KINETICS.

Kinetics is the science of the speed of processes.

Chemical kinetics considers speed and mechanism chemical reactions. The most important parameter of kinetics is the time of the process.

The reaction rates depend on many factors: the nature of the reactants, concentration, temperature, pressure, the presence of catalysts, and in the case of phase transformations, also on a number of other conditions (the state of the phase interface, conditions of heat and mass transfer, etc.). The task of kinetics is to elucidate the role of these factors and to establish the mechanism of reactions and phase transformations.

Chemical kinetics includes two sections:

1) a formal mathematical description of the reaction rate without taking into account the actual mechanism of the reaction itself (formal kinetics);

2) the doctrine of the mechanism of chemical interaction.

FORMAL KINETICS.

In formal kinetics, the rate of a chemical reaction is represented as a function of the concentration of the reactants only.

Regularities of formal kinetics allow:

1) determine the kinetic parameters of a chemical reaction (rate constant, half-life, etc.);

2) extend the patterns obtained to complex multi-stage chemical reactions typical of technological processes;

3) classify chemical reactions.

Substances that undergo a chemical transformation process are called starting materials.

Substances formed in the process of chemical transformation and not undergoing further chemical changes during this process are called reaction products.

Substances that are formed in one stage of the process of chemical transformation and consumed in other stages of the same process are called intermediates.

The reactions of formation and consumption of intermediate substances are called intermediate reactions.

A chemical reaction that takes place in one phase is called homogeneous chemical reaction(reaction in solution).

The chemical reaction that occurs at the interface is called heterogeneous chemical reaction(reaction on the catalyst surface). It should be noted that in a heterogeneous process, both reactants can be in the same phase. Yes, hydrogenation of ethylene

C 2 H 4 + H 2 → C 2 H 6

goes on the surface of the catalyst, for example, nickel. However, both reactants are in the same phase (in the gas phase above the catalyst surface).

Complex chemical reactions in which some steps are homogeneous and others are heterogeneous are called homogeneous-heterogeneous.

homophasic A process is called a process in which all components: initial, intermediate and final substances are within the limit of one phase. (For example, the reaction to neutralize an acid with an alkali in solution is homogeneous homophasic process).

A process is called heterophasic in which the components form more than one phase (for example, the hydrogenation of ethylene on a nickel catalyst is heterogeneous homophasic process- the process takes place at the interface of the metal and gas phases, and the initial substances and the reaction product are in the same gas phase).

The basic quantity in chemical kinetics is speed reaction.

The rate of a chemical reaction is the change in the concentration of a substance per unit time per unit volume. In general, the reaction rate varies with time and therefore it is better to define it as the derivative of the concentration of the reactant with respect to time (at a constant volume of the system):

Where
is the rate expressed by the decrease in concentration from the reactant; - time. Since over time the concentration of reacting substances decreases, therefore, a minus sign (“–”) is placed in front of the derivative (speed is a positive value).

When two or more substances interact, the reaction rate can be expressed in terms of the derivative of the concentration of any substance.

aA + bB → cc +dD

Equality takes place when the stoichiometric ratio between the participants in the reaction is observed.

The change in concentration with time is expressed by the kinetic curve (
).

Knowing the kinetic curve for any component, one can easily determine the rate of its accumulation or consumption by graphical differentiation of the kinetic curve.

The tangent of the slope of the tangent to the kinetic curve is a graphical interpretation of the rate of a chemical reaction.

The steepness of the kinetic curve characterizes the true rate of a chemical reaction at a certain point in time. In addition, the order and rate constant of the reaction can be determined from the kinetic curves.

In the general case, chemical kinetics studies the optimal conditions for conducting a process only for thermodynamically allowed reactions.

Chemical kinetics has 2 postulates:

I . On the independence of the reaction.

If the process proceeds through a number of stages, then it is assumed that the rate of each individual stage is independent of the rate of the remaining stages.

II . The rate of a chemical reaction is directly proportional to the concentration of the initial substances (ZDM).

aA + bB → cc +dD

This record of the reaction rate expression is called kinetic equation.

The rate of a chemical reaction depends on the concentration of the initial substances, temperature, time, catalyst and the nature of the substances.

k is the rate constant. It is numerically equal to the reaction rate at a concentration of substances equal to unity.

Rate constant k does not depend on the concentration of reagents and time (
). It depends on the temperature, the presence of a catalyst and the nature of the substances (
catalyst, nature of matter ).

Order is the exponent at the concentration of a given substance in the kinetic equation.

In the case of a one-stage process, the exponents are equal to the stoichiometric coefficients:
;
.

The sum of the orders of a reaction over all the reactants is called reaction order(
).

The rate constants of reactions of various orders have different dimensions and are different physical quantities; comparison of their absolute values ​​is meaningless.

First order rate constant: ;

Second order rate constant:
;

Third order rate constant:
.

CLASSIFICATION OF CHEMICAL REACTIONS:

I. By reaction order.n= 0, 1, 2, 3, fractional;

II. By molecularity.

Reaction molecularity is the number of molecules that take part at the same time in one collision event. Molecularity can only be determined by establishing the reaction mechanism. Depending on the number of reacting molecules (particles) participating in an elementary act, one-molecular (monomolecular), two-molecular, trimolecular reactions are distinguished.

TO unimolecular reactions of type A → P include the processes of decomposition of a molecule into simpler constituents and isomerization reactions. Bimolecular called elementary reactions of the form: A + B → P and 2A → P (H 2 + J 2 \u003d 2HJ, HJ + HJ \u003d H 2 + J 2, CH 3 COOCH 3 + H 2 O \u003d CH 3 COOH + CH 3 OH and t .d.). Much less common trimolecular reactions A+2B→P or 3A→P. In all cases, the type and quantity of the resulting reaction products does not matter, since the molecularity is determined only by the number of molecules of substances reacting in an elementary act.

The order of the reaction is established experimentally.

Molecularity and reaction order may or may not be the same. Molecularity and reaction order are the same only for simple reactions that occur in only one elementary stage without the participation of foreign molecules.

Molecularity and order of reaction do not match in three main cases:

1) for complex reactions;

2) for heterogeneous reactions;

3) for reactions with an excess of one of the reactants.

KINETIC EQUATIONS OF DIFFERENT ORDER REACTIONS.

The differentiation of reactions in order occurs according to a formal feature - the sum of the exponents in the kinetic equations of chemical reactions, which limits the possibilities of formal kinetics. Nevertheless, formal kinetics makes it possible to use mathematical relationships to find the kinetic parameters. All dependencies given below are valid for simple homogeneous reactions in closed systems at constant volume and temperature (V=const,T=const).

Zero-order reactions (n=0).

In this case, the reaction rate is constant, since the concentrations of the reaction components are constant.
.

Consider the ester saponification reaction:

The ester saponification rate will be described by the following equation:

1 excess

If we take a large excess of water, then its concentration will be constant and the kinetic equation will take the form:

We can say that the order of the reaction in the partial order of the water component will be zero.

Thus, a large excess of one of the reactants reduces the reaction order by a certain amount.

In the general case, the kinetic equation for a zero-order reaction has the form:

–kineticthe equationzeroorder

For example, the reaction A → P and its rate is described by the equation
, if substance A is taken in large excess, we get:

The rate constant of this reaction is:

Separate the variables and integrate this equation:

At
the integration constant is equal to the initial concentration C 0 (const = C 0), then we get:

;
at n=0

The half-conversion time is often used as a criterion for the rate of a reaction. called half-life.

Half life- this is the time during which half of the taken substance will react.

→
;

– half-life for a zero-order reaction

Zero order occurs in heterogeneous and photochemical reactions.

First order reactions (n=1).

An example of a reaction that strictly obeys a first-order equation is the thermal decomposition of acetone (although the reaction actually proceeds according to a complex mechanism):

CH 3 COCH 3 → CO + CH 3 CH 3

If we denote the concentration of acetone at each moment of time through C, then at constant temperature the reaction rate will be:

Dividing the variables and integrating the equation, we get:

At
constant of integrationconst=lnС 0 , then:

(1)

(2)

Equations (1) and (2) are different forms of the kinetic equation for a first-order reaction. They make it possible to calculate the concentration of a reactant at any time from a known rate constant or, conversely, to find the rate constant of a reaction at a given temperature by determining the concentration at any time. Express half-life for a first order reaction:

Thus, the half-life of a first-order reaction does not depend on the initial concentration of the starting material and is inversely proportional to the rate constant of the reaction.

This dependence can be represented graphically in coordinates
. Since the half-conversion time in this case will be the same, the concentration of the reactant can be determined at each time point.

For practical purposes, it is more advantageous to express the rate in terms of the loss of matter. Let V=const, at the start of the reaction
, the number of moles of the reactant is a. Through seconds reacted x moles of substance A. Then at this point in time the concentration of substance A will be
or
, Where
. After separation of variables and integration, the equation will look like:

At
, x=0

, That's why

A→P (V=const)

Initial number of moles ( =0)

The subject of chemical kinetics is the study of all factors affecting the rate of both the overall process and all intermediate stages.

Encyclopedic YouTube

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✪ Physical chemistry. Lecture 3. Chemical kinetics and catalysis

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✪ Chemistry. Kinetics of chemical reactions. The rate of a chemical reaction. Foxford Online Learning Center

✪ Introduction to kinetics

✪ Chemical kinetics

Subtitles

Basic concepts

homogeneous reaction - reaction in which the reactants are in the same phase

Heterogeneous reaction - a reaction that occurs at the interfaces - between a gaseous substance and a solution, between a solution and a solid substance, between a solid and gaseous substances

The reaction is called simple if the product is formed as a result of the direct interaction of the molecules (particles) of the reagents

A reaction is called complex if the final product is obtained as a result of the implementation of two or more simple reactions (elementary acts) with the formation of intermediate products

The rate of a chemical reaction

An important concept in chemical kinetics is chemical reaction rate. This value determines how the concentration of the reaction components changes over time. The rate of a chemical reaction is always positive, so if it is determined by the initial substance (the concentration of which decreases during the reaction), then the resulting value is multiplied by −1.
For example, for a reaction, the rate can be expressed as:

A + B → C + D , (\displaystyle A+B\to C+D,) v = ∂ C ∂ t = − ∂ A ∂ t . (\displaystyle v=(\frac (\partial C)(\partial t))=-(\frac (\partial A)(\partial t)).)

Order of a chemical reaction

Reaction order according to given substance- exponent at the concentration of this substance in the kinetic equation of the reaction.

Zero order reaction

The kinetic equation has the following form:

V 0 = k 0 (\displaystyle V_(0)=k_(0))

The rate of a zero-order reaction is constant in time and does not depend on the concentrations of reactants. Zero order is typical, for example, for heterogeneous reactions if the rate of diffusion of reactants to the interface is less than the rate of their chemical transformation.

First order reaction

Kinetic equation of the first order reaction:

V 1 = k 1 ⋅ C = − d C d τ (\displaystyle V_(1)=k_(1)\cdot C=-(\frac (dC)(d\tau )))

Reducing the equation to a linear form gives the equation:

ln ⁡ C = ln ⁡ C 0 − k 1 ⋅ τ (\displaystyle \ln C=\ln C_(0)-k_(1)\cdot \tau )

The reaction rate constant is calculated as the tangent of the slope of the straight line to the time axis:

k 1 = − t g α (\displaystyle k_(1)=-\mathrm (tg) \alpha )

Half life:

τ 1 2 = ln ⁡ 2 k 1 (\displaystyle \tau _(\frac (1)(2))=(\frac (\ln 2)(k_(1))))

Second order reaction

For second-order reactions, the kinetic equation has the following form:

V = k 2 C A 2 (\displaystyle V=k_(2)(C_(A))^(2)) V = k 2 C A ⋅ C B (\displaystyle V=k_(2)C_(A)\cdot C_(B))

In the first case, the reaction rate is determined by the equation

V = k 2 C A 2 = − d C d τ (\displaystyle V=k_(2)(C_(A))^(2)=-(\frac (dC)(d\tau )))

Linear form of the equation:

1 C = k 2 ⋅ τ + 1 C 0 (\displaystyle (\frac (1)(C))=k_(2)\cdot \tau +(\frac (1)(C_(0))))

The reaction rate constant is equal to the tangent of the slope of the straight line to the time axis:

k 2 = − t g α (\displaystyle k_(2)=-\mathrm (tg) \alpha ) k 2 = 1 τ (1 C − 1 C 0) (\displaystyle k_(2)=(\frac (1)(\tau ))\left((\frac (1)(C))-(\frac ( 1)(C_(0)))\right))

In the second case, the expression for the reaction rate constant will look like this:

k 2 = 1 τ (C 0 , A − C 0 , B) ln ⁡ C 0 , B ⋅ C A C 0 , A ⋅ C B (\displaystyle k_(2)=(\frac (1)(\tau (C_(0 ,A)-C_(0,B))))\ln (\frac (C_(0,B)\cdot C_(A))(C_(0,A)\cdot C_(B))))

Half-life (for the case of equal initial concentrations!):

τ 1 2 = 1 k 2 ⋅ 1 C 0 (\displaystyle \tau _(\frac (1)(2))=(\frac (1)(k_(2)))\cdot (\frac (1)( C_(0))))

Reaction molecularity

The molecularity of an elementary reaction is the number of particles that, according to the experimentally established reaction mechanism, participate in an elementary act of chemical interaction.

Monomolecular reactions- reactions in which a chemical transformation of one molecule occurs (isomerization, dissociation, etc.):

H 2 S → H 2 + S (\displaystyle (\mathsf (H_(2)S\rightarrow H_(2)+S)))

Bimolecular reactions- reactions, the elementary act of which is carried out when two particles (identical or different) collide:

C H 3 B r + K O H → C H 3 O H + K B r (\displaystyle (\mathsf (CH_(3)Br+KOH\rightarrow CH_(3)OH+KBr)))

Trimolecular reactions- reactions, the elementary act of which is carried out by the collision of three particles:

N O + N O + O 2 → 2 N O 2 (\displaystyle (\mathsf (NO+NO+O_(2)\rightarrow 2NO_(2))))

Reactions with a molecularity greater than three are unknown.

For elementary reactions carried out at close concentrations of the starting substances, the values ​​of molecularity and order of the reaction are the same. There is no clearly defined relationship between the concepts of molecularity and reaction order, since the reaction order characterizes the kinetic equation of the reaction, and molecularity characterizes the reaction mechanism.

Catalysis

. An example of a negative is a decrease in the corrosion rate when sodium nitrite, chromate and potassium dichromate are introduced into the liquid in which the metal is operated.

Many important chemical production, such as the production of sulfuric acid, ammonia, nitric acid, synthetic rubber, a number of polymers, etc., are carried out in the presence of catalysts.

Catalysis in biochemistry

Enzymatic catalysis is inextricably linked with the vital activity of plant and animal organisms. Many of the vital chemical reactions that take place in a cell (something like ten thousand) are controlled by special organic catalysts called ferments or enzymes. The term "special" should not be given close attention, since it is already known what these enzymes are built of. Nature has chosen for this one and only construction material- amino acids and connected them into polypeptide chains of various lengths and in different sequences

This is the so-called primary structure of the enzyme, where R are side residues, or the most important functional groups of proteins, possibly acting as active centers of enzymes. These side groups are the main load during the work of the enzyme, while the peptide chain plays the role of a supporting skeleton. According to the Pauling-Corey structural model, it is folded into a spiral, which in the normal state is stabilized by hydrogen bonds between acidic and basic centers:

For some enzymes, the complete amino acid composition and the sequence of their arrangement in the chain, as well as a complex spatial structure, have been established. But this still very often cannot help us answer two main questions: 1) why are enzymes so selective and accelerate the chemical transformations of molecules of only a very specific structure (which we also know); 2) how the enzyme lowers the energy barrier, that is, it chooses an energetically more favorable path, due to which the reactions can proceed at ordinary temperature.

Strict selectivity and high speed are the two main features of enzymatic catalysis, which distinguish it from laboratory and industrial catalysis. None of the man-made catalysts (with the possible exception of 2-hydroxypyridine) can be compared with enzymes in terms of the strength and selectivity of their effect on organic molecules. The activity of an enzyme, like that of any other catalyst, also depends on temperature: with an increase in temperature, the rate of the enzymatic reaction also increases. At the same time, a sharp decrease in the activation energy E compared to a non-catalytic reaction attracts attention. True, this does not always happen. There are many known cases where the speed increases due to an increase in the temperature-independent pre-exponential factor in the Arrhenius equation.

Types of enzymatic reactions

• Ping pong type- the enzyme first interacts with substrate A, taking away any chemical groups from it and converting it into the corresponding product. Substrate B is then attached to the enzyme, receiving these chemical groups. An example is the transfer of amino groups from amino acids to keto acids: transamination.
• Type of sequential reactions- substrates A and B are sequentially attached to the enzyme, forming a "triple complex", after which catalysis is carried out. The reaction products are also sequentially cleaved off from the enzyme.
• Type of random interactions- substrates A and B are attached to the enzyme in any order, randomly, and after catalysis they are also cleaved off.

In the 19th century As a result of the development of the fundamentals of chemical thermodynamics, chemists have learned to calculate the composition of an equilibrium mixture for reversible chemical reactions. In addition, on the basis of simple calculations, it was possible, without conducting experiments, to draw a conclusion about the fundamental possibility or impossibility of a particular reaction occurring under given conditions. However, the "fundamental possibility" of a reaction does not yet mean that it will take place. For example, the reaction C + O 2 → CO 2 is very favorable from the point of view of thermodynamics, in any case, at temperatures below 1000 ° C (at higher temperatures, the decomposition of CO 2 molecules already occurs), i.e. carbon and oxygen should (practically with 100% yield) turn into carbon dioxide. However, experience shows that a piece of coal can lie in the air for years, with free access to oxygen, without undergoing any changes. The same can be said about many other known reactions. For example, mixtures of hydrogen with chlorine or oxygen can be preserved for a very long time without any signs of chemical reactions, although in both cases the reactions are thermodynamically favorable. This means that after reaching equilibrium in the stoichiometric mixture H 2 + Cl 2, only hydrogen chloride should remain, and in the mixture 2H 2 + O 2 - only water. Another example: gaseous acetylene is quite stable, although the reaction C 2 H 2 → 2C + H 2 is not only thermodynamically allowed, but is also accompanied by a significant release of energy. Indeed, at high pressures, acetylene explodes, but under normal conditions it is quite stable.

Thermodynamically allowed reactions, like those considered above, can proceed only under certain conditions. For example, after ignition, coal or sulfur spontaneously combines with oxygen; hydrogen easily reacts with chlorine when the temperature rises or when exposed to ultraviolet light; a mixture of hydrogen and oxygen (explosive gas) explodes when ignited or when a catalyst is added. Why, then, for the implementation of all these reactions, special effects are necessary - heating, irradiation, the action of catalysts? Chemical thermodynamics does not give an answer to this question - the concept of time is absent in it. At the same time, for practical purposes, it is very important to know whether a given reaction will take place in a second, in a year, or in many millennia.

Experience shows that the rate of different reactions can vary greatly. Many reactions take place almost instantaneously in aqueous solutions. So, when an excess of acid is added to an alkaline solution of raspberry-colored phenolphthalein, the solution instantly becomes colorless, which means that the neutralization reaction, as well as the reaction of turning the colored form of the indicator into colorless, proceed very quickly. The reaction of oxidation of an aqueous solution of potassium iodide with atmospheric oxygen proceeds much more slowly: the yellow color of the reaction product, iodine, appears only after a long time. Corrosion processes of iron and especially copper alloys and many other processes proceed slowly.

Predicting the rate of a chemical reaction, as well as elucidating the dependence of this rate on the reaction conditions, is one of the important tasks of chemical kinetics, a science that studies the patterns of reactions in time. No less important is the second task facing chemical kinetics - the study of the mechanism of chemical reactions, that is, the detailed path for the transformation of starting materials into reaction products.

Speed ​​reaction.

The easiest way is to determine the rate for a reaction between gaseous or liquid reactants in a homogeneous (homogeneous) mixture in a vessel of constant volume. In this case, the reaction rate is defined as the change in the concentration of any of the substances involved in the reaction (it can be the starting substance or the reaction product) per unit time. This definition can be written as a derivative: v=d c/d t, Where v- speed reaction; t- time, c- concentration. This rate is easy to determine if there are experimental data on the dependence of the concentration of a substance on time. Based on these data, you can build a graph called the kinetic curve. The reaction rate at a given point on the kinetic curve is determined by the slope of the tangent at that point. Determining the slope of a tangent is always associated with some error. The initial reaction rate is most accurately determined, since at first the kinetic curve is usually close to a straight line; this makes it easier to draw a tangent at the starting point of the curve.

If the time is measured in seconds, and the concentration is in moles per liter, then the reaction rate is measured in units of mol / (l s). Thus, the reaction rate does not depend on the volume of the reaction mixture: under the same conditions, it will be the same both in a small test tube and in a large-tonnage reactor.

d value t is always positive, while the sign of d c depends on how the concentration changes with time - decreases (for the starting substances) or increases (for the reaction products). In order for the reaction rate to always remain a positive value, in the case of starting substances, a minus sign is placed in front of the derivative: v= -d c/d t. If the reaction proceeds in the gas phase, instead of the concentration of substances in the rate equation, pressure is often used. If the gas is close to ideal, then the pressure R associated with concentration simple equation: p = cRT.

In the course of the reaction, various substances can be consumed and formed with different speed, in accordance with the coefficients in the stoichiometric equation ( cm. STOICHIOMETRY), therefore, when determining the rate of a particular reaction, these coefficients should be taken into account. For example, in the ammonia synthesis reaction 3H 2 + N 2 → 2NH 3, hydrogen is consumed 3 times faster than nitrogen, and ammonia accumulates 2 times faster than nitrogen is consumed. Therefore, the rate equation for this reaction is written as follows: v= –1/3d p(H2)/d t= -d p(N 2)/d t= +1/2d p(NH3)/d t. In general, if the reaction is stoichiometric, i.e. flows exactly according to the written equation: aA + bB → cC + dD, its speed is defined as v= –(1/a)d[A]/d t= –(1/b)d[B]/d t= (1/c)d[C]/d t= (1/d)d[D]/d t(in square brackets it is customary to indicate the molar concentration of substances). Thus, the rates for each substance are rigidly interconnected and, having determined experimentally the rate for any participant in the reaction, it is easy to calculate it for any other substance.

Most of the reactions used in industry are heterogeneous catalytic. They flow at the interface between the solid catalyst and the gas or liquid phase. At the interface between the two phases, reactions such as the roasting of sulfides, the dissolution of metals, oxides and carbonates in acids, and a number of other processes also occur. For such reactions, the rate also depends on the size of the interface, so the rate of a heterogeneous reaction is referred not to a volume unit, but to a surface unit. It is not always easy to measure the size of the surface on which the reaction takes place.

If the reaction proceeds in a closed volume, then in most cases its rate is maximum at the initial moment of time (when the concentration of the initial substances is maximum), and then, as the initial reagents are converted into products and, accordingly, their concentration decreases, the reaction rate decreases. There are also reactions in which the rate increases with time. For example, if a copper plate is dipped into a solution of pure nitric acid, then the reaction rate will increase with time, which is easy to observe visually. The processes of aluminum dissolution in alkali solutions, oxidation of many organic compounds oxygen, a number of other processes. The reasons for this acceleration may be different. For example, this may be due to the removal of a protective oxide film from the metal surface, or with the gradual heating of the reaction mixture, or with the accumulation of substances that accelerate the reaction (such reactions are called autocatalytic).

In industry, reactions are usually carried out by continuously feeding the starting materials into the reactor and withdrawing the products. Under such conditions, a constant rate of a chemical reaction can be achieved. Photochemical reactions also proceed at a constant rate provided that the incident light is completely absorbed ( cm. PHOTOCHEMICAL REACTIONS).

limiting step of the reaction.

If the reaction is carried out by sequentially occurring stages (not necessarily all of them are chemical) and one of these stages requires much more time than the others, that is, it goes much more slowly, then such a stage is called rate-limiting. It is this slowest stage that determines the speed of the entire process. Consider, as an example, the catalytic oxidation of ammonia. There are two limiting cases here.

1. The flow of reagent molecules - ammonia and oxygen to the catalyst surface (physical process) is much slower than the catalytic reaction itself on the surface. Then, in order to increase the rate of formation of the target product, nitric oxide, it is completely useless to increase the efficiency of the catalyst, but care must be taken to accelerate the access of reagents to the surface.

2. The supply of reagents to the surface occurs much faster than the chemical reaction itself. Here it makes sense to improve the catalyst, to select the optimal conditions for the catalytic reaction, since the rate-limiting step in this case is the catalytic reaction on the surface.

collision theory.

Historically, the first theory on the basis of which it was possible to calculate the rates of chemical reactions was the theory of collisions. Obviously, in order for the molecules to react, they must first of all collide. It follows that the reaction should go faster, the more often the molecules of the starting substances collide with each other. Therefore, every factor that affects the frequency of collisions between molecules will also affect the rate of the reaction. Some important regularities concerning collisions between molecules were obtained on the basis of the molecular kinetic theory of gases.

In the gas phase, molecules move at high speeds (hundreds of meters per second) and very often collide with each other. The frequency of collisions is determined primarily by the number of particles per unit volume, that is, by the concentration (pressure). The frequency of collisions also depends on the temperature (as the temperature rises, the molecules move faster) and on the size of the molecules (large molecules collide with each other more often than small ones). However, concentration affects the frequency of collisions much more strongly. At room temperature and atmospheric pressure, each medium-sized molecule experiences several billion collisions per second.

Based on these data, it is possible to calculate the rate of the reaction A + B → C between two gaseous compounds A and B, assuming that a chemical reaction takes place with each collision of the reactant molecules. Let a liter flask at atmospheric pressure contain a mixture of reagents A and B at equal concentrations. In total, there will be 6 10 23 / 22.4 = 2.7 10 22 molecules in the flask, of which 1.35 10 22 molecules of substance A and the same number of molecules of substance B. Let each molecule A experience 10 9 collisions in 1 s with other molecules, of which half (5 10 8) falls on collisions with molecules B (collisions A + A do not lead to a reaction). Then, in total, 1.35 10 22 5 10 8 ~ 7 10 30 collisions of molecules A and B occur in the flask in 1 s. It is obvious that if each of them led to a reaction, it would take place instantly. However, many reactions are quite slow. From this we can conclude that only an insignificant fraction of collisions between reactant molecules leads to interaction between them.

To create a theory that would allow calculating the reaction rate based on the molecular kinetic theory of gases, it was necessary to be able to calculate the total number of collisions of molecules and the proportion of "active" collisions that lead to reactions. It was also necessary to explain why the rate of most chemical reactions greatly increases with increasing temperature - the speed of molecules and the frequency of collisions between them increase slightly with temperature - proportionally, that is, only 1.3 times with an increase in temperature from 293 K (20 ° C) to 373 K (100 ° C), while the reaction rate can increase thousands of times.

These problems were solved on the basis of collision theory in the following way. During collisions, molecules continuously exchange velocities and energies. So, as a result of a “successful” collision, a given molecule can noticeably increase its speed, while in an “unsuccessful” collision it can almost stop (a similar situation can be observed in the example of billiard balls). At normal atmospheric pressure, collisions, and therefore changes in speed, occur with each molecule billions of times per second. In this case, the velocities and energies of the molecules are to a large extent averaged. If at a given moment of time "count" in a given volume of gas the molecules that have certain velocities, it turns out that a significant part of them has a speed close to the average. At the same time, many molecules have a speed less than the average, and some move with speeds above the average. With increasing speed, the fraction of molecules having given speed, decreases rapidly. In accordance with the theory of collisions, only those molecules react that, upon collision, have a sufficiently high speed (and, consequently, a large supply of kinetic energy). Such an assumption was made in 1889 by the Swedish chemist Svante Arrhenius.

Activation energy.

Arrhenius brought into use chemists very important concept activation energy ( E a) is the minimum energy that a molecule (or a pair of reacting molecules) must have in order to enter into a chemical reaction. The activation energy is usually measured in joules and is not related to one molecule (this is a very small value), but to a mole of a substance and is expressed in units of J / mol or kJ / mol. If the energy of the colliding molecules is less than the activation energy, then the reaction will not proceed, and if it is equal to or greater, then the molecules will react.

The activation energies for different reactions are determined experimentally (from the dependence of the reaction rate on temperature). The activation energy can change over a fairly wide range, from units to several hundred kJ/mol. For example, for the reaction 2NO 2 → N 2 O 4, the activation energy is close to zero, for the reaction 2H 2 O 2 → 2H 2 O + O 2 in aqueous solutions E a = 73 kJ/mol, for thermal decomposition of ethane into ethylene and hydrogen E a = 306 kJ/mol.

The activation energy of most chemical reactions significantly exceeds the average kinetic energy of molecules, which at room temperature is only about 4 kJ / mol, and even at a temperature of 1000 ° C does not exceed 16 kJ / mol. Thus, in order to react, molecules usually must have a speed significantly greater than the average. For example, in the case E a = 200 kJ/mol colliding molecules small molecular weight should have a speed of the order of 2.5 km/s (the activation energy is 25 times the average energy of molecules at 20°C). And this - general rule: for most chemical reactions, the activation energy is much higher than the average kinetic energy of the molecules.

The probability for a molecule to store a lot of energy as a result of a series of collisions is very small: such a process requires for it an enormous number of successive "successful" collisions, as a result of which the molecule only gains energy without losing it. Therefore, for many reactions, only a tiny fraction of the molecules have enough energy to overcome the barrier. This share, in accordance with the Arrhenius theory, is determined by the formula: a \u003d e - E a / RT = 10 –E a /2,3 RT ~ 10 –E a /19 T, Where R= 8.31 J/(mol . TO). It follows from the formula that the fraction of molecules with energy E a , as well as the fraction of active collisions a, depends very strongly on both the activation energy and the temperature. For example, to react with E a = 200 kJ/mol at room temperature ( T~ 300 K), the fraction of active collisions is negligible: a = 10 –200000/(19 , 300) ~ 10 –35 . And if every second there are 7 10 30 collisions of molecules A and B in the vessel, then it is clear that the reaction will not go.

If we double the absolute temperature, i.e. heat the mixture to 600 K (327 ° C); in this case, the fraction of active collisions will sharply increase: a = 10 –200000/(19 , 600) ~ 4 10 -18 . Thus, a twofold increase in temperature increased the fraction of active collisions by a factor of 4·1017. Now every second of the total number of approximately 7·10 30 collisions, 7·10 30 ·4·10 -18 ~ 3·10 13 will lead to the reaction. Such a reaction, in which 3 10 13 molecules disappear every second (out of about 10 22), although very slowly, but still goes on. Finally, at a temperature of 1000 K (727° C) a ~ 3 10 -11 (out of every 30 billion collisions of a given reagent molecule, one leads to a reaction). This is already a lot, since in 1 s 7 10 30 3 10 -11 = 2 10 20 molecules will enter into the reaction, and such a reaction will take place in several minutes (taking into account the decrease in the frequency of collisions with a decrease in the concentration of reagents).

Now it is clear why increasing the temperature can increase the rate of the reaction so much. The average speed (and energy) of molecules increases insignificantly with increasing temperature, but on the other hand, the proportion of “fast” (or “active”) molecules that have a sufficient speed of movement or sufficient vibrational energy for the reaction to proceed increases sharply.

Calculation of the reaction rate, taking into account the total number of collisions and the fraction of active molecules (ie activation energy), often gives a satisfactory agreement with experimental data. However, for many reactions the experimentally observed rate turns out to be less than that calculated by the collision theory. This is explained by the fact that for the reaction to take place, it is necessary that the collision be successful not only energetically, but also “geometrically”, that is, the molecules must be oriented relative to each other at the moment of collision in a certain way. Thus, when calculating the rate of reactions according to the theory of collisions, in addition to the energy factor, the steric (spatial) factor for a given reaction is also taken into account.

Arrhenius equation.

The temperature dependence of the reaction rate is usually described by the Arrhenius equation, which in its simplest form can be written as v = v 0 a = v 0e- E a/ RT, Where v 0 is the rate that the reaction would have at zero activation energy (in fact, this is the frequency of collisions per unit volume). Because the v 0 weakly depends on temperature, everything is determined by the second factor - exponential: with increasing temperature, this factor increases rapidly, and the faster, the greater the activation energy E A. This dependence of the reaction rate on temperature is called the Arrhenius equation, it is one of the most important in chemical kinetics. For an approximate assessment of the effect of temperature on the reaction rate, the so-called "van't Hoff rule" is sometimes used ( cm. VANT HOFF RULE).

If the reaction obeys the Arrhenius equation, the logarithm of its rate (measured, for example, at the initial moment) should depend linearly on the absolute temperature, that is, the plot of ln v from 1/ T should be straight forward. The slope of this straight line is equal to the activation energy of the reaction. From such a graph, one can predict what the reaction rate will be at a given temperature, or at what temperature the reaction will proceed at a given rate.

Some practical examples using the Arrhenius equation.

1. The packaging of the frozen product says that it can be stored on the refrigerator shelf (5 ° C) for a day, in the freezer marked with one asterisk (-6 ° C) - a week, with two asterisks (-12 ° C) - a month , and in the freezer with the *** symbol (which means the temperature in it is -18 ° C) - 3 months. Assuming that the rate of product spoilage is inversely proportional to the guaranteed shelf life t xp, in coordinates ln t xp, 1/ T we obtain, in accordance with the Arrhenius equation, a straight line. From it, you can calculate the activation energy of biochemical reactions leading to spoilage of a given product (about 115 kJ/mol). From the same graph, you can find out to what temperature the product must be cooled so that it can be stored, for example, 3 years; it turns out -29 ° C.

2. Climbers know that in the mountains it is difficult to boil an egg, and in general any food that requires more or less long boiling. Qualitatively, the reason for this is clear: with a decrease atmospheric pressure the boiling point of water decreases. Using the Arrhenius equation, you can calculate how long it will take, for example, to boil an egg in Mexico City, located at an altitude of 2265 m, where a pressure of 580 mm Hg is considered normal, and water at this reduced pressure boils at 93 ° C The activation energy of the protein "folding" (denaturation) reaction has been measured and found to be very large compared to many other chemical reactions - about 400 kJ / mol (it may differ slightly for different proteins). In this case, lowering the temperature from 100 to 93 ° C (that is, from 373 to 366 K) will slow down the reaction by 10 (400,000/19) (1/366 - 1/373) = 11.8 times. That is why the inhabitants of the highlands prefer frying their food to cooking: the temperature of a frying pan, unlike the temperature of a pot of boiling water, does not depend on atmospheric pressure.

3. In a pressure cooker, food is cooked at elevated pressure and, therefore, at an elevated boiling point of water. It is known that in an ordinary saucepan beef is cooked for 2-3 hours, and apple compote - 10-15 minutes. Given that both processes have similar activation energies (about 120 kJ/mol), it can be calculated using the Arrhenius equation that meat will be cooked in a pressure cooker at 118°C for 25–30 minutes, and compote for only 2 minutes.

The Arrhenius equation is very important for the chemical industry. When an exothermic reaction occurs, the released thermal energy heats not only environment but also the reagents themselves. this can lead to an undesirable strong acceleration of the reaction. The calculation of the change in the reaction rate and the rate of heat release with increasing temperature makes it possible to avoid a thermal explosion ( cm. EXPLOSIVES).

Dependence of the reaction rate on the concentration of reagents.

The rate of most reactions gradually decreases over time. This result is in good agreement with the collision theory: as the reaction proceeds, the concentrations of the initial substances decrease, and the frequency of collisions between them also decreases; correspondingly, the frequency of collision of active molecules also decreases. This leads to a decrease in the reaction rate. This is the essence of one of the basic laws of chemical kinetics: the rate of a chemical reaction is proportional to the concentration of reacting molecules. Mathematically, this can be written as a formula v = k[A][B], where k is a constant called the rate constant of the reaction. The above equation is called the chemical reaction rate equation or the kinetic equation. The rate constant for this reaction does not depend on the concentration of the reactants and on time, but it depends on temperature in accordance with the Arrhenius equation: k = k 0e- E a/ RT .

The simplest equation of speed v = k[A][B] is always true in the case when molecules (or other particles, for example, ions) A, colliding with molecules B, can directly turn into reaction products. Similar reactions that take place in one step (as chemists say, in one stage) are called elementary reactions. There are few such reactions. Most reactions (even such seemingly simple ones as H 2 + I 2 ® 2HI) are not elementary, therefore, based on the stoichiometric equation of such a reaction, its kinetic equation cannot be written.

The kinetic equation can be obtained in two ways: experimentally - by measuring the dependence of the reaction rate on the concentration of each reactant separately, and theoretically - if the detailed reaction mechanism is known. Most often (but not always) the kinetic equation has the form v = k[A] x[B] y, Where x And y are called reaction orders for reactants A and B. These orders, in the general case, can be integer and fractional, positive and even negative. For example, the kinetic equation for the thermal decomposition reaction of acetaldehyde CH 3 CHO ® CH 4 + CO has the form v = k 1.5, i.e. the reaction is one and a half order. Occasionally, coincidence of stoichiometric coefficients and reaction orders is possible. So, the experiment shows that the reaction H 2 + I 2 ® 2HI has the first order both in hydrogen and in iodine, that is, its kinetic equation has the form v = k(which is why this reaction was considered elementary for many decades until its more complex mechanism was proven in 1967).

If the kinetic equation is known, i.e. it is known how the reaction rate depends on the concentrations of the reactants at each moment of time, and the rate constant is known, then it is possible to calculate the time dependence of the concentrations of the reactants and reaction products, i.e. theoretically obtain all kinetic curves. For such calculations, methods of higher mathematics or computer calculations are used, and they do not present fundamental difficulties.

On the other hand, the experimentally obtained kinetic equation helps to judge the reaction mechanism, i.e. about the totality of simple (elementary) reactions. Elucidation of reaction mechanisms is the most important task of chemical kinetics. This is very difficult task, since the mechanism of even a seemingly simple reaction can include many elementary stages.

The application of kinetic methods to determine the reaction mechanism can be illustrated by the example of alkaline hydrolysis of alkyl halides with the formation of alcohols: RX + OH – → ROH + X – . It was experimentally found that for R = CH 3 , C 2 H 5 , etc. and X = Cl the reaction rate is directly proportional to the concentrations of the reactants, i.e. has the first order with respect to the halide RX and the first order with respect to the alkali, and the kinetic equation has the form v = k 1 . In the case of tertiary alkyl iodides (R \u003d (CH 3) 3 C, X \u003d I), the order in RX is the first, and in alkali it is zero: v = k 2. In intermediate cases, for example, for isopropyl bromide (R = (CH 3) 2 CH, X = Br), the reaction is described by a more complex kinetic equation: v = k 1 + k 2. Based on these kinetic data, the following conclusion was made about the mechanisms of such reactions.

In the first case, the reaction proceeds in one step, by direct collision of alcohol molecules with OH ions - (the so-called SN 2 mechanism). In the second case, the reaction proceeds in two stages. The first stage is the slow dissociation of alkyl iodide into two ions: RI → R + + I – . The second is a very fast reaction between ions: R + + OH - → ROH. The rate of the overall reaction depends only on the slow (limiting) stage, so it does not depend on the concentration of alkali; hence the zero order in alkali (SN 1 mechanism). In the case of secondary alkyl bromides, both mechanisms occur simultaneously, so the kinetic equation is more complex.

Ilya Leenson

Literature:

History of the doctrine of the chemical process. M., Science, 1981
Leenson I.A. chemical reactions. M., AST - Astrel, 2002



Chemical kinetics

Chapter 6

Chemical kinetics. chemical balance.

Chemical kinetics.

Chemical kinetics - branch of chemistry that studies the rates and mechanisms of chemical processes, as well as their dependence on various factors.

The study of the kinetics of chemical reactions makes it possible both to determine the mechanisms of chemical processes and to control chemical processes in their practical implementation.

Any chemical process is the conversion of reactants into reaction products:

reactants → transition state → reaction products.

Reagents (source substances) - substances entering the process of chemical interaction.

reaction products- substances formed at the end of the chemical transformation process. In reversible processes, the products of the forward reaction are the reactants of the reverse reaction.

irreversible reactions- reactions occurring under given conditions in almost one direction (denoted by the sign →).

For example:

CaCO 3 → CaO + CO 2

Reversible reactions- reactions proceeding simultaneously in two opposite directions (denoted by a sign).

For example:

transition state (activated complex) is a state chemical system, which is intermediate between the starting materials (reagents) and the reaction products. In this state, old chemical bonds are broken and new chemical bonds are formed. Further, the activated complex is converted into reaction products.

Most chemical reactions are complex and consists of several stages, called elementary reactions .

elementary reaction- a single act of formation or rupture of a chemical bond. The set of elementary reactions that make up a chemical reaction determines mechanism of a chemical reaction.

The equation of a chemical reaction usually indicates the initial state of the system (initial substances) and its final state (reaction products). At the same time, the actual mechanism of a chemical reaction can be quite complex and include a number of elementary reactions. Complex chemical reactions include reversible, parallel, sequential and other multi-stage reactions (chain reactions, coupled reactions, etc.).

If the rates of various stages of a chemical reaction differ significantly, then the rate of a complex reaction as a whole is determined by the rate of its slowest stage. This stage (elementary reaction) is called limiting stage.

Depending on the phase state of the reacting substances, there are two types of chemical reactions: homogeneous And heterogeneous.

phase a part of a system that differs in its physical and chemical properties from other parts of the system and separated from them by the interface. Single phase systems are called homogeneous systems, from several phases - heterogeneous. An example of a homogeneous system can be air, which is a mixture of substances (nitrogen, oxygen, etc.) that are in the same gas phase. A suspension of chalk (solid) in water (liquid) is an example of a heterogeneous two-phase system.

Accordingly, reactions in which the interacting substances are in the same phase are called homogeneous reactions. The interaction of substances in such reactions occurs throughout the entire volume of the reaction space.

Heterogeneous reactions include reactions occurring at the phase boundary. In a heterogeneous system, the reaction always occurs at the interface between two phases, since only here the reacting substances that are in different phases can collide with each other.

Chemical reactions are usually distinguished by their molecularity, those. according to the number of molecules involved in each elementary act of interaction . On this basis, reactions are distinguished monomolecular, bimolecular and trimolecular.

Monomolecular called reactions in which the elementary act is a chemical transformation of one molecule , For example:

Bimolecular reactions are considered, in which the elementary act is carried out when two molecules collide, for example:

IN trimolecular reactions, an elementary act is carried out with the simultaneous collision of three molecules, for example:

The collision of more than three molecules at the same time is almost improbable, therefore reactions of greater molecularity do not occur in practice.

The rates of chemical reactions can vary significantly. Chemical reactions can proceed extremely slowly, over entire geological periods, such as rock weathering, which is the transformation of aluminosilicates:

K 2 O Al 2 O 3 6SiO 2 + CO 2 + 2H 2 O → K 2 CO 3 + 4SiO 2 + Al 2 O 3 2SiO 2 2H 2 O.

orthoclase - feldspar potash quartz. sand kaolinite (clay)

Some reactions proceed almost instantly, for example, the explosion of black powder, which is a mixture of coal, sulfur and nitrate:

3C + S + 2KNO 3 = N 2 + 3CO 2 + K 2 S.

The rate of a chemical reaction is a quantitative measure of the intensity of its occurrence.

In general under the speed of a chemical reaction understand the number of elementary reactions occurring per unit of time in a unit of reaction space.

Since for homogeneous processes the reaction space is the volume of the reaction vessel, then

for homogeneous reactions with The rate of a chemical reaction is determined by the amount of a substance that has reacted per unit time per unit volume.

Considering that the amount of a substance contained in a certain volume characterizes the concentration of a substance, then

the reaction rate is a value showing the change in the molar concentration of one of the substances per unit time.

If, at constant volume and temperature, the concentration of one of the reactants decreases from With 1 to With 2 for a period of time from t 1 to t 2 , then, in accordance with the definition, the reaction rate for a given period of time (average reaction rate) is equal to:

.

Usually, for homogeneous reactions, the dimension of the rate V[mol/l s].

Since for heterogeneous reactions the reaction space is surface , on which the reaction takes place, then for heterogeneous chemical reactions, the reaction rate refers to the unit area of ​​the surface on which the reaction takes place. Accordingly, the average rate of a heterogeneous reaction has the form:

,

Where S is the surface area on which the reaction takes place.

The unit of speed for heterogeneous reactions is [mol/l·s·m 2 ].

The rate of a chemical reaction depends on a number of factors:

the nature of the reactants;

concentrations of reactants;

pressure (for gas systems);

system temperature;

surface area (for heterogeneous systems);

the presence of a catalyst in the system and other factors.

Since each chemical interaction is the result of particle collisions, an increase in concentration (the number of particles in a given volume) leads to more frequent collisions, and as a result, to an increase in the reaction rate. The dependence of the rate of chemical reactions on the molar concentrations of the reactants is described by the basic law of chemical kinetics - law of acting masses , which was formulated in 1865 by N.N. Beketov and in 1867 by K.M. Guldberg and P. Waage.

Law of acting masses reads: the rate of an elementary chemical reaction at a constant temperature is directly proportional to the product of the molar concentrations of the reactantsin powers equal to their stoichiometric coefficients.

The equation expressing the dependence of the reaction rate on the concentration of each substance is called reaction kinetic equation .

It should be noted that the law of mass action is fully applicable only to the simplest homogeneous reactions. If the reaction proceeds in several stages, then the law is valid for each of the stages, and the rate of a complex chemical process is determined by the rate of the slowest reaction, which is limiting stage the whole process.

In the general case, if an elementary reaction enters simultaneously T substance molecules A And n substance molecules IN:

mA + nV = WITH,

then the equation for the reaction rate (kinetic equation) looks like:

Where k is the coefficient of proportionality, which is called rate constant chemical reaction; [ A A; [B] - molar concentration of a substance B;m And n- stoichiometric coefficients in the reaction equation.

To understand physical meaning reaction rate constants , must be taken in the above equations for the concentration of reactants [ A] = 1 mol/l and [ IN] = 1 mol/l (or equate their product to unity), and then:

Hence it is clear that reaction rate constant k is numerically equal to the reaction rate in which the concentrations of reactants (or their product in kinetic equations) are equal to unity.

Reaction rate constant k depends on the nature of the reactants and temperature, but does not depend on the value of the concentration of the reactants.

For heterogeneous reactions, the concentration of the solid phase is not included in the expression for the rate of a chemical reaction.

For example, in the methane synthesis reaction:

C (t) + 2H 2 (g) CH 4 (g),

according to the law of mass action, the reaction rate is determined only by the concentration of hydrogen, and the surface area of ​​solid carbon is taken into account by the rate constant of the chemical reaction k. The kinetic equation for this reaction will be:

If the reaction proceeds in the gas phase, then a change in the pressure in the system has a significant effect on its rate, since a change in pressure in the gas phase leads to a proportional change in concentration. Thus, an increase in pressure leads to a proportional increase in concentration, and a decrease in pressure, respectively, reduces the concentration of the gaseous reactant.

A change in pressure has practically no effect on the concentration of liquid and solids(condensed state of matter) and does not affect the rate of reactions occurring in liquid or solid phases.

Chemical reactions are carried out due to the collision of particles of reacting substances. However, not every collision of reactant particles is effective; leads to the formation of reaction products. Only particles with increased energy - active particles capable of carrying out a chemical reaction. With increasing temperature increases kinetic energy particles and the number of active particles increases, therefore, the rate of chemical processes increases.

The dependence of the reaction rate on temperature is determined van't Hoff's rule : for every 10 0 C increase in temperature, the rate of a chemical reaction increases by two to four times.

,

Where V 1 is the reaction rate at the initial temperature of the system t 1 , V 2 is the reaction rate at the final temperature of the system t 2 ,

γ is the temperature coefficient of the reaction, equal to 2÷4.

Knowing the value of the temperature coefficient γ makes it possible to calculate the change in the reaction rate with increasing temperature from T 1 to T 2. In this case, you can use the formula:

Obviously, as the temperature rises, arithmetic progression the reaction rate increases with geometric progression. The effect of temperature on the reaction rate is the greater, the greater the value of the reaction temperature coefficient g.

It should be noted that the van't Hoff rule is approximate and is applicable only for an approximate assessment of the effect of small temperature changes on the reaction rate.

The energy required for the reactions to proceed can be provided by various influences (heat, light, electricity, laser radiation, plasma, radiation, high pressure, etc.). Reactions can be divided into thermal, photochemical, electrochemical, radiation-chemical, etc. With all these effects, the proportion of active molecules increases, which have an energy equal to or greater than the minimum energy required for this interaction E min.

When active molecules collide, the so-called activated complex , within which the redistribution of atoms takes place.

The energy required to increase the energy of the molecules of the reacting substances to the energy of the activated complex is called the activation energy Ea.

The activation energy can be considered as some additional energy that the reactant molecules must acquire in order to overcome a certain energy barrier . Thus, E a ra on the difference between the average energy of the reacting particles E ref and the energy of the activated complex E min. The activation energy is determined by the nature of the reactants. Meaning E a ranges from 0 to 400 kJ. If the value E a exceeds 150 kJ, then such reactions do not occur at temperatures close to the standard.

The change in the energy of a system during a reaction can be graphically represented using the following energy diagram (Fig. 6.1).

Activated complex reaction path

Rice. 6.1. Energy diagram of an exothermic reaction:

E ref is the average energy of the initial substances; E prod is the average energy of the reaction products; E min is the energy of the activated complex; E act - activation energy; Δ H p is the thermal effect of a chemical reaction

According to the Arrhenius equation, the higher the value of activation energy E act, the more the rate constant of a chemical reaction k temperature dependent:

,

Where E- activation energy (J/mol), R is the universal gas constant, T- temperature in K, A- Arrhenius constant, e= 2.718 is the base of natural logarithms.

Catalysts- These are substances that increase the rate of a chemical reaction. They interact with reagents to form an intermediate chemical compound and are released at the end of the reaction. The effect that catalysts have on chemical reactions is called catalysis.

For example, a mixture of aluminum powder and crystalline iodine at room temperature shows no noticeable signs of interaction, but a drop of water is enough to cause a violent reaction:

2Al + 3J 2 2AlJ 3 .

Distinguish homogeneous catalysis (the catalyst forms a homogeneous system with the reactants, for example, a gas mixture) and heterogeneous catalysis (the catalyst and the reactants are in different phases and the catalytic process takes place at the interface).

To explain the mechanism of homogeneous catalysis, the most widely used intermediate theory (proposed by the French researcher Sabatier and developed in the works of N.D. Zelinsky). According to this theory, a slow process, such as a reaction:

in the presence of a catalyst, it proceeds rapidly, but in two stages. In the first stage of the process, an intermediate compound of one of the reactants with a catalyst is formed A…cat.

First stage:

A + kat = A... kat.

The resulting compound at the second stage forms an activated complex with another reagent [ A... kat... B], which turns into the final product AB with catalyst regeneration kat.

Second stage:

A...kat + B = = AB + kat.

The intermediate interaction of the catalyst with the reactants directs the process to a new path, characterized by a lower energy barrier. Thus, the mechanism of action of catalysts is associated with a decrease in the activation energy of the reaction due to the formation of intermediate compounds.

An example is a slow reaction:

2SO 2 + O 2 \u003d 2SO 3 slowly.

In the industrial nitrous method for producing sulfuric acid, nitric oxide (II) is used as a catalyst, which significantly speeds up the reaction:

2SO2 +O2 2SO3 fast.

In the presence of a catalyst (NO), the reaction proceeds rapidly in two steps:

O 2 + 2NO \u003d 2NO 2 fast,

2NO 2 + 2SO 2 = 2SO 3 + 2NO fast.

Solutions of acids, bases, and salts (primarily salts) often serve as homogeneous catalysts. d-elements: Cr, Mn, Fe, Co, Ni, Cu, etc.).

In the case of heterogeneous catalysis, the reactants and the catalyst form several phases. Reactions take place at the interface: usually at the interface between the solid and liquid or solid and gas phases.

For example, the synthesis of ammonia from nitrogen and hydrogen is carried out using a catalyst, which is a mixture of metallic iron with the addition of potassium and aluminum oxide.

N 2 + 3H 2 2NH 3.

Heterogeneous catalysis is widely used in oil refining processes. The catalysts are platinum, nickel, aluminum oxide, etc. The hydrogenation of vegetable oil proceeds on a nickel catalyst (nickel on kieselguhr), etc.

An example of heterogeneous catalysis is the oxidation of SO 2 to SO 3 on a V 2 O 5 catalyst in the production of sulfuric acid by the contact method.

Substances that increase the activity of a catalyst are called promoters. (or activators). In this case, the promoters themselves may not have catalytic properties.

Catalytic poisons - foreign impurities in the reaction mixture, leading to partial or complete loss of catalyst activity. Thus, traces of phosphorus and arsenic cause a rapid loss of activity in the V 2 O 5 catalyst in the oxidation of SO 2 to SO 3.

Many of the most important chemical industries, such as the production of sulfuric acid, ammonia, nitric acid, synthetic rubber, a number of polymers, etc., are carried out in the presence of catalysts.

Biochemical reactions in plant and animal organisms are accelerated biochemical catalysts - enzymes.

Sharp it is possible to slow down the course of undesirable chemical processes by adding special substances to the reaction medium - inhibitors . For example, to retard undesirable processes of corrosion destruction of metals, various methods are widely used. metal corrosion inhibitors .

Questions for self-control of theory knowledge

on the topic "Chemical kinetics"

1. What does chemical kinetics study?

2. What is commonly understood by the term "reagents"?

3. What is commonly understood by the term "reaction products"?

4. How are reversible processes indicated in chemical reactions?

5. What is commonly understood by the term "activated complex"?

6. What is an elementary reaction?

7. What reactions are considered complex?

8. What stage of reactions is called the limiting stage?

9. Define the concept of "phase"?

10. What systems are considered homogeneous?

11. What systems are considered heterogeneous?

12. Give examples of homogeneous systems.

13. Give examples of heterogeneous systems.

14. What is considered the "molecularity" of the reaction?

15. What is meant by the term "rate of a chemical reaction"?

16. Give examples of fast and slow reactions.

17. What is meant by the term "rate of a homogeneous chemical reaction"?

18. What is meant by the term "rate of a heterogeneous chemical reaction"?

19. What factors determine the rate of a chemical reaction?

20. Formulate the basic law of chemical kinetics.

21. What is the rate constant of chemical reactions?

22. On what factors does the rate constant of chemical reactions depend?

23. The concentration of what substances is not included in the kinetic equation of chemical reactions?

24. How does the rate of a chemical reaction depend on pressure?

25. How does the rate of a chemical reaction depend on temperature?

26. How is the Van't Hoff Rule formulated?

27. What is the "temperature coefficient of a chemical reaction"?

28. Define the term "activation energy".

29. Give the definition of the concept of "catalyst of a chemical reaction"?

30. What is homogeneous catalysis?

31. What is heterogeneous catalysis?

32. How is the mechanism of action of a catalyst in homogeneous catalysis explained?

33. Give examples of catalytic reactions.

34. What are enzymes?

35. What are promoters?

General chemistry: textbook / A. V. Zholnin; ed. V. A. Popkova, A. V. Zholnina. - 2012. - 400 p.: ill.

Chapter 2. FUNDAMENTALS OF THE KINETICS OF CHEMICAL REACTIONS

Chapter 2. FUNDAMENTALS OF THE KINETICS OF CHEMICAL REACTIONS

The difference between breathing and burning is only in the speed of the process.

A.-L. Lavoisier

2.1. CHEMICAL KINETICS. SUBJECT AND BASIC CONCEPTS OF CHEMICAL KINETICS. SPEED REACTION

The direction, depth and fundamental possibility of the process is judged by the magnitude of the change in free energy (ΔG ≤0). However, this value does not indicate the real possibility of the reaction occurring under these conditions.

For example, the reaction of interaction of nitrous oxide with oxygen proceeds instantly at room temperature:

At the same time, 2H 2 (g) + O 2 (g) \u003d 2H 2 O (g), Δ °G\u003d -286.8 kJ / mol - a reaction characterized by a significantly large decrease in free energy, under normal conditions, the interaction does not occur, but at 700 ° C or in the presence of a catalyst, the process proceeds instantly. Consequently, thermodynamics does not answer the question of the conditions and rate of the process. This shows the limitations of the thermodynamic approach. To describe a chemical reaction, it is also necessary to know the regularities of its course in time, which are studied by kinetics.

Kinetics is a branch of chemistry that studies the rate, mechanism of chemical reactions and the influence of various factors on them.

Depending on whether the reaction components are in one or more phases, the kinetics of homogeneous and heterogeneous reactions are distinguished. According to the reaction mechanism, they are divided into simple and complex, therefore, the kinetics of simple and complex reactions are distinguished.

The basic concept of reaction kinetics is the rate of a chemical reaction. Determining the rate of chemical reactions is of biological and national economic importance.

The rate of a chemical reaction is determined by the amount of a substance that has reacted per unit time per unit volume (in the case of homogeneous reactions, when the reactants are in the same phase) or per unit interface(in the case of heterogeneous reactions, when the reactants are in different phases).

The reaction rate is characterized by a change in the concentration of any of the initial or final reaction products as a function of time. The equation describing the dependence of the reaction rate (v) on concentration (With) reactants are called kinetic. The reaction rate is often expressed in mol/l-s, in biochemistry in mg/100 ml-s, or in mass fraction, in %/100 ml-s. Distinguish average speed reactions in a time interval and the true rate of the reaction at a certain point in time. If in the time interval t1 And t2 the concentration of one of the starting substances or reaction products is equal to c 1 and c 2, respectively, then the average reaction rate (v) in the time interval t1 And t2 can be expressed:

Since in this case we are talking about a decrease in the concentration of the starting substance, i.e. the change in the concentration of a substance is taken in this case with a minus sign (-). If the reaction rate is estimated by a change (increase) in the concentration of one of the reaction products, then with a plus sign (+):

According to equation (2.2) determine average speed chemical reaction. True (instantaneous) speed reactions are determined graphically. Build a graph of the dependence of the concentration of the starting substance or reaction product (Ca) on time (t) - the kinetic curve of the reaction of Ca - f(t) for a non-linear process (Fig. 2.1).

At any point in time (eg. t1) the true reaction rate is equal to the tangent of the slope of the tangent to the kinetic curve at the point corresponding to present moment time. According to the schedule instantaneous speed reactions will be calculated by the formula:

In biochemistry, to describe the kinetics of enzymatic reactions, Michaelis-Menten equation, which shows the dependence of the rate of the reaction catalyzed by the enzyme on the concentration of the substrate and the enzyme. The simplest kinetic scheme for which the Michaelis equation is valid: E+ S↔ES→E+ P:

Rice. 2.1. Kinetic curve

Where Vm- maximum reaction rate; K m - Michaelis constant, equal to the concentration of the substrate, at which the reaction rate is half of the maximum; S- substrate concentration.

The study of the rate of a chemical reaction provides information about its mechanism. In addition to the concentration, the reaction rate depends on the nature of the reactants, external conditions, and the presence of a catalyst.

2.2. MOLECULARITY AND ORDER OF THE REACTION. HALF-LIFE

In kinetics, chemical reactions differ in terms of molecularity and reaction order. Reaction molecularity is determined by the number of particles (atoms, molecules or ions) simultaneously participating in the elementary act of chemical transformation. One, two or three molecules can take part in the elementary act of the reaction. Impact probability more particles are very small. On this basis, monomolecular, bimolecular and trimolecular reactions are distinguished. Experimentally, the molecularity of a reaction can only be determined for elementary (simple) reactions proceeding in one stage in accordance with the stoichiometric equation. Most of these reactions require a large activation energy (150-450 kJ/mol).

Most reactions are complex. The set of elementary steps that make up a complex reaction is called reaction mechanism

tions. Therefore, to characterize the reaction kinetics, the concept is introduced reaction order, which is determined by the stoichiometric equation.

The sum of the stoichiometric indicators of all the initial substances included in the reaction equation (2.5) (a+ b), defines general order reactions. The indicator with which this reagent enters the equation is called the order of the reaction with respect to the substance (particular order of the reaction), for example, the indicator A- reaction order for substance A, b- for substance B. The reaction order and molecularity are the same only for simple reactions. The order of the reaction is determined by those substances that affect the rate of the reaction.

Monomolecular reactions include decomposition and isomerization reactions.

Reactions whose rate equation includes the concentration of one reactant to the first power are called first-order reactions.

The kinetic equation includes substances whose concentration changes during the reaction. The concentrations of substances that are in significant excess do not change during the reaction.

Water in the sodium carbonate hydrolysis reaction is in significant excess and is not included in the kinetic equation.

In heterogeneous systems, the collision of particles occurs at the interface, so the mass of the solid phase does not affect the reaction rate and is therefore not taken into account in the expression for the reaction rate.

Bimolecular reactions include dimerization reactions and substitution reactions that proceed through the stage activated complex.

Reactions whose rate is proportional to the product of the concentrations of two substances to the first power or the square of the concentration of one substance are called second-order reactions.

Trimolecular reactions are rare, and four-molecular reactions are not known.

Third-order reactions do not occur among biochemical processes.

Reactions whose rate does not depend on the concentration of the starting substances are called zero-order reactions (v = k).

An example of zero-order reactions is catalytic reactions, the rate of which depends only on the concentration of the catalyst. Enzymatic reactions are a special case of such reactions.

As a rule, several reagents (substrate, coenzyme, cofactor) are involved in biochemical processes. Sometimes not all of them are known. Therefore, the course of the process is judged by one substance. In this case, the quantitative characteristic of the course of reactions in time is half-life (time) reagent - the time during which the amount or concentration of the starting substance is halved (by 50%) or half of the reaction products are formed. In this way, in particular, the decay of radionuclides is characterized, since their half-life does not depend on the initial amount.

By analyzing the dependence of the half-life of the reaction on the initial concentration, it is possible to determine the order of the reaction (the Ostwald-Noyes method). The constancy of the half-life (at a given temperature) characterizes many decomposition reactions and, in general, first-order reactions. As the reagent concentration increases, the half-life decreases for second-order reactions and increases for zero-order reactions.

2.3. REACTION RATE CONSTANT, ITS DEFINITION. LAW OF MASS ACTION

The rate of homogeneous reactions depends on the number of encounters of reacting particles per unit time per unit volume. The probability of collision of interacting particles is proportional to the product of the concentrations of the reacting substances. Thus, the reaction rate is directly proportional to the product of the concentrations of the reactants, taken in powers equal to the stoichiometric coefficients of the corresponding substances in the reaction equation. This pattern is called law of acting masses(the law of the rate of a chemical reaction), which is

fundamental law of chemical kinetics. The law of mass action was established by the Norwegian scientists K. Guldberg and P. Wage in 1867.

For example, for a reaction proceeding in a general form, according to the scheme

the kinetic equation will be valid:

Where v- the rate of a chemical reaction; with A And with B- concentration of substances A And IN[mol/l]; v a And vb- order indicators for reagents A and B; k- rate constant of a chemical reaction - a coefficient that does not depend on the concentration of reactants.

Chemical reaction rate constant (k) is the rate of a chemical reaction under conditions where the product of the concentrations of the reactants is 1 mol/L. In this case v = k.

For example, if in the reaction H 2 (g) + I 2 (g) \u003d 2НI (g) c (H 2) and c (I 2) are equal to 1 mol / l or if c (H 2) is 2 mol / l , and c(I 2) 0.5 mol/l, then v= k.

The units of the equilibrium constant are determined by the stoichiometry of the reaction. It is incorrect to compare the rate constants of reactions of different orders with each other, since they are quantities that are different in meaning and have different dimensions.

2.4. MECHANISM OF CHEMICAL REACTIONS. CLASSIFICATION OF COMPLEX REACTIONS

The reaction mechanism considers all collisions of individual particles that occur simultaneously or sequentially. The mechanism gives a detailed stoichiometric picture of each reaction step, i.e. understanding the mechanism means establishing the molecularity of each reaction step. Studying the mechanism of chemical reactions is a very difficult task. After all, we cannot conduct direct observations of the course of the interaction of molecules. The results obtained sometimes depend on the size and shape of the vessel. In some cases, the same results can be explained using different mechanisms.

The reaction of gaseous hydrogen with iodine H 2 (g) + I 2 (g) \u003d 2HI (g) was considered a classic example of a bimolecular reaction of the second

order, but in 1967 N.N. Semenov, G. Eyring and J. Sullivan showed that it has a complex character and consists of 3 elementary reactions: I 2 = 2I; 2I = I 2 ; 2I + H 2 = 2HI. Although the reaction can formally be classified as a trimolecular reaction, its rate is described by a kinetic equation resembling a second-order reaction equation:

In complex reactions, the molecularity and order of the reaction, as a rule, do not match. An unusual - fractional or negative - order of the reaction clearly indicates its complex mechanism.

The kinetic equation for the reaction of carbon monoxide oxidation with oxygen 2CO (g) + O 2 (g) \u003d CO 2 (g) has a negative (minus the first) order in CO:

as the carbon monoxide concentration increases, the reaction rate decreases.

According to the mechanism of the reaction, it can be divided into several types.

successive reactions call complex reactions, in each of which the product (X 1) of the first elementary stage reacts with the product of the second stage, the product (X 2) of the second stage enters the third, etc., until the final product is formed:

Where S- substrate (initial reagent); k 1 , k 2 , k 3 ... - rate constant 1, 2, etc. reaction steps; P- final product.

The stages of successive reactions proceed at different rates. The stage with the lowest rate constant is called the limiting stage. It determines the kinetic regularity of the reaction as a whole. Substances formed in intermediate stages are called intermediate products or intermediates which are the substrates of subsequent stages. If an intermediate is slowly formed and quickly decomposes, then its concentration does not change for a long time. Almost all metabolic processes are sequential reactions (for example, glucose metabolism).

Parallel reactions are reactions that have the same initial reagents, which correspond to different products. WITH the rate of parallel reactions is equal to the sum of the rates of individual reactions. This rule also applies to bimolecular parallel chemical reactions.

Series-parallel reactions called reactions that have the same initial reagents that can react in two ways (mechanisms) or more, including with a different number of intermediate stages. This case underlies the phenomenon catalysis, when the intermediate of one of the paths will increase the speed of other paths.

Competing reactions called complex reactions in which the same substance A reacts simultaneously with one or more reagents B 1 , B 2 etc., participates in simultaneously occurring reactions: A+ B 1 → X 1; A+ B 2 → X 2 . These reactions compete with each other for the reactant A.

Conjugated reactions complex reactions are called in which one reaction occurs only in the presence of another. In coupled reactions, the intermediate serves as a connecting link between the primary and secondary processes and causes both to occur.

A living cell needs energy for its existence. The universal source of energy in living organisms is adenosine triphosphoric acid (ATP). This compound performs the function of an energy accumulator, since when it interacts with water, i.e. hydrolysis, adenosine diphosphoric (ADP) and phosphoric (P) acids are formed and energy is released. That is why ATP is called macro-ergic compound, and bursting during its hydrolysis R-O-R connection- macroergic. macroergic bond called chemical bond, upon rupture of which, as a result of the hydrolysis reaction, significant energy is released:

As is known, the breaking of any bond (including macroergic) always requires the expenditure of energy. In the case of ATP hydrolysis, in addition to the process of breaking the bond between phosphate groups, for which Δ G>0, the processes of hydration, isomerization and neutralization of the products formed during hydrolysis occur. As a result of all these processes, the total change in the Gibbs energy has a negative

meaning. Consequently, it is not the breaking of the bond that is macroergic, but the energy result of its hydrolysis.

In order for endergonic reactions to occur in living systems (ΔG > 0), it is necessary that they be coupled with exergonic reactions (ΔG<0). Такое сопряжение возможно, если обе реакции имеют какое-либо общее промежуточное соединение, и на всех стадиях сопряженных реакций суммарный процесс характеризуется отрицательным значением изменения энергии Гиббса (∑ΔG сопр.р <0). Например, синтез сахарозы является эндэргонической реакцией и самопроизвольно происходить не может:

However, the conjugation of this reaction with the exergonic reaction of ATP hydrolysis, accompanied by the formation of a common intermediate compound glucose-1-phosphate, leads to the fact that the overall process has ∑ΔG<0:

chain reactions are called chemical and nuclear reactions in which the appearance of an active particle (a free radical or atom in chemical processes, a neutron in nuclear processes) causes a large number (chain) of successive transformations of inactive molecules or nuclei. Chain reactions are common in chemistry. Many photochemical reactions, oxidation processes (combustion, explosion), polymerization, cracking proceed according to the chain mechanism. The theory of chain reactions was developed by academician H.H. Semenov, S.N. Hinshelwood (England) and others. The main stages of chain reactions are: nucleation (initiation), continuation (elongation) and chain termination (termination). There are two types of chain reactions: straight chain reactions and branched chain reactions. A feature of chain reactions is that one primary act of activation leads to the transformation of a huge number of molecules of the starting substances. Biochemical reactions of free radical oxidation are chain reactions.

Periodic (self-oscillatory) reactions called complex multi-stage autocatalytic reactions involving several substances, in which there is a periodic fluctuation in the concentrations of the oxidized and reduced forms. Vibrational reactions were discovered by B.P. Belousov, studied by A.M. Zhabotinsky and others. The frequency and form of oscillations depend on the concentrations of the starting substances, acids

ness, temperature. An example of such reactions can be the interaction of bromomalonic acid with potassium bromate in an acidic medium, the catalyst is a salt of cerium (III). Periodic reactions are of great importance for biological objects, where reactions of this kind are widespread.

Solid phase combustion reactions(reactions of self-propagating high-temperature synthesis, SHS) were discovered in 1967 at the Institute of Chemical Physics of the USSR Academy of Sciences by A.G. Merzhanov and I.G. Borovinskaya. The essence of the SHS method lies in the fact that after local initiation of the interaction reaction of reagents, the combustion reaction front spontaneously propagates throughout the system due to heat transfer from hot products to the starting materials, initiating the interaction reaction in them. Thus, the combustion process is carried out, which is both the cause and the consequence of the reaction. The mechanism of SHS reactions is quite complex and includes the processes reaction diffusion. The term "reactive diffusion" defines a set of phenomena that occur during the interaction of two chemically different components capable of forming chemical compounds in the form of solid phases. The products of chemical interaction form a continuous layer, which differs in its structure from the initial components, but does not interfere with further interaction.

2.5. THE THEORY OF ACTIVE IMPACTS. ENERGY OF ACTIVATION. DEPENDENCE OF THE REACTION RATE ON THE NATURE OF REACTING SUBSTANCES AND TEMPERATURE

In order for an elementary act of chemical interaction to take place, the reacting particles must collide with each other. However, not every collision results in a chemical interaction. The latter occurs when the particles approach at distances at which the redistribution of the electron density and the emergence of new chemical bonds are possible. Interacting particles must have enough energy to overcome the repulsive forces that arise between their electron shells.

transition state- the state of the system, in which the destruction and creation of a connection are balanced. In a transitional state, the system

stays for a short (10 -15 s) time. The energy required to bring the system into a transition state is called activation energy. In multistep reactions that include several transition states, the activation energy corresponds to the highest energy value. After overcoming the transition state, the molecules fly apart again with the destruction of old bonds and the formation of new ones or with the transformation of the original bonds. Both options are possible, as they occur with the release of energy. There are substances that can reduce the activation energy for a given reaction.

Active molecules A 2 and B 2 upon collision combine into an intermediate active complex A 2 ... B 2 with weakening and then breaking of the A-A and B-B bonds and strengthening of the A-B bonds.

The "activation energy" of the HI formation reaction (168 kJ/mol) is much less than the energy required to completely break the bond in the initial H 2 and I 2 molecules (571 kJ/mol). Therefore, the reaction path through the formation active (activated) complex energetically more favorable than the path through the complete breaking of bonds in the original molecules. The vast majority of reactions occur through the formation of intermediate active complexes. The provisions of the active complex theory were developed by G. Eyring and M. Polyani in the 30s of the XX century.

Activation energy represents the excess of the kinetic energy of the particles relative to the average energy required for the chemical transformation of the colliding particles. Reactions are characterized by different values ​​of activation energy (E a). In most cases, the activation energy of chemical reactions between neutral molecules ranges from 80 to 240 kJ/mol. For biochemical processes, the values ​​of E and are often lower - up to 20 kJ / mol. This is explained by the fact that the vast majority of biochemical processes proceed through the stage of enzyme-substrate complexes. Energy barriers limit the reaction. Due to this, in principle, possible reactions (at G<0) практически всегда не протекают

or slow down. Reactions with an activation energy above 120 kJ/mol are so slow that they are difficult to see.

In order for a reaction to occur, the molecules must be oriented in a certain way and have sufficient energy upon collision. The probability of proper orientation in a collision is characterized by activation entropyΔ S a . The redistribution of the electron density in the active complex is favored by the condition that, upon collision, the molecules A 2 and B 2 are oriented, as shown in Fig. 2.2, a, while with the orientation shown in Fig. 2.2, b, the probability of a reaction is still much less - in fig. 2.2, c.

Rice. 2.2. Favorable (a) and unfavorable (b, c) orientations of A 2 molecules

and B 2 on collision

The equation characterizing the dependence of the rate and reaction on temperature, activation energy and activation entropy has the form:

Where k- reaction rate constant; A - in the first approximation, the total number of collisions between molecules per unit time (second) per unit volume; e is the base of natural logarithms; R- universal gas constant; T- absolute temperature; E a- activation energy; Δ S a- change in entropy of activation.

Equation (2.8) was derived by Arrhenius in 1889. The pre-exponential factor A is proportional to the total number of collisions between molecules per unit time. Its dimension coincides with the dimension of the rate constant and, therefore, depends on the total order of the reaction. The exponent is equal to the fraction of active collisions from their total number, i.e. the colliding molecules must have enough

exact interaction energy. The probability of their desired orientation at the moment of impact is proportional to e ΔSa/R

When discussing the law of mass action for velocity (2.6), it was specially stipulated that the rate constant is a constant value that does not depend on the concentrations of reagents. It was assumed that all chemical transformations proceed at a constant temperature. At the same time, it is well known that the rate of chemical transformation can change significantly with a decrease or increase in temperature. From the point of view of the law of mass action, this change in velocity is due to the temperature dependence of the rate constant, since the concentrations of reactants change only slightly due to thermal expansion or contraction of the liquid.

The most well known fact is that the rate of reactions increases with increasing temperature. This type of temperature dependence of the velocity is called normal (Fig. 2.3, a). This type of dependence is characteristic of all simple reactions.

Rice. 2.3. Types of temperature dependence of the rate of chemical reactions: a - normal; b - abnormal; c - enzymatic

However, chemical transformations are now well known, the rate of which decreases with increasing temperature. An example is the gas-phase reaction of nitrogen (II) oxide with bromine (Fig. 2.3, b). This type of temperature dependence of the velocity is called anomalous.

Of particular interest to physicians is the temperature dependence of the rate of enzymatic reactions, i.e. reactions involving enzymes. Almost all reactions occurring in the body belong to this class. For example, during the decomposition of hydrogen peroxide in the presence of the enzyme catalase, the rate of decomposition depends on temperature. In the range 273–320 °K, the temperature dependence has a normal character. As the temperature increases, the speed increases, and as the temperature decreases, it decreases. When the temperature rises above

320 °K, a sharp anomalous drop in the peroxide decomposition rate is observed. A similar picture takes place for other enzymatic reactions (Fig. 2.3, c).

From the Arrhenius equation for k it is clear that, since T included in the exponent, the rate of a chemical reaction is very sensitive to changes in temperature. The dependence of the rate of a homogeneous reaction on temperature can be expressed by the van't Hoff rule, according to which with an increase in temperature for every 10 °, the reaction rate increases by 2-4 times; the number showing how many times the rate of a given reaction increases with an increase in temperature by 10 ° is called reaction rate temperature coefficient- γ.

Where k- rate constant at temperature t°C Knowing the value of γ, one can calculate the change in the reaction rate with a change in temperature from T1 before T2 according to the formula:

As the temperature rises in an arithmetic progression, the speed increases exponentially.

For example, if γ = 2.9, then with an increase in temperature by 100 ° the reaction rate increases by a factor of 2.9 10, i.e. 40 thousand times. Deviations from this rule are biochemical reactions, the rate of which increases tenfold with a slight increase in temperature. This rule is valid only in a rough approximation. Reactions involving large molecules (proteins) are characterized by a large temperature coefficient. The rate of protein denaturation (ovalbumin) increases 50 times with a temperature increase of 10 °C. After reaching a certain maximum (50-60 °C), the reaction rate decreases sharply as a result of thermal denaturation of the protein.

For many chemical reactions, the law of mass action for velocity is unknown. In such cases, the following expression can be used to describe the temperature dependence of the conversion rate:

pre-exponent A with does not depend on temperature, but depends on concentration. The unit of measure is mol/l s.

The theoretical dependence makes it possible to pre-calculate the velocity at any temperature if the activation energy and the pre-exponential are known. Thus, the effect of temperature on the rate of chemical transformation is predicted.

2.6. REVERSIBLE AND IRREVERSIBLE REACTIONS. STATE OF CHEMICAL EQUILIBRIUM. REACTION ISOTHERM EQUATION

A chemical reaction does not always "come to an end", in other words, the starting materials are not always completely converted into reaction products. This is because as the reaction products accumulate, conditions can be created for the reaction to proceed in the opposite direction. Indeed, if, for example, iodine vapor is mixed with hydrogen at a temperature of ~200 ° C, then the following reaction will occur: H 2 + I 2 = 2HI. However, it is known that hydrogen iodine, even when heated to 180 ° C, begins to decompose into iodine and hydrogen: 2HI \u003d H 2 + I 2.

Chemical reactions that can go in opposite directions under the same conditions are called reversible. When writing the equations of reversible reactions, two oppositely directed arrows are put instead of the equal sign. A reaction proceeding from left to right is called straight(forward reaction rate constant k1), from right to left - reverse(reverse reaction rate constant k2).

In reversible reactions, the rate of the direct reaction initially has a maximum value, and then decreases due to a decrease in the concentration of the starting substances. Conversely, the reverse reaction at the initial moment has a minimum rate, which increases as the concentration of the reaction products increases. Finally, there comes a moment when the rates of the forward and reverse reactions become equal. The state in which the rate of the reverse reaction becomes equal to the rate of the forward reaction is called chemical balance.

The state of chemical equilibrium of reversible processes is quantitatively characterized by equilibrium constant. At the moment the state of chemical equilibrium is reached, the rates of the forward and reverse reactions are equal to (kinetic condition).

where K - equilibrium constant, which is the ratio of the rate constants of the forward and reverse reactions.

On the right side of the equation are those concentrations of interacting substances that are established at equilibrium - equilibrium concentrations. This equation is a mathematical expression of the law of mass action in chemical equilibrium. It should be especially noted that, in contrast to the law of mass action for the reaction rate in this equation, the exponents a, b, d, f and etc. are always equal to the stoichiometric coefficients in the equilibrium reaction.

The numerical value of the equilibrium constant of a given reaction determines its yield. Reaction yield called the ratio of the amount of product actually obtained to the amount that would have been obtained if the reaction had proceeded to the end (usually expressed as a percentage). So, at K >> 1, the reaction yield is high, and vice versa, at K<<1 выход реакции очень мал.

The equilibrium constant is related to standard Gibbs energy reactions by the following ratio:

Using equation (2.12), one can find the value of the Gibbs energy of the reaction in terms of equilibrium concentrations:

This equation is called chemical reaction isotherm equation. It allows you to calculate the change in the Gibbs energy during the course of the process and determine the direction of the reaction:

at ∆G<0 - реакция идет в прямом направлении, слева направо;

At ΔG = 0 - the reaction has reached equilibrium (thermodynamic condition);

when ΔG>0 - the reaction goes in the opposite direction.

It is important to understand that the equilibrium constant does not depend on the concentrations of substances. The converse statement is true: in a state of equilibrium, the concentrations themselves take on such values ​​that the ratio of their products in powers of stoichiometric coefficients

is constant at a given temperature. This statement corresponds to the law of mass action and can even be used as one of its formulations.

As mentioned above, reversible reactions do not proceed to the end. However, if one of the products of a reversible reaction leaves the reaction sphere, then the essentially reversible process proceeds almost to the end. If electrolytes are involved in a reversible reaction and one of the products of this reaction is a weak electrolyte, precipitate or gas, then in this case the reaction also proceeds almost to the end. irreversible reactions called such reactions, the products of which do not interact with each other with the formation of starting substances. Irreversible reactions, as a rule, "reach the end", i.e. until the complete consumption of at least one of the starting substances.

2.7. LE CHATELIER PRINCIPLE

The state of chemical equilibrium under constant external conditions can theoretically be maintained indefinitely. In reality, with a change in temperature, pressure or concentration of reagents, the equilibrium can “shift” in one direction or another of the process.

Changes occurring in the system as a result of external influences are determined by the principle of mobile equilibrium - Le Chatelier's principle.

An external impact on a system that is in a state of equilibrium leads to a shift in this equilibrium in the direction in which the effect of the produced impact is weakened.

With regard to the three main types of external influence - changes in concentration, pressure and temperature - Le Chatelier's principle is interpreted as follows.

With an increase in the concentration of one of the reacting substances, the equilibrium shifts towards the consumption of this substance, with a decrease in concentration, the equilibrium shifts towards the formation of this substance.

The influence of pressure is very similar to the effect of changing the concentrations of reactants, but it affects only gas systems. Let us formulate a general proposition on the effect of pressure on chemical equilibrium.

With an increase in pressure, the equilibrium shifts towards a decrease in the amount of gaseous substances, i.e. in the direction of decreasing pressure; when the pressure decreases, the equilibrium shifts in the direction of increasing

quantities of gaseous substances, i.e. towards increasing pressure. If the reaction proceeds without changing the number of molecules of gaseous substances, then the pressure does not affect the equilibrium position in this system.

When the temperature changes, both the forward and reverse reactions change, but to different degrees. Therefore, to clarify the effect of temperature on chemical equilibrium, it is necessary to know the sign of the thermal effect of the reaction.

As the temperature rises, the equilibrium shifts towards an endothermic reaction, and as the temperature decreases, it shifts towards an exothermic reaction.

As applied to biosystems, Le Chatelier's principle states that in a biosystem, for each action, a counteraction of the same strength and nature is formed, which balances biological regulatory processes and reactions and forms an associated level of their disequilibrium.

In pathological processes, the existing closedness of the regulatory circuit is violated. Depending on the level of disequilibrium, the quality of intersystem and interorgan relations changes, they become more and more non-linear. The structure and specificity of these relationships is confirmed by the analysis of the relationship between the indicators of the lipid peroxidation system and the level of antioxidants, between harmonic indicators in conditions of adaptation and pathology. These systems are involved in maintaining antioxidant homeostasis.

2.8. QUESTIONS AND TASKS FOR SELF-CHECKING OF PREPAREDNESS FOR LESSONS AND EXAMS

1. What reactions are called homogeneous and which are heterogeneous? Give one example of each type of reaction.

2. What reactions are called simple and which are complex? Give two examples of simple and complex reactions.

3. In what case can the molecularity and the order of the kinetic equation coincide numerically?

4. The speed of some reaction does not change over time. Will the half-life of this reaction change over time, and if so, how? Give an explanation.

5. In what case can the true (instantaneous) rate and the average reaction rate (in a sufficiently large time interval) coincide?

6. Calculate the rate constant of the reaction A + B → AB, if at concentrations of substances A and B equal to 0.5 and 0.1 mol/l, respectively, its rate is 0.005 mol/l min.

7. The half-life of some first-order reaction is 30 minutes. What part of the original amount of the substance will remain after an hour?

8. Give the concept of the general order of the reaction and the order of the reaction by substance.

9.Methods for determining the reaction rate.

10.Basic law of chemical kinetics.

11. Give the concept of the mechanism of chemical reactions.

12. Simple and complex reactions.

13. Conjugated reactions. What factors affect the rate constant of a chemical reaction?

14. Is the reaction rate really proportional to the product of the concentrations of the reactants to the power of their stoichiometric coefficients?

15. What experimental data are required to determine the order of reactions?

16. Write the kinetic equation for the reaction H 2 O 2 + 2HI → I 2 + + 2H 2 O if equal volumes of 0.02 mol / l H 2 O 2 solution and 0.05 mol / l HI solution are mixed. Rate constant 0.05 l/mol s.

17. Write the kinetic equation for the reaction H 2 O 2 + 2HI → I 2 + + 2H 2 O, given that it is characterized by the first order of the reaction in terms of the concentrations of both starting substances.

18. Prove that the rate of a chemical reaction is maximum at a stoichiometric ratio of components.

19. List possible explanations for the effect of temperature on the reaction rate.

2.9. TESTS

1. According to the van't Hoff rule, with an increase in temperature by 10 °, the rate of many reactions:

a) decreases by 2-4 times;

b) decreases by 5-10 times;

c) increases by 2-4 times;

d) increases by 5-10 times.

2. The number of elementary acts of interaction per unit of time determines:

a) the order of the reaction;

b) reaction rate;

c) molecularity of the reaction;

d) half-life.

3. What factors increase the rate of a reaction?

a) the nature of the reactants;

b) temperature, concentration, catalyst;

c) only a catalyst;

d) only concentration;

d) temperature only.

4. How many times will the reaction rate 2A(g) + B(g) increase?→A 2 B (g) with an increase in the concentration of substance A by 2 times?

a) the speed will not change;

b) will increase 18 times;

c) increase by 8 times;

d) will increase by 4 times;

d) doubled.

5. Elementary reaction A(tv) + 2B(g)→AB 2 (d). Indicate the correct kinetic equation for this reaction:

a)k[A][B] 2 ;

b)k[A][B];

c) to [B];

d) to [B] 2;

e) to [A].

6. How to change the pressure in the system in order to increase the reaction rate A (tv) + 2B (g)→AB 2 (d) 9 times?

a) increase the pressure by 9 times;

b) reduce the pressure by 9 times;

c) increase the pressure by 3 times;

d) reduce the pressure by 3 times.

7. What is the temperature coefficient of the reactionγ 10 , if, when the reaction mixture is cooled by 30 °, the reaction rate decreases by 8 times?

a) 16;

b) 8;

at 6;

d) 4;

D 2.

8. Which reaction is faster?

A) E act= 40 kJ/mol;

b) E act = 80 kJ/mol;

V) E act = 160 kJ/mol;

G) E act \u003d 200 kJ / mol.