principle of the stationary concentration of intermediate products of chain reactions. In accordance with this principle, the rate of generation of active particles at the generation stage is equal to the rate of their disappearance at the termination stage. Indeed, if the rate of termination were greater than the rate of chain nucleation, the number of active particles would drop to zero, and the reaction would stop by itself. If the nucleation rate prevailed, the number of active particles would increase with time, which would lead to an explosion.
However, elucidating the chemical mechanism for each stage of the reaction of hydrogen with chlorine has proven difficult. Bodenstein proposed the theory of energy branching: the HCl molecules formed in the primary reaction carry excess energy and therefore contribute to the occurrence of further reactions by transferring excess energy to the molecules of the starting substances. However, this theory turned out to be incorrect in this case. The correct reaction mechanism was given in 1918 by the German physical chemist Nobel Prize winner Walter Nernst. He suggested that the active particles were hydrogen and chlorine atoms; the chain reaction diagram looked like this. The nucleation of a chain occurs during the thermal dissociation of chlorine molecules at high temperatures or when they absorb light quanta at room temperature: Cl 2 → 2Cl. This is followed by two rapidly repeating stages of chain continuation: Cl + H 2 → HCl + H and H + Cl 2 → HCl + Cl. Chain termination occurs when active hydrogen or chlorine atoms react with impurity molecules, or “stick” to the wall of the vessel, or react (recombine) with each other, turning into inactive H 2 and Cl 2 molecules.
It was subsequently shown that hydrogen atoms are much more active than chlorine atoms; Accordingly, hydrogen atoms react much faster and therefore their stationary concentration is much lower. Thus, at room temperature, the steady-state concentration of hydrogen atoms is approximately 100 times less than that of chlorine atoms. As a result, the probability of meeting two hydrogen atoms or hydrogen atoms and chlorine atoms is much less than for two chlorine atoms, so practically the only chain termination reaction is the recombination of chlorine atoms: Cl + Cl → Cl 2. If the pressure in the reaction vessel is very low and its dimensions are small, the active particles can reach the wall of the vessel even before they react with H 2 and Cl 2 molecules; under these conditions, the breaking of chains on the walls of the reaction vessel can play an important role.
Nernst's scheme was confirmed by various experiments. One of the most ingenious ones was carried out by the English physical chemist Michael Polyani. In his experiments, a stream of hydrogen passed over slightly heated sodium metal and carried away with it a very small amount of its vapor. Then the stream fell into a vessel with chlorine in the dark. At the temperature of the experiment, pure hydrogen did not react with chlorine, but an insignificant admixture of sodium vapor completely changed the matter: a rapid reaction of the formation of hydrogen chloride took place. Here, instead of light, sodium plays the role of initiator of the chain reaction: Na + Cl 2 → NaCl + Cl. Just as in a photochemical reaction there are many reacted molecules for each absorbed quantum of light, so here for each reacted sodium atom there are many HCl molecules formed. Polyani obtained similar results for the reaction of chlorine with methane. In this case, the chain initiation and termination reactions were the same as in the reaction of chlorine with hydrogen, and the chain continuation reactions looked like this: Cl + CH 4 → HCl + CH 3 and CH 3 + Cl 2 → CH 4 + Cl . These reactions also involve particles with unpaired electrons (indicated by a dot) - free radicals.
Many reactions involving free radicals turned out to be chain reactions, the mechanism of which was generally similar to the mechanism of the reaction of hydrogen with chlorine. The splitting reactions at high temperatures (pyrolysis) of hydrocarbons, for example, ethane, proceed by a chain mechanism: C 2 H 6 C 2 H 4 + H 2 ; Such reactions are of great importance in the industrial processing of petroleum hydrocarbons. Chain reactions turned out to be the oxidation of organic compounds with oxygen, reactions of addition to unsaturated compounds of halogens (chlorine and bromine), hydrogen bromide and other compounds, polymerization reactions, and a number of other processes. Polymerization chain reactions are interesting because in them the chain continuation steps leave behind “real chains” in the form of monomer unit residues linked to each other. In a thickened and hardened polymer (for example, in polystyrene or in polymethyl methacrylate - “organic glass”), terminal free radicals can sometimes be found, which, due to their high viscosity, were unable to react with the free monomer molecule.
Semenov-Hinshelwood chains.
At the end of 1924, at the Leningrad Institute of Physics and Technology, in the Laboratory of Electronic Chemistry, headed by N.N. Semenov, they began to measure the intensity of the glow of phosphorus vapors during their oxidation with oxygen. In their first experiments, a young university graduate, Zinaida Valta, and her immediate supervisor, Yu.B. Khariton, came across a completely unexpected phenomenon. It turned out that when there is little oxygen, the oxidation of phosphorus does not occur at all. But as soon as the oxygen pressure exceeded a certain critical value, intense oxidation began with the emission of light. Previously, the theory assumed that the reaction rate should increase smoothly with increasing concentration. Here there is a sharp transition from a complete lack of reaction to a very fast process with an insignificant change in pressure. Another, quite strange fact was revealed: at pressure below critical, i.e. in the absence of a reaction, it was enough to introduce argon into the vessel to cause a bright flash. It turned out that the inert gas argon, incapable of any chemical reactions, made oxygen reactive! This was already a real miracle...
Later it turned out that oxygen can completely lose its activity not only when the pressure decreases, but also when the pressure increases above a certain critical value. This second (upper) limit of oxygen pressure depended significantly on the impurities of various substances. Some of these impurities made “passive” oxygen very active, causing phosphorus to burn. This behavior contradicted all then existing ideas about the mechanisms and rates of chemical reactions.
The results of strange experiments, without any attempt to explain them, were published in the German Physical Journal. The consequences were quick and disappointing: the work was subjected to extremely sharp criticism from the famous Bodenstein, who by that time was considered the head of world chemical kinetics. He wrote that all the results on the oxidation of phosphorus are not a discovery, but an illusion, and even pointed out its reason - the incorrect design of the installation in which the experiments were carried out. At the conclusion of his short article, Bodenstein noted that the so-called “limiting” phenomena were repeatedly observed in the past for different reactions, but when tested each time it turned out that they were all associated with various experimental errors.
The objections were very serious. But a thorough check (without Khariton - he was on a business trip abroad and without Valt - she moved to another institute) showed the correctness of the first publication. Moreover, new, no less “heretical” data were obtained. It turned out, for example, that the critical oxygen pressure strongly depends on the size of the reaction vessel.
Semenov felt that he was on the verge of a discovery. He understood that the reaction was a chain reaction, like the reaction of hydrogen with chlorine. However, the Bodenstein–Nernst chain reaction mechanism, based on the “domino principle,” never led (and could not lead) to critical phenomena. There was something different here. At the same time, Cyril Hinshelwood began working in this direction in England. In both laboratories, critical phenomena were discovered in the combustion reactions of hydrogen and a number of other substances. It turned out, for example, that in glass heat-resistant vessels at temperatures of 500–600 ° C, the reaction of hydrogen with oxygen does not occur at all until the pressure reaches 3–4 mm Hg. Art. When the pressure exceeded this lower limit, a rapid reaction suddenly began, accompanied by a glow. But the most amazing phenomenon was that the flames could be extinguished simply by increasing the pressure. At temperatures below 400° C, ignition in a pure mixture of hydrogen and oxygen was not observed at any pressure. However, it was enough to add an inert gas to the mixture, and a flash occurred!
All these new phenomena were explained by Semenov (and independently by Hinshelwood) under the assumption of branching chains. If in the reaction of hydrogen with chlorine at each stage of chain continuation one active particle is consumed and one appears (unbranched chain), then in the reaction of hydrogen (and other reagents) with oxygen for one disappeared active particle two or more new ones are formed, for example,
H + O 2 OH + O
O + H 2 OH + H
OH + H 2 H 2 O + H
If we add up these three consecutive reactions, we get H + O 2 + 2H 2 OH + 2H, that is, one active particle turns into three. As a result, the number of active centers rapidly increases (chains branch), and if the rate of chain termination is not high enough, the reaction very quickly goes into an explosive mode (at low pressure, instead of an explosion, a flash is observed). Such reactions, which occur with an increase in the number of active particles, are called branched-chain reactions. If we consider that these processes are highly exothermic, and the reaction of each active particle with a molecule of the starting substance requires billionths of a second, then it is easy to understand why branched chain reactions at high concentrations (pressures) of reagents cause destructive explosions.
It is important to note that the avalanche of a branched chain reaction ends very quickly: a fraction of a second after it starts, there are no longer enough starting materials to continue the reaction - almost all of them have turned into reaction products. Here we can give the following analogy: various rumors spread through a “branched chain mechanism” if everyone who learns the news tells it to more than one person. And just like rumors and gossip, various “branched-chain” financial and other pyramids (such as the famous “Vlastilina”, MMM, “chain letters”, etc.) and various “tempting” pyramids quickly end, but also spread quickly. offers to earn 100 thousand for 100 rubles and other scams that require attracting new “clients” at each stage. At first glance, everything looks fair, but an increasing number of participants quickly become involved in the pyramid, and soon there are not enough “raw materials” - there is no one else to buy the shares, and they rapidly depreciate. Similar financial pyramids date back to the 19th century. were in use in different countries; in France they were called “snowballs”, in our country they were called avalanches. Their mechanism (and mathematical description) are very reminiscent of branched chain chemical reactions.
Semenov and Hinshelwood gave this explanation for the processes studied. At low pressures, the majority of active particles - atoms, free radicals, without having time to collide with many reagent molecules and “multiply”, reach the wall of the reaction vessel and “die” on them - the chains break. The smaller the diameter of the reactor, the greater the chance for radicals to reach its walls. This is where the dependence on the size of the vessel comes from! As the concentration increases, the chances of radicals colliding with reagent molecules become greater than the chances of reaching the wall—an avalanche of reactions occurs. This explains the existence of a lower pressure limit. Inert gas molecules, as Semenov aptly put it, “get tangled up in the legs” of the active particle and slow down its movement towards the wall; This explains the amazing effect of argon on the critical pressure.
When the upper pressure limit is reached, the chains break off again faster than their branching occurs; however, the reason for chain termination here is different - active radicals disappear as a result of “mutual destruction” - recombination in the volume of the vessel (the rate of this reaction increases very quickly with increasing pressure). Thus, all experimental facts received a logical explanation within the framework of the branched chain reaction theory. In 1956, N.N. Semenov and S. Hinshelwood received the Nobel Prize in Chemistry for these studies.
The theory of branched chain reactions is of great practical importance, as it explains the behavior of many industrially important processes, such as combustion, oil cracking, and ignition of the combustible mixture in internal combustion engines. The presence of upper and lower pressure limits means that mixtures of oxygen with hydrogen, methane, and other flammable gases explode only at certain ratios. For example, mixtures of hydrogen with air explode at a hydrogen content of 4 to 75%, and mixtures of methane with air explode at a methane content of 5 to 15%. This is why gas leaks are so dangerous: if there is more than 5% methane in the air, an explosion can occur even from a tiny spark in the switch when turning on or off the lights in the kitchen.
Chain processes have acquired particular importance in connection with the work of physicists to obtain nuclear energy. It turned out that the fission of uranium, plutonium, and other fissile materials obeys the same laws as branched chain chemical reactions. Thus, the fission reaction of uranium is caused by neutrons, which split uranium nuclei, releasing enormous energy. Chain branching occurs due to the fact that when a nucleus splits, several active particles are released - neutrons, capable of splitting new nuclei.
Reactions with degenerate branching.
The oxidation of some compounds produces peroxides, which themselves are capable of decomposing under certain conditions with the formation of active particles - free radicals. As a result, chain branching occurs, although not so fast: in order for the decomposition of peroxides to occur at a noticeable rate, they must first accumulate. Such processes were called degenerate ramifications.
A typical example of a branched chain reaction with degenerate branching is the oxidation reaction of hydrocarbons. It begins with the fact that an oxygen molecule abstracts a hydrogen atom from a molecule of an organic compound: RH + O 2 R· + HO 2 ·. The hydroperoxide radical formed at the initiation stage is converted into the radical R· with an unpaired electron on the carbon atom as a result of the reaction HO 2 · + RH H 2 O 2 + R·. So the HO 2 · radical does not participate further in the reaction. The radical R· has several possibilities. Firstly, it can combine (recombine) with other radicals, including similar ones: R· + R· R 2. Secondly, it can tear off a hydrogen atom from a molecule of the original substance: R· + R"H RH + R"·. Finally, it can attach to the double bond of an oxygen molecule: R· + O=O R–O–O·. The first reaction can be ignored: the probability of meeting two active radicals is very small, since their concentration is negligible. The second reaction only results in the exchange of a hydrogen atom. But as a result of the third reaction, the peroxide radical RO 2· is formed, which, together with the radical R·, leads the chain. It consists of two repeating oxidation chain reaction steps: RO 2 · + RH ROOH + R· and R· + O 2 RO 2 ·.
It can be seen that the chain is led by the radicals RO 2 · and R ·, since they are constantly generated during the reaction. RO 2 · radicals are less active, their concentration is much higher, so the chain breaks when two peroxide radicals meet. This meeting can produce various products, including ROOR peroxides (they are formed by the recombination of peroxide radicals), alcohols, and carbonyl compounds. If the chains are long, there will be few of these substances - recombination products - and the main product of the chain reaction will be ROOH hydroperoxide, which can sometimes be obtained in high yield. The O–O bond in hydroperoxides is relatively weak (more than twice as weak as the C–O bond in alcohols). When it breaks, two radicals are formed at once - RO· and OH·, which initiate new chains. It turns out that the reaction product, hydroperoxide, simultaneously accelerates it. Such reactions are called autocatalytic.
Revival of "energy chains".
The assumption made by Bodenstein and a number of other chemists about “energy chains” did not receive experimental confirmation and was forgotten for many decades. However, in 1963, V.I. Vedeneev, A.M. Chaikin and A.E. Shilov discovered that “energy branching” is possible in the fluorination reactions of a number of compounds. An example is the reaction of fluorine with hydrogen. In this reaction, at the stage of chain continuation H + F 2 HF* + F, so much energy is released that the resulting “hot” hydrogen fluoride molecule (marked with an asterisk) can cause the chain to branch. This happens by transferring excess energy to the starting substances; The energy carrier in this case is the hydrogen molecule. The reaction mechanism is as follows:
F 2 + H 2 H + HF + F – slow stage of chain nucleation
F + H 2 HF + H – two reactions
H + F 2 HF* + F – continuation of the chain
HF* + H 2 HF + H 2 * – excitation transfer
H 2 * + F 2 H + HF* + F – chain branching
Chain breaks occur on impurity molecules or on the walls of the vessel. The study of the mechanism of this reaction made it possible to create a chemical fluorine-hydrogen laser, in which the source of light (in the infrared range) is excited HF molecules.
Ilya Leenson