Systems. But this value does not reflect the real possibility of the reaction occurring, its speed and mechanism.

To fully understand a chemical reaction, you must have knowledge of what time patterns exist during its implementation, i.e. rate of chemical reaction and its detailed mechanism. Study the speed and mechanism of the reaction chemical kinetics- the science of the chemical process.

From the point of view of chemical kinetics, reactions can be classified into simple and complex.

Simple reactions– processes occurring without the formation of intermediate compounds. According to the number of particles taking part in it, they are divided into monomolecular, bimolecular, trimolecular. The collision of more than 3 particles is unlikely, so trimolecular reactions are quite rare, and four-molecular reactions are unknown. Complex reactions– processes consisting of several elementary reactions.

Any process proceeds at its inherent speed, which can be determined by the changes that occur over a certain period of time. Average rate of chemical reaction expressed by changing the amount of substance n substance consumed or received per unit volume V per unit time t.

υ = ± dn/ dt· V

If a substance is consumed, then we put a “-” sign; if it accumulates, we put a “+” sign.

At constant volume:

υ = ± dC/ dt,

Reaction rate unit mol/l s

In general, υ is a constant value and does not depend on which substance involved in the reaction we are monitoring.

The dependence of the concentration of a reagent or product on the reaction time is presented in the form kinetic curve, which looks like:

It is more convenient to calculate υ from experimental data if the above expressions are transformed into the following expression:

Law of mass action. Order and rate constant of reaction

One of the formulations law of mass action sounds like this: The rate of an elementary homogeneous chemical reaction is directly proportional to the product of the concentrations of the reactants.

If the process under study is represented in the form:

a A + b B = products

then the rate of a chemical reaction can be expressed kinetic equation:

υ = k [A] a [B] b or

υ = k·C a A ·C b B

Here [ A] And [B] (C A AndC B) - concentrations of reagents,

a andb– stoichiometric coefficients of a simple reaction,

k– reaction rate constant.

Chemical meaning of quantity k- This speed reaction at single concentrations. That is, if the concentrations of substances A and B are equal to 1, then υ = k.

It must be taken into account that in complex chemical processes the coefficients a andb do not coincide with stoichiometric ones.

The law of mass action is satisfied if a number of conditions are met:

• The reaction is activated thermally, i.e. energy of thermal motion.
• The concentration of reagents is distributed evenly.
• The properties and conditions of the environment do not change during the process.
• The properties of the environment should not affect k.

To complex processes law of mass action cannot be applied. This can be explained by the fact that a complex process consists of several elementary stages, and its speed will not be determined by the total speed of all stages, but only by one slowest stage, which is called limiting.

Each reaction has its own order. Define private (partial) order by reagent and general (full) order. For example, in expressing the rate of a chemical reaction for a process

a A + b B = products

υ = k·[ A] a·[ B] b

a– order by reagent A

b — order by reagent IN

General procedure a + b = n

For simple processes the reaction order indicates the number of reacting species (coincides with the stoichiometric coefficients) and takes integer values. For complex processes the order of the reaction does not coincide with the stoichiometric coefficients and can be any.

Let us determine the factors influencing the rate of chemical reaction υ.

1. Dependence of the reaction rate on the concentration of reactants

   determined by the law of mass action: υ = k[ A] a·[ B] b

It is obvious that with increasing concentrations of reactants, υ increases, because the number of collisions between substances participating in the chemical process increases. Moreover, it is important to take into account the order of the reaction: if it n=1 for some reagent, then its speed is directly proportional to the concentration of this substance. If for any reagent n=2, then doubling its concentration will lead to an increase in the reaction rate by 2 2 = 4 times, and increasing the concentration by 3 times will speed up the reaction by 3 2 = 9 times.

The rate of a chemical reaction is the change in the concentration of reactants per unit time.

In homogeneous reactions, the reaction space refers to the volume of the reaction vessel, and in heterogeneous reactions, the surface on which the reaction occurs. The concentration of reacting substances is usually expressed in mol/l - the number of moles of a substance in 1 liter of solution.

The rate of a chemical reaction depends on the nature of the reactants, concentration, temperature, pressure, contact surface of the substances and its nature, and the presence of catalysts.

An increase in the concentration of substances that enter into a chemical interaction leads to an increase in the rate of the chemical reaction. This happens because all chemical reactions take place between a certain number of reacting particles (atoms, molecules, ions). The more of these particles there are in the volume of the reaction space, the more often they collide and chemical interaction occurs. A chemical reaction can occur through one or more elementary acts (collisions). Based on the reaction equation, we can write down an expression for the dependence of the reaction rate on the concentration of the reactants. If only one molecule participates in an elementary act (during a decomposition reaction), the dependence will have the following form:

v= k*[A]

This is the equation for a monomolecular reaction. When two different molecules interact in an elementary act, the dependence has the form:

v= k*[A]*[B]

The reaction is called bimolecular. In the case of a collision of three molecules, the expression is valid:

v= k*[A]*[B]*[C]

The reaction is called trimolecular. Coefficient designations:

v — speed reaction;

[A], [B], [C] are the concentrations of reacting substances;

k—proportionality coefficient; called the reaction rate constant.

If the concentrations of the reactants are equal to one (1 mol/l) or their product is equal to one, then v = k.. The rate constant depends on the nature of the reactants and on the temperature. The dependence of the rate of simple reactions (i.e. reactions occurring through one elementary act) on concentration is described by the law of mass action: the rate of a chemical reaction is directly proportional to the product of the concentration of reactants raised to the power of their stoichiometric coefficients.

For example, let's look at the reaction 2NO + O 2 = 2NO 2.

In it v= k* 2 *

In the case when the equation of a chemical reaction does not correspond to the elementary act of interaction, but reflects only the relationship between the mass of the substances that reacted and the substances formed, then the powers of the concentrations will not be equal to the coefficients appearing in front of the formulas of the corresponding substances in the reaction equation. For a reaction that occurs in several stages, the rate of the reaction is determined by the rate of the slowest (limiting) stage.

This dependence of the reaction rate on the concentration of reactants is valid for gases and reactions taking place in solution. Reactions involving solids do not obey the law of mass action, since the interaction of molecules occurs only at the interface. Consequently, the rate of a heterogeneous reaction also depends on the size and nature of the contact surface of the reacting phases. The larger the surface, the faster the reaction will occur.

The effect of temperature on the rate of a chemical reaction

The effect of temperature on the rate of a chemical reaction is determined by the Van't Hoff rule: with an increase in temperature for every 10 ° C, the reaction rate increases 2-4 times. Mathematically, this rule is expressed by the following equation:

v t2= v t1*g(t2-t1)/10

Where v t1 And v t2 — reaction rates at temperatures t2 and t1; g - temperature coefficient of reaction - a number showing how many times the reaction rate increases with an increase in temperature for every 10 ° C. Such a significant dependence of the rate of a chemical reaction on temperature is explained by the fact that the formation of new substances does not occur with every collision of reacting molecules. Only those molecules (active molecules) interact that have sufficient energy to break the bonds in the original particles. Therefore, each reaction is characterized by an energy barrier. To overcome it, the molecule needs activation energy - some excess energy that a molecule must have in order for its collision with another molecule to lead to the formation of a new substance. With increasing temperature, the number of active molecules increases rapidly, which leads to a sharp increase in the reaction rate according to Van't Hoff's rule. The activation energy for each specific reaction depends on the nature of the reactants.

Active collision theory allows us to explain the influence of certain factors on the rate of a chemical reaction. The main provisions of this theory:

• Reactions occur when particles of reactants that have a certain energy collide.
• The more reactant particles there are, the closer they are to each other, the more likely they are to collide and react.
• Only effective collisions lead to a reaction, i.e. those in which “old connections” are destroyed or weakened and therefore “new ones” can be formed. To do this, the particles must have sufficient energy.
• The minimum excess energy required for effective collision of reactant particles is called activation energy Ea.
• The activity of chemicals is manifested in the low activation energy of reactions involving them. The lower the activation energy, the higher the reaction rate. For example, in reactions between cations and anions, the activation energy is very low, so such reactions occur almost instantly

Catalyst influence

One of the most effective means of influencing the rate of chemical reactions is the use of catalysts. TO atalizers - These are substances that change the rate of a reaction, but at the end of the process they themselves remain unchanged in composition and mass. In other words, at the moment of the reaction itself, the catalyst is actively involved in the chemical process, but by the end of the reaction, the reactants change their chemical composition, turning into products, and the catalyst is released in its original form. Typically, the role of a catalyst is to increase the rate of a reaction, although some catalysts slow down the process rather than speed it up. The phenomenon of acceleration of chemical reactions due to the presence of catalysts is called catalysis, and slowdowns - inhibition.

Some substances do not have a catalytic effect, but their additions dramatically increase the catalytic ability of catalysts. Such substances are called promoters. Other substances (catalytic poisons) reduce or even completely block the action of catalysts, this process is called catalyst poisoning.

There are two types of catalysis: homogeneous And heterogeneous. At homogeneous catalysis the reactants, products and catalyst form one phase (gas or liquid). In this case, there is no interface between the catalyst and the reactants.

Peculiarity heterogeneous catalysis is that catalysts (usually solids) are in a different phase state than the reactants and products of the reaction. The reaction usually develops on the surface of a solid.

In homogeneous catalysis, intermediate products are formed between the catalyst and the reactant as a result of a reaction with a lower activation energy. In heterogeneous catalysis, the increase in rate is explained by the adsorption of reactants on the surface of the catalyst. As a result, their concentration increases and the reaction rate increases.

A special case of catalysis is autocatalysis. Its meaning is that a chemical process is accelerated by one of the reaction products.

Chemical reaction rate- change in the amount of one of the reacting substances per unit of time in a unit of reaction space.

The speed of a chemical reaction is influenced by the following factors:

• the nature of the reacting substances;
• concentration of reactants;
• contact surface of reacting substances (in heterogeneous reactions);
• temperature;
• action of catalysts.

Active collision theory allows us to explain the influence of certain factors on the rate of a chemical reaction. The main provisions of this theory:

• Reactions occur when particles of reactants that have a certain energy collide.
• The more reactant particles there are, the closer they are to each other, the more likely they are to collide and react.
• Only effective collisions lead to a reaction, i.e. those in which “old connections” are destroyed or weakened and therefore “new ones” can be formed. To do this, the particles must have sufficient energy.
• The minimum excess energy required for effective collision of reactant particles is called activation energy Ea.
• The activity of chemicals is manifested in the low activation energy of reactions involving them. The lower the activation energy, the higher the reaction rate. For example, in reactions between cations and anions, the activation energy is very low, so such reactions occur almost instantly

The influence of the concentration of reactants on the reaction rate

As the concentration of reactants increases, the reaction rate increases. In order for a reaction to occur, two chemical particles must come together, so the rate of the reaction depends on the number of collisions between them. An increase in the number of particles in a given volume leads to more frequent collisions and an increase in the reaction rate.

An increase in the rate of reaction occurring in the gas phase will result from an increase in pressure or a decrease in the volume occupied by the mixture.

Based on experimental data in 1867, Norwegian scientists K. Guldberg and P. Waage, and independently of them in 1865, Russian scientist N.I. Beketov formulated the basic law of chemical kinetics, establishing dependence of the reaction rate on the concentrations of the reactants -

Law of mass action (LMA):

The rate of a chemical reaction is proportional to the product of the concentrations of the reacting substances, taken in powers equal to their coefficients in the reaction equation. (“effective mass” is a synonym for the modern concept of “concentration”)

aA +bB =cС +dD, Where k– reaction rate constant

ZDM is performed only for elementary chemical reactions occurring in one stage. If a reaction proceeds sequentially through several stages, then the total speed of the entire process is determined by its slowest part.

Expressions for the rates of various types of reactions

ZDM refers to homogeneous reactions. If the reaction is heterogeneous (reagents are in different states of aggregation), then the ZDM equation includes only liquid or only gaseous reagents, and solid ones are excluded, affecting only the rate constant k.

Molecularity of the reaction is the minimum number of molecules involved in an elementary chemical process. Based on molecularity, elementary chemical reactions are divided into molecular (A →) and bimolecular (A + B →); trimolecular reactions are extremely rare.

Rate of heterogeneous reactions

• Depends on surface area of ​​contact between substances, i.e. on the degree of grinding of substances and the completeness of mixing of reagents.
• An example is wood burning. A whole log burns relatively slowly in air. If you increase the surface of contact between wood and air, splitting the log into chips, the burning rate will increase.
• Pyrophoric iron is poured onto a sheet of filter paper. During the fall, the iron particles become hot and set fire to the paper.

Effect of temperature on reaction rate

In the 19th century, the Dutch scientist Van't Hoff experimentally discovered that with an increase in temperature by 10 o C, the rates of many reactions increase by 2-4 times.

Van't Hoff's rule

For every 10 ◦ C increase in temperature, the reaction rate increases by 2-4 times.

Here γ (Greek letter "gamma") - the so-called temperature coefficient or van't Hoff coefficient, takes values ​​from 2 to 4.

For each specific reaction, the temperature coefficient is determined experimentally. It shows exactly how many times the rate of a given chemical reaction (and its rate constant) increases with every 10 degree increase in temperature.

Van't Hoff's rule is used to approximate the change in the reaction rate constant with increasing or decreasing temperature. A more precise relationship between the rate constant and temperature was established by the Swedish chemist Svante Arrhenius:

How more E a specific reaction, so less(at a given temperature) will be the rate constant k (and rate) of this reaction. An increase in T leads to an increase in the rate constant, this is explained by the fact that an increase in temperature leads to a rapid increase in the number of “energetic” molecules capable of overcoming the activation barrier Ea.

Effect of catalyst on reaction rate

You can change the rate of a reaction by using special substances that change the reaction mechanism and direct it along an energetically more favorable path with a lower activation energy.

Catalysts- these are substances that participate in a chemical reaction and increase its speed, but at the end of the reaction they remain unchanged qualitatively and quantitatively.

Inhibitors– substances that slow down chemical reactions.

Changing the rate of a chemical reaction or its direction using a catalyst is called catalysis .

Kinetics– the science of the rates of chemical reactions.

Chemical reaction rate– the number of elementary acts of chemical interaction occurring per unit time per unit volume (homogeneous) or per unit surface (heterogeneous).

True reaction speed:

For homogeneous, heterogeneous reactions:

1) concentration of reacting substances;

2) temperature;

3) catalyst;

4) inhibitor.

Only for heterogeneous:

1) the rate of supply of reacting substances to the phase interface;

2) surface area.

The main factor is the nature of the reactants - the nature of the bonds between atoms in the molecules of the reactants.

NO 2 – nitrogen oxide (IV) – fox tail, CO – carbon monoxide, carbon monoxide.

If they are oxidized with oxygen, then in the first case the reaction will occur instantly, as soon as you open the cap of the vessel, in the second case the reaction is extended over time.

The concentration of reactants will be discussed below.

Blue opalescence indicates the moment of sulfur precipitation; the higher the concentration, the higher the speed.

Rice. 10

The higher the concentration of Na 2 S 2 O 3, the less time the reaction takes. The graph (Fig. 10) shows a directly proportional relationship. The quantitative dependence of the reaction rate on the concentration of the reacting substances is expressed by the LMA (law of mass action), which states: the rate of a chemical reaction is directly proportional to the product of the concentrations of the reacting substances.

So, basic law of kinetics is an empirically established law: the rate of a reaction is proportional to the concentration of the reactants, example: (i.e. for a reaction)

For this reaction H 2 + J 2 = 2HJ – the rate can be expressed in terms of a change in the concentration of any of the substances. If the reaction proceeds from left to right, then the concentration of H 2 and J 2 will decrease, and the concentration of HJ will increase as the reaction progresses. For the instantaneous reaction rate, we can write the expression:

square brackets indicate concentration.

Physical meaning k– molecules are in continuous motion, collide, fly apart, and hit the walls of the vessel. In order for the chemical reaction to form HJ to occur, the H2 and J2 molecules must collide. The number of such collisions will be greater, the more molecules of H 2 and J 2 are contained in the volume, i.e., the greater the values ​​[H 2 ] and . But the molecules move at different speeds, and the total kinetic energy of the two colliding molecules will be different. If the fastest molecules H 2 and J 2 collide, their energy can be so high that the molecules break into atoms of iodine and hydrogen, which fly apart and then interact with other molecules H 2 + J 2 > 2H+2J, then H + J 2 > HJ + J. If the energy of the colliding molecules is less, but high enough to weaken the H – H and J – J bonds, the formation reaction of hydrogen iodide will occur:

For most colliding molecules, the energy is less than that required to weaken the bonds in H 2 and J 2. Such molecules will “quietly” collide and also “quietly” disperse, remaining what they were, H 2 and J 2. Thus, not all, but only part of the collisions lead to a chemical reaction. The proportionality coefficient (k) shows the number of effective collisions leading to a collision reaction at concentrations [H 2 ] = 1 mol. Magnitude k–const speed. How can speed be constant? Yes, the speed of uniform rectilinear motion is a constant vector quantity equal to the ratio of the movement of a body over any period of time to the value of this interval. But molecules move chaotically, then how can the speed be const? But a constant speed can only be at a constant temperature. With increasing temperature, the proportion of fast molecules whose collisions lead to a reaction increases, i.e., the rate constant increases. But the increase in the rate constant is not unlimited. At a certain temperature, the energy of the molecules will become so great that almost all collisions of the reactants will be effective. When two fast molecules collide, a reverse reaction will occur.

There will come a moment when the rates of formation of 2HJ from H 2 and J 2 and decomposition will be equal, but this is already a chemical equilibrium. The dependence of the reaction rate on the concentration of the reactants can be traced using the traditional reaction of interaction of a solution of sodium thiosulfate with a solution of sulfuric acid.

Na 2 S 2 O 3 + H 2 SO 4 = Na 2 SO 4 + H 2 S 2 O 3, (1)

H 2 S 2 O 3 = Sv+H 2 O+SO 2 ^. (2)

Reaction (1) occurs almost instantly. The rate of reaction (2) depends at a constant temperature on the concentration of the reactant H 2 S 2 O 3. This is exactly the reaction we observed - in this case, the speed is measured by the time from the beginning of the solutions to merge until the appearance of opalescence. In the article L. M. Kuznetsova The reaction of sodium thiosulfate with hydrochloric acid is described. She writes that when solutions are drained, opalescence (turbidity) occurs. But this statement by L.M. Kuznetsova is erroneous since opalescence and turbidity are two different things. Opalescence (from opal and Latin escentia– suffix meaning weak effect) – scattering of light by turbid media due to their optical inhomogeneity. Light scattering– deviation of light rays propagating in a medium in all directions from the original direction. Colloidal particles are capable of scattering light (Tyndall-Faraday effect) - this explains opalescence, a slight turbidity of the colloidal solution. When carrying out this experiment, it is necessary to take into account the blue opalescence, and then the coagulation of the colloidal suspension of sulfur. The same density of the suspension is noted by the visible disappearance of any pattern (for example, a grid on the bottom of a cup) observed from above through the layer of solution. Time is counted using a stopwatch from the moment of draining.

Solutions of Na 2 S 2 O 3 x 5H 2 O and H 2 SO 4.

The first is prepared by dissolving 7.5 g of salt in 100 ml of H 2 O, which corresponds to a 0.3 M concentration. To prepare a solution of H 2 SO 4 of the same concentration, you need to measure 1.8 ml of H 2 SO 4 (k), ? = = 1.84 g/cm 3 and dissolve it in 120 ml of H 2 O. Pour the prepared Na 2 S 2 O 3 solution into three glasses: 60 ml in the first, 30 ml in the second, 10 ml in the third. Add 30 ml of distilled H 2 O to the second glass, and 50 ml to the third glass. Thus, in all three glasses there will be 60 ml of liquid, but in the first the salt concentration is conditionally = 1, in the second – ½, and in the third – 1/6. After the solutions have been prepared, pour 60 ml of H 2 SO 4 solution into the first glass with a salt solution and turn on the stopwatch, etc. Considering that the reaction rate decreases with dilution of the Na 2 S 2 O 3 solution, it can be determined as a quantity inversely proportional to time v = 1/? and construct a graph, plotting the concentration on the abscissa axis, and the reaction rate on the ordinate axis. The conclusion from this is that the reaction rate depends on the concentration of substances. The data obtained are listed in Table 3. This experiment can be performed using burettes, but this requires a lot of practice from the performer, because the graph may be incorrect.

Table 3

Speed ​​and reaction time

The Guldberg-Waage law is confirmed - professor of chemistry Gulderg and young scientist Waage).

Let's consider the next factor - temperature.

As temperature increases, the rate of most chemical reactions increases. This dependence is described by Van't Hoff's rule: “For every 10 °C increase in temperature, the rate of chemical reactions increases by 2 to 4 times.”

Where ? – temperature coefficient showing how many times the reaction rate increases when the temperature increases by 10 °C;

v 1 – reaction rate at temperature t 1 ;

v 2 – reaction rate at temperature t2.

For example, a reaction at 50 °C takes two minutes, how long will it take for the process to complete at 70 °C if the temperature coefficient ? = 2?

t 1 = 120 s = 2 min; t 1 = 50 °C; t 2 = 70 °C.

Even a slight increase in temperature causes a sharp increase in the reaction rate of active collisions of the molecule. According to activation theory, only those molecules whose energy is greater than the average energy of molecules by a certain amount participate in the process. This excess energy is activation energy. Its physical meaning is the energy that is necessary for the active collision of molecules (rearrangement of orbitals). The number of active particles, and therefore the reaction rate, increases with temperature according to an exponential law, according to the Arrhenius equation, which reflects the dependence of the rate constant on temperature

Where A - Arrhenius proportionality coefficient;

k– Boltzmann's constant;

E A – activation energy;

R – gas constant;

T- temperature.

A catalyst is a substance that accelerates the rate of a reaction without being consumed.

Catalysis– the phenomenon of changing the reaction rate in the presence of a catalyst. There are homogeneous and heterogeneous catalysis. Homogeneous– if the reagents and the catalyst are in the same state of aggregation. Heterogeneous– if the reagents and catalyst are in different states of aggregation. About catalysis, see separately (further).

Inhibitor– a substance that slows down the rate of reaction.

The next factor is surface area. The larger the surface area of ​​the reactant, the greater the speed. Let us consider, using an example, the effect of the degree of dispersion on the reaction rate.

CaCO 3 – marble. Dip the tiled marble into hydrochloric acid HCl, wait five minutes, it will dissolve completely.

Powdered marble - we will do the same procedure with it, it will dissolve in thirty seconds.

The equation for both processes is the same.

CaCO 3 (s) + HCl (g) = CaCl 2 (s) + H 2 O (l) + CO 2 (g) ^.

So, when adding powdered marble, the time is less than when adding slab marble, for the same mass.

With an increase in the interface surface, the rate of heterogeneous reactions increases.

The rate of a chemical reaction depends on the following factors:

1) The nature of the reacting substances.

2) The contact surface of the reagents.

3) Concentration of reactants.

4) Temperature.

5) Presence of catalysts.

The rate of heterogeneous reactions also depends on:

a) the size of the phase interface (with an increase in the phase interface, the rate of heterogeneous reactions increases);

b) the rate of supply of reacting substances to the phase interface and the rate of removal of reaction products from it.

Factors affecting the rate of a chemical reaction:

1. Nature of the reagents. The nature of the chemical bonds in compounds and the structure of their molecules play an important role. For example, the release of hydrogen by zinc from a solution of hydrochloric acid occurs much faster than from a solution of acetic acid, since the polarity of the H-C1 bond is greater than the O-H bond in the CH 3 COOH molecule, in other words, due to the fact that HCl - is a strong electrolyte, and CH 3 COOH is a weak electrolyte in aqueous solution.

2. Contact surface of reagents. The larger the contact surface of the reacting substances, the faster the reaction proceeds. The surface area of ​​solids can be increased by grinding them, and for soluble substances by dissolving them. Reactions in solutions occur almost instantly.

3. Concentration of reagents. For interaction to occur, particles of reacting substances in a homogeneous system must collide. When increasing concentrations of reactants the speed of reactions increases. This is explained by the fact that as the amount of substance per unit volume increases, the number of collisions between particles of reacting substances increases. The number of collisions is proportional to the number of particles of reacting substances in the volume of the reactor, i.e., their molar concentrations.

The quantitative dependence of the reaction rate on the concentration of reactants is expressed law of mass action (Guldberg and Waage, Norway, 1867): the rate of a chemical reaction is proportional to the product of the concentrations of the reacting substances.

For reaction:

aA + bB ↔ cC + dD

the reaction rate in accordance with the law of mass action is equal to:

υ = k[A]υ a ·[B]υ b,(9)

where [A] and [B] are the concentrations of the starting substances;

k-reaction rate constant, which is equal to the reaction rate at the concentrations of the reactants [A] = [B] = 1 mol/l.

The reaction rate constant depends on the nature of the reactants, temperature, but does not depend on the concentration of substances.

Expression (9) is called kinetic equation of the reaction. The kinetic equations include the concentrations of gaseous and dissolved substances, but do not include the concentrations of solid substances:

2SO 2 (g) + O 2 (g) = 2SO 3 (g); υ = k 2 · [O 2 ];

CuO (tv.) + H 2 (g) = Cu (tv.) + H 2 O (g); υ = k.

Using kinetic equations, you can calculate how the reaction rate changes when the concentration of reactants changes.

The influence of the catalyst.

5. Reaction temperature. Active collision theory

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 reaction. Chemical interaction occurs when particles approach distances at which redistribution of electron density and the formation of new chemical bonds are possible. The interacting particles must have sufficient energy to overcome the repulsive forces that arise between their electron shells.

Transition state- a state of the system in which the destruction and creation of connections are balanced. The system remains in a transition state for a short time (10–15 s). The energy that must be expended 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 scatter 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 reaction for the formation of НI (168 kJ/mol) is significantly less than the energy required to completely break the bond in the initial molecules of Н2 and I2 (571 kJ/mol). Therefore, the reaction path through the formation active (activated) complex energetically more favorable than the path through the complete rupture of bonds in the original molecules. The vast majority of reactions occur through the formation of intermediate active complexes. The principles of the theory of the active complex were developed by G. Eyring and M. Polyani in the 30s of the 20th century.

Activation energy represents the excess kinetic energy of particles relative to the average energy required for the chemical transformation of colliding particles. Reactions are characterized by different activation energies (E a). In most cases, the activation energy of chemical reactions between neutral molecules ranges from 80 to 240 kJ/mol. For biochemical processes values E a 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 (with Q< 0) практически всегда не протекают или замедляются. Реакции с энергией активации выше 120 кДж/моль настолько медленны, что их протекание трудно заметить.

For a reaction to occur, the molecules must be oriented in a certain way and have sufficient energy when they collide. The probability of proper collision orientation is characterized by activation entropy ∆S a. The redistribution of electron density in the active complex is favored by the condition when, upon collision, molecules A 2 and B 2 are oriented, as shown in Fig. 3a, whereas with the orientation shown in Fig. 3b, the probability of reaction is even much less - in Fig. 3c.

Rice. 3. Favorable (a) and unfavorable (b, c) orientations of molecules A 2 and B 2 upon collision

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

(10)

Where k- reaction rate constant;

A- to a first approximation, the total number of collisions between molecules per unit of time (second) per unit volume;

e- the base of natural logarithms;

R- universal gas constant;

T- absolute temperature;

E a- activation energy;

∆S a- change in activation entropy.

Equation (11) was derived by Arrhenius in 1889. Pre-exponential factor A proportional to the total number of collisions between molecules per unit time. Its dimension coincides with the dimension of the rate constant and depends on the total order of the reaction.

Exhibitor equal to the proportion of active collisions from their total number, i.e. colliding molecules must have sufficient interaction energy. The probability of their desired orientation at the moment of impact is proportional to .

When discussing the law of mass action for speed (9), it was specifically stated that the rate constant is a constant value that does not depend on the concentrations of the reagents. It was assumed that all chemical transformations occur at a constant temperature. At the same time, the rate of chemical transformation can change significantly with decreasing or increasing temperature. From the point of view of the law of mass action, this change in speed is due to the temperature dependence of the rate constant, since the concentrations of the reacting substances change only slightly due to thermal expansion or compression of the liquid.

The most well-known fact is that the rate of reactions increases with increasing temperature. This type of temperature dependence of speed is called normal (Fig. 3 a). This type of dependence is characteristic of all simple reactions.

Rice. 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; this type of temperature dependence of the rate is called abnormal . An example is the gas-phase reaction of nitrogen (II) oxide with bromine (Fig. 3 b).

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, when hydrogen peroxide decomposes in the presence of the enzyme catalase, the rate of decomposition depends on temperature. In the range 273-320 TO The temperature dependence is normal. As the temperature increases, the speed increases, and as the temperature decreases, it decreases. When the temperature rises above 320 TO there is a sharp anomalous drop in the rate of peroxide decomposition. A similar picture occurs for other enzymatic reactions (Fig. 3c).

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 temperature changes. The dependence of the rate of a homogeneous reaction on temperature can be expressed by the van’t Hoff rule, according to which with every 10° increase in temperature, the reaction rate increases by 2-4 times; a number showing how many times the rate of a given reaction increases with an increase in temperature by 10° is called temperature coefficient of reaction rate -γ.

This rule is expressed mathematically by the following formula:

(12)

where γ is the temperature coefficient, which shows how many times the reaction rate increases when the temperature increases by 10 0; υ 1 –t 1 ; υ 2 – reaction rate at temperature t2.

As the temperature increases in an arithmetic progression, the speed increases in a geometric progression.

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

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

Pre-exponent And with does not depend on temperature, but depends on concentration. The unit of measurement is mol/l∙s.

The theoretical dependence allows the speed to be calculated in advance at any temperature if the activation energy and pre-exponential are known. Thus, the influence of temperature on the rate of chemical transformation is predicted.

Complex reactions

The principle of independence. Everything discussed above related to relatively simple reactions, but in chemistry so-called complex reactions are often encountered. Such reactions include those discussed below. When deriving kinetic equations for these reactions, the principle of independence is used: If several reactions occur in a system, then each of them is independent of the others and its rate is proportional to the product of the concentrations of its reactants.

Parallel reactions- These are reactions that occur simultaneously in several directions.

The thermal decomposition of potassium chlorate occurs simultaneously in two reactions:

Sequential reactions- These are reactions that occur in several stages. These are the majority of reactions in chemistry.

.

Conjugate reactions. If several reactions occur in a system and the occurrence of one of them is impossible without the other, then these reactions are called conjugated , and the phenomenon itself - by induction .

2HI + H 2 CrO 4 → I 2 + Cr 2 O 3 + H 2 O.

This reaction is practically not observed under normal conditions, but if FeO is added to the system, the following reaction occurs:

FeO + H 2 CrO 4 → Fe 2 O 3 + Cr 2 O 3 + H 2 O

and at the same time the first reaction occurs. The reason for this is the formation in the second reaction of intermediate products involved in the first reaction:

FeO 2 + H 2 CrO 4 → Cr 2 O 3 + Fe 5+;

HI + Fe 5+ → Fe 2 O 3 + I 2 + H 2 O.

Chemical induction- a phenomenon in which one chemical reaction (secondary) depends on another (primary).

A+ IN- primary reaction,

A+C- secondary reaction,

then A is an activator, IN- inductor, C - acceptor.

During chemical induction, unlike catalysis, the concentrations of all reaction participants decrease.

Induction factor determined from the following equation:

.

Depending on the magnitude of the induction factor, the following cases are possible.

I > 0 - damping process. The reaction rate decreases over time.

I < 0 - ускоряющийся процесс. Скорость реакции увеличи­вается со временем.

The phenomenon of induction is important because in some cases the energy of the primary reaction can compensate for the energy consumed in the secondary reaction. For this reason, for example, it turns out to be thermodynamically possible to synthesize proteins by polycondensation of amino acids.

Chain reactions. If a chemical reaction proceeds with the formation of active particles (ions, radicals), which, entering subsequent reactions, cause the appearance of new active particles, then this sequence of reactions is called chain reaction.

The formation of free radicals is associated with the expenditure of energy to break bonds in the molecule. This energy can be imparted to molecules by illumination, electrical discharge, heating, irradiation with neutrons, α- and β-particles. To carry out chain reactions at low temperatures, initiators are introduced into the reaction mixture - substances that easily form radicals: sodium vapor, organic peroxides, iodine, etc.

The reaction of the formation of hydrogen chloride from simple compounds, activated by light.

Total reaction:

H 2 + C1 2 2HC1.

Individual stages:

Сl 2 2Сl∙ photoactivation of chlorine (initiation)

Cl∙ + H 2 = HCl + H∙ chain development

H∙ + Cl 2 = HCl + Cl∙, etc.

H∙ + Cl∙ = HCl open circuit

Here H∙ and Cl∙ are active particles (radicals).

In this reaction mechanism, three groups of elementary stages can be distinguished. The first is a photochemical reaction chain nucleation. Chlorine molecules, having absorbed a light quantum, dissociate into free atoms that are highly reactive. Thus, during the nucleation of a chain, the formation of free atoms or radicals from valence-saturated molecules occurs. The process of chain nucleation is also called initiation. Chlorine atoms, having unpaired electrons, are able to react with molecular hydrogen, forming molecules of hydrogen chloride and atomic hydrogen. Atomic hydrogen, in turn, interacts with a chlorine molecule, as a result of which a hydrogen chloride molecule and atomic chlorine are again formed, etc.

These processes, characterized by the repetition of the same elementary stages (links) and proceeding with the preservation of free radicals, lead to the consumption of starting substances and the formation of reaction products. Such groups of reactions are called reactions of development (or continuation) of the chain.

The stage of the chain reaction in which the death of free radicals occurs is called open circuit. Chain termination can occur as a result of the recombination of free radicals, if the energy released during this process can be given to some third body: the wall of the vessel or molecules of inert impurities (stages 4, 5). That is why the rate of chain reactions is very sensitive to the presence of impurities, to the shape and size of the vessel, especially at low pressures.

The number of elementary links from the moment the chain begins to break is called the chain length. In the example under consideration, up to 10 5 HCl molecules are formed for each quantum of light.

Chain reactions during which there is no “multiplying” of the number of free radicals are called unbranched or simple chain reactions . In each elementary stage of an unbranched chain process, one radical “gives birth” to one molecule of the reaction product and only one new radical (Fig. 41).

Other examples of simple chain reactions: a) chlorination of paraffin hydrocarbons Cl∙ + CH 4 → CH 3 ∙ + HC1; CH 3 ∙ + Cl - → CH 3 Cl + Cl ∙ etc.; b) radical polymerization reactions, for example, polymerization of vinyl acetate in the presence of benzoyl peroxide, which easily decomposes into radicals; c) the interaction of hydrogen with bromine, which occurs according to a mechanism similar to the reaction of chlorine with hydrogen, only with a shorter chain length due to its endothermicity.

If, as a result of the act of growth, two or more active particles appear, then this chain reaction is branched.

In 1925, N. N. Semenov and his collaborators discovered reactions containing elementary stages, as a result of which not one, but several chemically active particles - atoms or radicals - appear. The appearance of several new free radicals leads to the appearance of several new chains, i.e. one chain branches. Such processes are called branched chain reactions (Fig. 42).

An example of a highly branched chain process is the oxidation of hydrogen at low pressures and temperatures of about 900°C. The reaction mechanism can be written as follows.

1. H 2 + O 2 OH∙ + OH∙ chain initiation

2. OH∙ + H2 → H2O + H∙ chain development

3. H∙ + O 2 → OH∙ + O: chain branching

4. O: + H 2 → OH∙ +H∙

5. OH∙ +H 2 → H 2 O + H∙ continuation of the chain

6. Н∙ + Н∙ + wall → Н 2 open circuit on the wall of the vessel

7. H∙ + O 2 + M → HO 2 ∙ + M open circuit in the volume.

M is an inert molecule. The radical HO 2 ∙, formed during a triple collision, is inactive and cannot continue the chain.

At the first stage of the process, hydroxyl radicals are formed, which ensure the development of a simple chain. In the third stage, as a result of interaction with the original molecule of one radical, two radicals are formed, and the oxygen atom has two free valences. This ensures branching of the chain.

As a result of chain branching, the reaction rate rapidly increases in the initial period of time, and the process ends with a chain ignition-explosion. However, branched chain reactions end in explosion only when the rate of branching is greater than the rate of chain termination. Otherwise, the process is slow.

When the reaction conditions change (changes in pressure, temperature, mixture composition, size and condition of the walls of the reaction vessel, etc.), a transition from a slow reaction to an explosion may occur and vice versa. Thus, in chain reactions there are limiting (critical) states at which chain ignition occurs, from which thermal ignition that occurs in exothermic reactions as a result of ever-increasing heating of the reacting mixture with weak heat removal should be distinguished.

Oxidation of vapors of sulfur, phosphorus, carbon monoxide (II), carbon disulfide, etc. occurs through a branched chain mechanism.

The modern theory of chain processes was developed by Nobel Prize laureates (1956) Soviet academician N. N. Semenov and English scientist Hinshelwood.

Chain reactions should be distinguished from catalytic reactions, although the latter are also cyclic in nature. The most significant difference between chain reactions and catalytic ones is that with a chain mechanism, the reaction can flow in the direction of increasing the energy of the system due to spontaneous reactions. A catalyst does not cause a thermodynamically impossible reaction. In addition, in catalytic reactions there are no process stages such as chain nucleation and chain termination.

Polymerization reactions. A special case of a chain reaction is a polymerization reaction.

Polymerization is a process in which the reaction of active particles (radicals, ions) with low-molecular compounds (monomers) is accompanied by the sequential addition of the latter with an increase in the length of the material chain (molecule length), i.e., with the formation of a polymer.

Monomers are organic compounds, usually containing unsaturated (double, triple) bonds in the molecule.

Main stages of the polymerization process:

1. Initiation(under the influence of light, heat, etc.):

A: A→A" + A"- homolytic decomposition with the formation of radicals (active valence-unsaturated particles).

A: B→ A - + B +- heterolytic decomposition with the formation of ions.

2. Chain height: A" + M→ AM"

(or A - + M→ AM", or IN + + M→ VM +).

3. Open circuit: AM" + AM"→ polymer

(or AM" + B +→ polymer, VM + + A"→ polymer).

The speed of a chain process is always greater than that of a non-chain process.