The rate of a chemical reaction: dependence on the concentration of reactants, temperature, catalyst action. Activation energy. Chemical balance.

Chemical kinetics is the science of the mechanisms and rates of chemical reactions.

Chemical reaction rate

The rate of a chemical reaction is equal to the change in the amount of substance that reacts or is formed as a result of the reaction per unit time in a unit of reaction space. The rate of reaction is denoted by the letter V, usually expressed in moles per liter (mol/L), and the time in seconds or minutes.

The rate of a chemical reaction depends on:

1) from the nature of the reacting substances;

2) concentrations of reacting substances;

3) temperature;

4) the presence of a catalyst.

Dependence of the reaction rate on the concentration of reactants

The quantitative dependence of the reaction rate on the concentration of reactants is expressed by the law of mass action: the reaction rate is proportional to the product of the molar concentrations of all reactants, to powers equal to the stoichiometric coefficient for the corresponding reactant of the reaction equation. In general terms for a homogeneous reaction:

aA + bB = dD + fF

υ = k[A] a -[B] in or v = k С a А·С in in.

To indicate the concentrations of reactants or reaction products, it is customary to use the letter “C” or square brackets; CA, St - concentrations of substances A and B, mol/l; [A], [B] – equilibrium concentrations of substances A and B; A And V- stoichiometric coefficients for substances A and B in the reaction equation; k - proportionality coefficient, called the reaction rate constant, depends on the nature of the reactants, temperature and the presence of a catalyst.

For example, the expression for the reaction rate: 2CO (g) + O 2 (g) = 2CO 2 (g)

has the form: υ ​​= k 2 [O 2 ]

For heterogeneous reactions of the type: aA(g) + bB(k) = cC(k)

the expression for the reaction rate is: υ = kC a A or υ = k[A] a.

Dependence of the rate of a chemical reaction on temperature

The rate of reactions increases with increasing temperature. The reason for this is an increase in the energy of the colliding particles, as a result of which the likelihood that a chemical transformation will occur during a collision increases. Determined by Van't Hoff's rule: with an increase in temperature by 10°, the rate of most chemical reactions increases by 2-4 times.

The mathematical expression of Van't Hoff's rule:

υ2/ υ1 = T/10 ( - Van't Hoff coefficient)

where υ1 and υ2 are the reaction rates at temperatures T1 and T2; γ is the temperature coefficient of the reaction, showing how many times the reaction rate increases when the temperature increases by 10°.

The activation energy of the reaction E A is the threshold energy. If the energy of the colliding particles is less than E A, then during the collision the reaction will not occur; if the energy exceeds E A, the reaction will occur.

A chemical transformation occurs only when conditions arise for a redistribution of the electron density of the colliding particles. This process takes place over time and requires energy. Let's consider the interaction of gaseous substances A 2 and B 2:

A 2 (g) + B 2 (g) = 2AB (g)

The reaction path can be characterized by three successive states of the system:

A B A……B A - B

│ + │ → : : → +

A B A…....B A - B

initial state transition state final state

(initial reagents) (activated complex) (reaction products)

In the transition state, a rearrangement of atoms occurs, accompanied by a redistribution of electron density. The energy required for the transition of substances into the state of an activated complex is called the Gibbs activation energy.

It is determined by the relation

Therefore, we can similarly write the Gibbs activation energy

∆G ≠ = ∆H ≠ - T∆S ≠ ,

where ∆H ≠ is the enthalpy of activation of the reaction; T – temperature; ∆S ≠ - entropy of reaction activation.

The formation of an activated complex requires energy. The probability that when particles collide, an activated complex is formed and a reaction occurs depends on the energy of the colliding particles. Only those molecules whose energy is sufficient for this react. Such molecules are called active. The energy required for the transition of substances into the state of an activated complex is called the enthalpy of activation ∆H ≠.

Solving typical problems.

Example 1. How will the rate of interaction of the starting substances change when the temperature increases from 20 to 66°C if the temperature coefficient of the reaction is 2.5?

Solution. According to the conditions of the problem, the change in temperature T2 - T1 == 66 - 20 = 46°. Therefore, as a result of an increase in temperature by 46°, the ratio v 2 / v 1 = T/10 = 4.6 log2.5 = 4.6 0.398 = 1.831, then υ66/ υ20 = 67.7. The reaction rate increases by 67.7 times.

Dependence of reaction rate on catalyst

In the presence of a catalyst, the rate of a chemical reaction changes. The rate increases in the presence of some catalysts, and slows down in the presence of others.

A catalyst is a substance that participates in a reaction and changes its rate, but remains unchanged after the chemical reaction ends. A catalyst that slows down a chemical reaction is called an inhibitor. Biological catalysts of protein nature are called enzymes.

The mechanism of action of catalysts is due to the fact that they form intermediate compounds with the starting substances and thereby change the reaction path, and the new reaction path is characterized by a lower energy barrier, that is, a lower activation energy compared to a non-catalyzed reaction.

Chemical equilibrium

Reactions that occur simultaneously in two opposite directions (forward and reverse) are called reversible. Generally speaking, there are no irreversible reactions. It’s just that under certain conditions, some reactions can be brought almost to completion, for example, if products are removed from the reaction sphere - precipitation, gas evolution, formation of slightly dissociated products, etc. For any reversible homogeneous reaction:

aA + bB ↔ cC -dD

at the initial moment of time, according to the law of mass action, the rate of the forward reaction: υ = k ·C a A-C b B , has a maximum value, and the rate of the reverse reaction υ = k -C c C-C d D is equal to zero. Over time, the concentration of the starting substances - reactants (A and B) decreases, and the reaction products (C and D) increase and, therefore, the rate of the forward reaction decreases and the rate of the reverse reaction increases. There comes a moment when both speeds become equal, which corresponds to the equilibrium state of the system.

The concentrations of reagents and reaction products established at the moment of equilibrium are called equilibrium [A], [B], [C], [D], they remain constant until chemical equilibrium is disrupted. The equilibrium state of a chemical system is characterized according to the law of mass action by the equilibrium constant (Kp), for the reaction

aA + bB « dD + fF.

This expression makes it possible to calculate Kp from the known equilibrium concentrations of all substances of a homogeneous reaction or the concentration of an individual substance from the known concentrations of other substances and KP. For the same temperature, the ratio of the products of equilibrium concentrations (in powers of their stoichiometric coefficients) of substances in the right and left sides the equation of a chemical reaction represents a constant value. The equilibrium constant shows the depth of the process. If K>>1, the process is strongly shifted towards the production of reaction products. If K<<1, наоборот, процесс сдвинут влево и практически не идет. К=1 - равновесие установилось.

When the reaction proceeds in the forward direction to a state of equilibrium, the concentration of the reactants decreases by the values ​​ΔС A and ΔС B and the concentrations of the products increase by the values ​​ΔС C and ΔС D, determined by the expressions for the reagents:

ΔС А = С 0(А) - [А], ΔС в = С 0(в) - [В], ΔС с = С 0(С) + [С] = 0 + [С] = [С], ΔС D = C 0(D) + [D] = 0 + [D] = [D],

where C 0 (A), C 0, (B), C 0, (C), C 0, (D) are the initial concentrations of reagents and reaction products.

The solution of the problem

Example 1. At a certain temperature, the equilibrium constant of the reaction: H 2 (g) + I 2 (g) ↔ 2HI (g) is equal to 1. Determine the composition of the equilibrium reaction mixture if 1 mol/l H 2 and 2 mol/l I 2 were taken for the reaction .

Solution. The problem comes down to determining the equilibrium concentrations of reactants and reaction products through the equilibrium constant. Equilibrium concentrations are the concentrations of reactants that did not react at the time of equilibrium, and the concentrations of reaction products formed at the time of equilibrium. These concentrations can be calculated from the reaction equation:

H 2 (g) +I 2 (g) ↔2HI (g).

Initial concentration: 1 2 0

At the moment of equilibrium:

1) reacted, ∆С x x

2) 1 2 left

3) 2x were formed

Thus, the equilibrium concentrations of the starting materials and reaction products are:

C H2 -x = (l-x),

C I 2 - x = (2-x),

2x, since it is clear from the equation that HI is formed 2 times more than H 2 or I 2 reacts. C 0, (H2) and C 0, (I 2) are the initial concentrations of H 2 and I 2. Upon reaching equilibrium, the composition of the reaction mixture was as follows:

[H 2 ] = (1 - 0.45) = 0.55 mol/l,

= (2 - 0.45) = 1.55 mol/l,

2·0.45 = 0.9 mol/l.

Chemical equilibrium shift

Each chemical equilibrium is established at a certain value of three parameters that characterize it: 1) the concentration of the reacting substances; 2) temperature; 3) pressure (for gases). A change in one of these parameters leads to an imbalance: (υ≠ υ). If υ> υ, then the equilibrium shifts to the right, in the direction of the formation of reaction products, which is denoted by (→). If υ< υ, то равновесие смещается влево (←), в направлении образования исходных веществ.

The direction of the equilibrium shift is determined by Le Chatelier's principle: if an external influence is exerted on a system in a state of equilibrium, then the equilibrium will shift in the direction that will weaken the external influence.

1. If an external influence on the system manifests itself in a decrease in the concentration of one of the substances participating in the reaction, then this shifts the equilibrium towards its formation. When the concentration of one of the substances increases, the equilibrium of the system shifts towards the reaction that reduces it.

2. An increase in temperature shifts the equilibrium towards an endothermic reaction (∆H > 0), and a decrease - towards an exothermic reaction (∆H< 0).

3. A change in pressure affects the equilibrium if at least one gaseous substance is involved in the reaction and the number of moles of the initial gaseous substances and gaseous reaction products is not the same. As the pressure decreases or increases, the equilibrium will mix, respectively, towards the formation of more or less moles of gas.

Example 1. Under what conditions is the reaction equilibrium:

4Fe (k) + 3O 2 (g) ↔2Fe 2 O 3 (k), ΔH 0 r = -1644.4 kJ

will shift towards the decomposition of the oxide?

Solution. 1. A shift in equilibrium towards the decomposition of the oxide means a shift to the left, i.e. increasing the rate of the reverse reaction, which is endothermic. The direct reaction is exothermic (ΔН 0 r< О). Такое смещение, согласно принципу Ле-Шателье, достигается повышением температуры.

2. The given reversible reaction is heterogeneous. It involves one gaseous substance - oxygen, which is the starting material. To shift the equilibrium in the direction of O 2 formation (←), its concentration must be reduced, which is equivalent to a decrease in pressure in the system.

Lecture 8. Solutions

Types of solutions, thermodynamics of dissolution. Solubility. Dissolution of gases in liquids. Henry's Law. Supersaturated solutions. Raoult's law. Colligative properties of solutions. Electrolytes. Arrhenius' theory of electrolytic dissociation. Degree of electrolytic dissociation. Features of electrolyte solutions. Degree and constant of dissociation. Weak electrolytes. Ionic product of water. Hydrogen index. Product of solubility. Hydrolysis of salts. Various cases of hydrolysis. Degree and constant of hydrolysis. Hydrolysis shift.

A molecular or true solution is a homogeneous system consisting of two or more components. Colloidal solutions should be distinguished from molecular solutions: suspensions, emulsions, aerosols. Colloidal solutions differ from molecular solutions in that they are multicomponent heterogeneous systems. Examples of molecular solutions are an aqueous solution of sodium chloride, air, and an alloy of silver and gold. It is clear from the examples given that the types of solutions can be different.

Solutions can be in three states of aggregation: gaseous, liquid, solid. Therefore, a distinction is made between liquid solutions, gas solutions, and solid solutions. A solution consists of a solute and a solvent. A solvent is considered to be that component that is in the same state of aggregation as the solution itself. If all components are in the same state, then the solvent is considered to be the component that is in greater quantity.

A solution that is in equilibrium with the solute is called a saturated solution. In such solutions, at a given temperature, a larger amount of solute cannot dissolve. A saturated solution is in dynamic equilibrium with the insoluble part of the solute.

Solubility The ability of a substance to dissolve in a particular solvent is called. A measure of the solubility of a substance is its content in a saturated solution under certain conditions. Numerically, solubility is expressed in the same ways as composition. For example, the percentage of the mass of a solute to the mass of a saturated solution, or the amount of solute contained in 1 liter of a saturated solution. Sometimes the solubility coefficient is used to characterize solubility. The solubility coefficient is the number of units of mass of an anhydrous substance that saturates 100 units of mass of solvent under given conditions.

Typically, substances consisting of polar molecules and substances with an ionic type of bond are better dissolved in polar solvents (water, alcohols, liquid ammonia), and non-polar substances are better dissolved in non-polar solvents (benzene, carbon disulfide). This confirms the rule of thumb “like dissolves like.”

Solubility depends on temperature. For some substances this dependence is weak. For example, the solubility of potassium, lead, and silver nitrates (KNO 3, Pb(NO 3) 2, AgNO 3) in water increases sharply with increasing temperature. The solubility of sodium chloride (NaCl) in water changes only slightly as the temperature increases.

When solids are dissolved in water, the volume of the system usually changes slightly. Therefore, the solubility of solids in water is practically independent of pressure.

Liquids can dissolve in each other. Some of them, for example, alcohol - water, are unlimitedly soluble, that is, they can be mixed with each other in any proportions. There are liquids that are mutually soluble only to a certain limit; they are called partially miscible. If you shake diethyl ether with water, two layers are formed: the upper one is a saturated solution of water in ether, and the lower one is a saturated solution of ether in water. For such systems, with increasing temperature, the mutual solubility of liquids increases until a temperature is reached at which both liquids mix in any proportions. The temperature at which the limited mutual solubility of liquids becomes unlimited is called the critical solution temperature. Thus, at temperatures below 66.4 0 C, phenol is limitedly soluble in water, and water is limitedly soluble in phenol. For the water-phenol system, a temperature of 66.4 0 C is the critical dissolution temperature, since from this temperature and above, both liquids are infinitely soluble in each other.

The mutual dissolution of liquids is usually not accompanied by significant changes in volume, and therefore depends little on pressure. Only at very high pressures of the order of thousands of atmospheres does the mutual solubility of liquids increase significantly.

There are liquids that are completely insoluble in each other; they are called immiscible liquids. If both liquids are insoluble in one another, then when mixed in any proportions, two separate layers are formed. Examples of such liquids are the following: mercury - water, carbon disulfide - water, chlorobenzene - water, phenylamine - water.

If a third substance is introduced into a system consisting of two immiscible liquids, capable of dissolving in each of these liquids, then the solute will be distributed between both liquids in proportion to its solubility in each of them. For such systems, the distribution law is satisfied according to which: a substance capable of dissolving in two immiscible solvents is distributed between them so that the ratio of its concentrations in these solvents at a constant temperature remains constant, regardless of the total amount of solute

K(V) = s 1V / s 2V,

where c 1B and c 2B are the concentrations of the dissolved substance in the first and second solvents, K(B) is the distribution constant of substance B between two immiscible liquids.

The dissolution of gases in water is an exothermic process. Therefore, the solubility of gases decreases with increasing temperature. When a gas dissolves in a liquid, equilibrium is established

Gas + Liquid ↔ Saturated solution

In this case, the volume of the system is significantly reduced. Consequently, an increase in pressure should lead to a shift of equilibrium to the right, that is, to an increase in gas solubility.

Henry formulated this pattern more generally: The partial vapor pressure of a solute above a solution is proportional to the mole fraction of the solute in the solution.

The solubility of most substances decreases with decreasing temperature. Therefore, when hot saturated solutions are cooled, excess solute is released. However, if cooling is carried out carefully and slowly, then there will be no release of the solute from the solution. In this case, the resulting solution contains significantly more solute than is required for saturation at a given temperature. Such solutions are called oversaturated. Such solutions in a calm state can remain unchanged for years. But if you throw a crystal of the substance that is dissolved in it into a solution, then immediately other crystals begin to grow around it and after a short time the entire excess of the dissolved substance crystallizes out. Crystallization sometimes begins from simple shaking of the solution. Supersaturated solutions of Glauber's salt (Na 2 SO 4 ∙10H 2 O), sodium thiosulfate (Na 2 S 2 O 3 ∙5H 2 O) very easily form.

Raoult's law. Colligative properties of solutions.

The colligative properties of solutions are those properties that depend only on the concentration of particles of the dissolved substance, but not on its chemical composition. The most common are the following four colligative properties of solutions:

1) decrease in steam pressure;

2) increase in boiling point;

3) lowering the freezing temperature;

4) osmotic pressure.

All these four properties apply to solutions containing non-volatile soluble substances, that is, those soluble substances whose vapor pressure is negligible.

At a given temperature, the pressure of saturated vapor above the liquid is a constant value. When a substance is dissolved in a given liquid, the saturated vapor pressure of that liquid above the solution decreases. The saturated vapor pressure of a solvent above a solution is always lower than above a pure solvent at the same temperature. The difference between these pressures is called reduction of steam above the solution.

In 1887, the French physical chemist Raoult established a law linking a decrease in vapor pressure over dilute solutions of nonelectrolytes with an increase in the concentration of the dissolved substance. It is called Raoult's law: the relative decrease in the saturated vapor pressure of the solvent above the solution is equal to the mole fraction of the solute

(P 0 – P)/P 0 = X

The consequence of a decrease in the saturated vapor pressure of a solvent above a solution is an increase in its boiling point compared to a pure solvent and a decrease in its freezing point.

Any liquid begins to boil at the same temperature at which the pressure of its saturated vapor reaches the external pressure. Water at a pressure of 101 kPa begins to boil at a temperature of 100 0 C because the saturated vapor pressure is 101 kPa. Since at a given temperature the pressure of saturated water vapor above the solution will be lower than above the pure solvent, the solution does not boil at 100 0 C. The boiling point of an aqueous solution is more than 100 0 C, and the higher the concentration of the solution, the higher it is.

When liquids freeze, crystallization begins at the temperature at which the saturated vapor pressure above the liquid phase becomes equal to the saturated vapor pressure above the solid phase. Water freezes at 0 0 C because at this temperature the saturated vapor pressure of water above the liquid and above the ice is the same and equal to 0.61 kPa.

For dilute solutions, the increase in boiling point and decrease in freezing point does not depend on the nature of the solute and is directly proportional to the amount of substance n:

Dt deputy = K k ×C m; Dt bale =K e ×C m,

where Dt dep. and Dt bal. – respectively, the decrease in the freezing point and the increase in the boiling point of the solution are found using the formula

Dt deputy = T deputy. r-la – T deputy. solution; Dt kip = T kip. solution – T boil. r-la;

K k and K e are the cryoscopic and ebullioscopic constants of the solvent, respectively; С m – molal concentration of solution (mol/kg) can be found using the formula

,

where m 1 is the mass of the dissolved substance, g; M is its molar mass, g/mol; m 2 – mass of solvent, g.

Osmosis- This is the spontaneous transition of a solvent through a semi-permeable membrane from a dilute solution or pure solvent to a concentrated solution. A membrane that allows solvent particles to pass through but does not allow solute particles to pass through is called a semipermeable membrane. An example of such a membrane is the bull bladder. A semi-permeable membrane allows solvent particles to pass in both directions. But on the side of the membrane where the solution concentration is higher, the solvent concentration is lower. Therefore, the resulting transition of the solvent into a concentrated solution occurs. This leads to the establishment of a pressure difference on both sides of the membrane, which is called osmotic pressure.

Osmotic pressure is a colligative property, since it depends only on the concentration of dissolved particles and does not depend on their chemical composition. For osmotic pressure, the Van't Hoff equation is satisfied.

,

where n is the amount of dissolved substance, mol; V – solution volume, m3; R – gas constant equal to 8.31 J/(mol K); T – temperature, K; m is the mass of the dissolved substance, g; M – its molar mass, g/mol; cm - molar concentration, mol/l

Osmotic pressure plays an important role in biological systems. In animals, some cells, such as red blood cells, contain a saline solution. These cells are surrounded by a plasma membrane. In an aquatic environment, red blood cells undergo osmosis, swell and burst. However, if they are exposed to a more concentrated salt solution, the cells shrink.

If the pressure applied to a concentrated solution exceeds the osmotic pressure, then the solvent passes from the concentrated solution through the membrane into the dilute solution. This process is called reverse osmosis. It is used in industry to obtain drinking water from sea water.

Electrolyte solutions

Substances that break down into ions and conduct electric current are called electrolytes. An electrolyte conducts electric current as a result of the directional movement of its ions creating a flow of electrical charges. Thus, passing an electric current through an electrolyte is accompanied by a transfer of substance.

Electrolytes are acids, bases and salts that are in a molten state or in an aqueous solution.

The ability of electrolytes to conduct electric current is called electrolytic conductivity. It differs from the electronic conductivity of metals or other conductors of electric current. In substances with electronic conductivity, the flow of charge is due to the movement of electrons. Therefore, passing an electric current through conductors with electronic conductivity is not accompanied by the transfer of matter. Electrolytes are conductors of the second kind. In a solution or melt, they break up into ions, which is why electric current flows.

The breakdown of electrolytes into ions when dissolved in water is called electrolytic dissociation.

To explain the characteristics of aqueous solutions of electrolytes, the Swedish chemist S. Arrhenius proposed in 1887 electrolytic dissociation theory. The main provisions of the theory are as follows:

1. Electrolytes, when dissolved in water, break down into ions - positive and negative. Ions are in more stable electronic states than atoms. Among such ions there are simple ones, for example, Na +, Mg 2+, Al 3+, and complex ones, consisting of several atoms, for example, NO 3 -, SO 4 2-, PO 4 3-. In a solution, ions move randomly in different directions.

2. Under the influence of an electric current, ions acquire directional movement: positively charged ions move towards the cathode, negatively charged ions move towards the anode. The former are called cations, and the latter anions.

3. Dissociation is a reversible process. Simultaneously with the decomposition of molecules into ions, the reverse process occurs - the combination of ions into a molecule.

Therefore, in the equations of electrolytic dissociation there is not an equal sign, but a reversibility sign ↔.

Under degree of electrolyte dissociation refers to the ratio of the number of molecules dissociated into ions n to the total number of molecules of dissolved electrolyte N, that is

Depending on the degree of dissociation, weak and strong electrolytes are distinguished. Strong electrolytes at high concentrations are dissociated by more than 1/2. The degree of dissociation of weak electrolytes is very small compared to 1. Strong electrolytes– these are the majority of soluble salts (except for CuCl 2, Pb(CH 3 COO) 2, Fe(CNS) 3), alkalis and strong acids (HCI, HBr, HI, HNO 3, H2SO4, HClO 4, HMnO 4). Weak electrolytes are most organic acids, inorganic weak acids and weak bases, some neutral salts CdCl 2, Fe(CH 3 COO) 3. Particularly weak electrolytes are water, hydrogen sulfide, hydrocyanic and boric acids.

The degree of dissociation depends on the nature of the electrolyte and solvent, as well as on the concentration of the electrolyte. With decreasing concentration, the degree of dissociation increases, and with a strong dilution of the solution, A→1, and the differences between strongly and weakly dissociating electrolytes are smoothed out.

Ionic product.

The electrical conductivity of water is explained by the fact that water dissociates to a very small extent, forming hydrogen ions and hydroxide ions:

This process is equilibrium and, like any equilibrium process, it can be characterized by an equilibrium constant, which is dissociation constant:

At room temperature, only one out of 108 water molecules breaks down into ions.

In dilute solutions, the concentration of water changes very little and can be considered constant, then

Since is a constant, they enter into K D and denote To W:

This quantity is called ionic product of water and is a constant value at a given temperature.

In pure water at room temperature, the concentrations of hydrogen ions and hydroxide ions are equal to each other and equal to 10 –7 mol/l. Hence:

Equilibrium constant To W depends on temperature and does not depend on the concentration of H + cations and OH – anions.

If an acid is added to water, the concentration of hydrogen cations will increase, the equilibrium will shift to the left, and the concentration of hydroxide ions will decrease so that the ionic product of water remains unchanged.

Thus, in aqueous solutions, at a constant temperature, the concentrations of hydrogen cations and hydroxide ions are related to each other. When calculating for aqueous solutions of strong electrolytes, not concentrations are used, but activities:

Ion activity a i is expressed as the product of the concentration of an ion with i and its activity coefficient i:

and i= i with i

It is impossible to experimentally determine the activities of the a + cation and the a - anion, since they do not exist separately. Therefore, the concept of average ionic activity a is introduced. For an electrolyte producing n+ cations and n- anions

a ± = (a + n+ ∙a - n-) 1/n

where n = n + + n - .

The average ion activity coefficient γ ± is determined similarly

γ ± = (γ + n + ∙γ - n -) 1/ n

To characterize the acidity (alkalinity) of the environment, a special parameter has been introduced - hydrogen index or pH. Hydrogen index or pH is the reversed decimal logarithm of the concentration of hydrogen ions in a solution:

The pH value determines the nature of the reaction of the solution. For example, at 295K it is neutral and pH = 7 (the concentration of hydrogen ions is [H + ] = 10 –7 mol/l). At pH<7 (концентрация ионов водорода [Н + ] >10 –7 mol/l) the reaction of the solution is acidic, at pH>7 (concentration of hydrogen ions [H + ]<10 –7 моль/л) – щелочная. С изменением температуры величина ионного произведения воды To W changes.

The pH value can serve as a criterion for the strength of an acid or base. In the series of acids, the one with the higher activity of H + ions (pH lower) at the same molar concentration will be strong. For foundations, such a dependence is inverse.

Product of solubility.

In a saturated electrolyte solution, the product of the concentrations of its ions is a constant value at a given temperature. This value quantitatively characterizes the ability of an electrolyte to dissolve; it is called the solubility product

The solubility product depends on the nature of the solute and solvent, as well as on temperature and does not depend on the activity of the ions of the sparingly soluble electrolyte in solution.

The solubility products for most electrolytes are calculated and contained in tables. Knowing the solubility product, it is possible to calculate whether a substance precipitates under given conditions. The condition for the formation of a precipitate of a poorly soluble electrolyte is that the product of the activities of the ions of this electrolyte in solution exceeds the tabulated value of the solubility product.

Hydrolysis of salts

Salt hydrolysis is the exchange interaction of salt ions with water molecules, leading to an increase in the acidity or alkalinity of the solution and the formation of weakly dissociable compounds. The essence of hydrolysis reactions is the interaction of salt ions with water ions to form weak electrolytes. During the hydrolysis process, one of the water ions binds into a weak electrolyte, and the other, as a rule, accumulates in the solution. The ion that accumulates in the solution determines the reaction of the medium. If H + ions accumulate, then the environment will be acidic, if the OH - groups are alkaline. When electrolytes of equal strength are formed, the medium can be neutral.

Hydrolysis equations are written similarly to other ionic equations: slightly dissociated (including water) and slightly soluble, as well as gaseous substances are written in the form of molecules, strong electrolytes are written in the form of ions. The equations for the hydrolysis of salts of polybasic acids and polyacid bases are written in steps, similar to stepwise dissociation.

There are four cases of interaction between salt and water.

1. Salts formed by a strong acid and a weak base.

Salts of this type, when dissolved in water, form an acidic solution. An example is ammonium chloride NH 4 Cl. The equation for the hydrolysis reaction of this salt has the form

NH 4 Cl + H 2 O ↔ NH 4 OH + HCl.

In ionic form, the hydrolysis reaction equation has the form

NH 4 + + H 2 O ↔ NH 4 OH + H + .

Due to the binding of OH ions by ammonium ions into weakly dissociating NH 4 OH molecules, an excess of hydrogen ions appears in the solution and the solution becomes acidic.

2. A salt is formed by a weak acid and a strong base.

Salts of this type, when dissolved in water, form an alkaline solution. When a salt formed by a weak acid and a strong base is hydrolyzed, a weak acid and an excess of hydroxyl ions OH - are formed. An example is the hydrolysis of potassium cyanide, the reaction equation has the form

KCN + H 2 O ↔ HCN + KOH

or in ionic form

Chemical kinetics and equilibrium

Goal of the work: study of the influence of temperature on the reaction rate, concentration on the shift in chemical equilibrium.

Theoretical rationale:

Speed ​​of chemical reaction is the amount of a substance that reacts or is formed as a result of a reaction per unit time per unit volume (for homogeneous reactions) or per unit interface surface (for heterogeneous reactions).

If over a period of time?f = f 2 f 1 the concentration of one of the substances participating in the reaction decreases by?C = C2C1, then the average rate of the chemical reaction for the specified period of time is equal to

The value V expresses the rate of a chemical process over a certain period of time. Therefore, the smaller?f, the closer the average speed will be to the true one.

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

1) the nature and concentration of the reacting substances;

2) temperature of the reaction system;

3) presence of a catalyst;

4) pressure,

5) the magnitude of the phase interface and the mixing rate of the system (for heterogeneous reactions);

6) type of solvent.

Effect of reagent concentrations. The rate of a reaction is proportional to the number of collisions of molecules of the reacting substances. The number of collisions, in turn, is greater, the higher the concentration of each of the starting substances.

A general formulation of the effect of concentration on the rate of a chemical reaction is given by law of mass action(1867, Guldberg, Waage, Beketov).

At a constant temperature, the rate of a chemical reaction is proportional to the product of the concentrations of the reacting substances, taken in powers of their equalizing (stoichiometric) coefficients.

For the reaction aA + bB = cC V = K[A] a [B] b,

where K is the proportionality coefficient or speed constant;

If [A] = 1 mol/l, [B] = 1 mol/l, then V = K, hence the physical meaning

rate constants K: the rate constant is equal to the reaction rate at concentrations of reactants equal to unity.

The effect of temperature on the reaction rate. As the temperature increases, the frequency of collisions of reacting molecules increases, and therefore the reaction rate increases.

The quantitative effect of temperature on the rate of homogeneous reactions can be expressed by Van't Hoff's rule.

In accordance with Van't Hoff's rule, when the temperature increases (decreases) by 10 degrees, the rate of a chemical reaction increases (decreases) by 2-4 times:

where V (t 2 ) and V (t 1 ) - the rate of chemical reaction at appropriate temperatures; f(t 2 ) And f(t 1 ) - duration of the chemical reaction at appropriate temperatures; G - van't Hoff temperature coefficient, which can take a numerical value in the range of 2-4.

Activation energy. The excess energy that molecules must have in order for their collision to lead to the formation of a new substance is called the activation energy of a given reaction (expressed in kJ/mol). One of the methods of activation is to increase the temperature: with increasing temperature, the number of active particles increases greatly, due to which the reaction rate sharply increases.

The dependence of the reaction rate on temperature is expressed by the Arrhenius equation:

where K is the rate constant of the chemical reaction; E a - activation energy;

R - universal gas constant; A - constant; exp is the base of natural logarithms.

The magnitude of the activation energy can be determined if two values ​​of the rate constant K 1 and K 2 are known at temperatures T 1 and T 2, respectively, according to the following formula:

Chemical balance.

All chemical reactions can be divided into two groups: irreversible and reversible. Irreversible reactions proceed to completion - until one of the reactants is completely consumed, i.e. flow in only one direction. Reversible reactions do not proceed to completion. In a reversible reaction, none of the reactants are completely consumed. A reversible reaction can occur in both the forward and reverse directions.

Chemical equilibrium is a state of a system in which the rates of forward and reverse reactions are equal.

For a reversible reaction

m A+ n B? p C+ q D

the chemical equilibrium constant is

In reversible chemical reactions, equilibrium is established at the moment when the ratio of the product of concentrations of products raised to powers equal to the stoichiometric coefficients to the product of concentrations of starting substances, also raised to the corresponding powers, is equal to some constant value called the chemical equilibrium constant.

The chemical equilibrium constant depends on the nature of the reactants and on temperature. The concentrations at which equilibrium is established are called equilibrium. A change in external conditions (concentration, temperature, pressure) causes a shift in the chemical equilibrium in the system and its transition to a new equilibrium state.

Such a transition of a reaction system from one state to another is called a displacement (or shift) of chemical equilibrium.

The direction of the shift in chemical equilibrium is determined by Le Chatelier’s principle: If any external influence is applied to a system that is in a state of chemical equilibrium (change concentration, temperature, pressure), then processes spontaneously arise in this system that tend to weaken the effect produced.

An increase in the concentration of one of the starting reagents shifts the equilibrium to the right (the direct reaction increases); An increase in the concentration of reaction products shifts the equilibrium to the left (the reverse reaction intensifies).

If a reaction proceeds with an increase in the number of gas molecules (i.e., on the right side of the reaction equation, the total number of gas molecules is greater than the number of molecules of gaseous substances on the left side), then an increase in pressure prevents the reaction, and a decrease in pressure favors the reaction.

When the temperature increases, the equilibrium shifts towards the endothermic reaction, and when the temperature decreases, it shifts towards the exothermic reaction.

The catalyst changes the rate of both forward and reverse reactions by the same number of times. Therefore, the catalyst does not cause a shift in equilibrium, but only shortens or increases the time required to achieve equilibrium.

Experiment No. 1 Dependence of the speed of a homogeneous reaction on the concentration of the initial reagents.

b Instruments, equipment: test tubes, stopwatch, solutions of sodium thiosulfate (III), dil. sulfuric acid (1M), water.

b Methodology: This dependence can be studied using the classic example of a homogeneous reaction between sodium thiosulfate and sulfuric acid, proceeding according to the equation

Na 2 S 2 O 3 + H 2 SO 4 = Na 2 SO 4 + Sv + SO 2 ^ + H 2 O.

At first, sulfur forms a colloidal solution with water (barely perceptible turbidity). It is necessary to measure the time from the moment of draining until a barely noticeable turbidity appears with a stopwatch. Knowing the reaction time (in seconds), you can determine the relative speed of the reaction, i.e. reciprocal of time:

chemical homogeneous kinetics

For the experiment, you should prepare three dry, clean test tubes and number them. Add 4 drops of sodium thiosulfate solution and 8 drops of water to the first; in the second - 8 drops of sodium thiosulfate and 4 drops of water; in the third - 12 drops of sodium thiosulfate. Shake the test tubes.

If we conditionally designate the molar concentration of sodium thiosulfate in test tube 1 as “c”, then accordingly in test tube 2 there will be 2 s mole, in test tube 3 - 3 s mol.

Add one drop of sulfuric acid into test tube 1, and at the same time turn on the stopwatch: shaking the test tube, watch for the appearance of turbidity in the test tube, holding it at eye level. When the slightest cloudiness appears, stop the stopwatch, note the reaction time and write it down in the table.

Perform similar experiments with the second and third test tubes. Enter the experimental data in the laboratory journal in the form of a table.

b Conclusion: with increasing concentration of sodium thiosulfate, the rate of this reaction increases. The dependence graph is a straight line passing through the origin.

Experience No. 2. Study of the dependence of the rate of a homogeneous reaction on temperature.

b Instruments and equipment: test tubes, stopwatch, thermometer, solutions of sodium thiosulfate (III), sulfuric acid (1M)

b Methodology:

Prepare three clean, dry test tubes and number them. Add 10 drops of sodium thiosulfate solution to each of them. Place test tube No. 1 in a glass of water at room temperature and after 1...2 minutes note the temperature. Then add one drop of sulfuric acid to the test tube, simultaneously turn on the stopwatch and stop it when a weak, barely noticeable turbidity appears. Record the time in seconds from the moment the acid is added to the test tube until turbidity appears. Record the result in the table.

Then increase the temperature of the water in the glass by exactly 10 0 either by heating it on a hot plate or by mixing it with hot water. Place test tube No. 2 in this water, hold for several minutes and add one drop of sulfuric acid, turning on the stopwatch at the same time, shake the test tube with its contents in a glass of water until turbidity appears. If a barely noticeable cloudiness appears, turn off the stopwatch and enter the stopwatch readings into the table. Carry out a similar experiment with the third test tube. First increase the temperature in the beaker by another 10 0, place test tube No. 3 in it, hold for several minutes and add one drop of sulfuric acid, while turning on the stopwatch and shaking the test tube.

Express the results of the experiments in a graph, plotting speed on the ordinate axis, and temperature on the abscissa axis.

Determine the temperature coefficient of the reaction g

b Conclusion: during the experiment, the average temperature coefficient was calculated, which turned out to be equal to 1.55. Ideally it is

2-4. The deviation from the ideal can be explained by the error in measuring the time of turbidity of the solution. The graph of the reaction rate versus temperature has the form of a parabola branch that does not pass through 0. With increasing temperature, the reaction rate increases

Experiment No. 3 The influence of the concentration of reactants on chemical equilibrium.

b Instruments and equipment: test tubes, potassium chloride (crystal), solutions of iron (III) chloride, potassium thiocyanate (saturated), distilled water, cylinder

b Methodology:

A classic example of a reversible reaction is the interaction between ferric chloride and potassium thiocyanate:

FeCl3+ 3 KCNS D Fe(CNS) 3 + 3 KCl.

Red

The resulting iron thiocyanate has a red color, the intensity of which depends on the concentration. By changing the color of the solution, one can judge the shift in chemical equilibrium depending on the increase or decrease in the content of iron thiocyanate in the reaction mixture. Create an equation for the equilibrium constant of this process.

Pour 20 ml of distilled water into a measuring cup or cylinder and add one drop of a saturated solution of iron (III) chloride and one drop of a saturated solution of potassium thiocyanate . Pour the resulting colored solution equally into four test tubes. Number the test tubes.

Add one drop of a saturated solution of iron (III) chloride to the first test tube. Add one drop of a saturated solution of potassium thiocyanate to the second test tube. Add crystalline potassium chloride to the third test tube and shake vigorously. The fourth test tube is for comparison.

Based on Le Chatelier's principle, explain what causes the color change in each individual case.

Write the results of the experiment in a table in the form

In the first and second case, we increased the concentration of the starting substances, so a more intense color is obtained. Moreover, in the second case the color is darker, because the concentration of KSCN changes at a cubic rate. In the third experiment, we increased the concentration of the final substance, so the color of the solution became lighter.

Conclusion: with an increase in the concentration of the starting substances, the equilibrium shifts towards the formation of reaction products. As the concentration of products increases, the equilibrium shifts towards the formation of starting substances.

General conclusions: during the experiments, we experimentally established the dependence of the reaction rate on the concentration of the starting substances (the higher the concentration, the higher the reaction rate); dependence of the reaction rate on temperature (the higher the temperature, the greater the reaction rate); how the concentration of reacting substances affects the chemical equilibrium (with an increase in the concentration of starting substances, the chemical equilibrium shifts towards the formation of products; with an increase in the concentration of products, the chemical equilibrium shifts towards the formation of starting substances)

The branch of chemistry that studies the rate of chemical reactions and its dependence on various factors is called chemical kinetics.

System in chemistry, the substance or collection of substances in question.

Phase- part of a system that is separated from other parts by an interface.

Systems consisting of one phase are called homogeneous, or homogeneous(gas mixtures, solutions).

Systems consisting of two or more phases are called heterogeneous or heterogeneous(gas + solid, liquid + solid).

Speed ​​of chemical reaction is the number of elementary acts occurring per unit time in a unit volume (homogeneous reactions) or on a unit surface (heterogeneous reactions).

Quantitatively, the reaction rate is usually characterized by a change in the concentration of any of the initial or final reaction products per unit time.

The reaction rate depends on the nature of the reacting substances and the reaction conditions: concentration, temperature, presence of catalysts, as well as on some other factors (pressure - for gas reactions, grinding - for solids, radioactive irradiation).

The quantitative dependence of the reaction rate on the concentration of reactants is expressed law of action of (acting) masses

(1867): the rate of a chemical reaction is directly proportional to the product of the concentrations of the reacting substances, taken in their stoichiometric coefficients. For example, for the reaction

aA + bB = cC + dD the reaction rate in accordance with the law of mass action is equal to

v = k[A] a -[B] b

Where [A] And [IN] - concentrations of starting substances; k- reaction rate constant, which depends on the nature of the reactants, temperature and the presence of catalysts, but does not depend on the concentrations of substances.

Addiction reaction rate versus temperature is expressed van't Hoff's rule:When the temperature increases by 10 °C, the rate of most chemical reactions increases by 2 - 4 times:

v t 2 = v t 1 γ (t 1 - t 2)/10

Where v t 1, v t 2 - reaction speed respectively at t 1 - initial temperature and t 2 - final temperature; γ- temperature coefficient of reaction rate.

The excess energy that molecules must have in order for their collision to lead to the formation of a new substance is called activation energy.

Addiction reaction rate constants from activation energy is expressed Arrhenius equation:

k =A e (- E act / RT)

Where A - constant, independent of temperature; E act- activation energy; R- universal gas constant; e - base of natural logarithms (e=2.7 18...); T- absolute temperature, K.

Catalysts- substances that increase the rate of a reaction, but are not consumed themselves. Inhibitors- substances that slow down the rate of reaction, but are not consumed.

The phenomenon of a change in reaction rate in the presence of catalysts is called catalysis.

Reactions that occur with the participation of catalysts are called catalytic reactions.

Homogeneous catalysts are in the same state of aggregation as the reactants. Heterogeneous catalysts are in a different state of aggregation than the reactants.

Chemical reactions occurring under given conditions in mutually opposite directions are called reversible.

The state of the system in which the rate of the forward reaction is equal to the rate of the reverse reaction is called chemical equilibrium.

Le Chatelier's principle.When external conditions change, the chemical equilibrium shifts towards the reaction (direct or reverse) that weakens this external influence.

Various factors influence the shift in equilibrium.

Temperature influence: As the temperature increases, the chemical equilibrium shifts toward the endothermic reaction.

Pressure influence: with increasing pressure, the chemical equilibrium shifts towards a decrease in the number of gas molecules.

Effect of concentration: When the concentration of any of the substances participating in the reaction increases, the chemical equilibrium shifts towards the consumption of this substance.

101. How will the reaction rate N 2 + 3H 2 = 2NH 3 change if the volume of gas
triple the mixture?

102. Rate of reaction A + B = C when the temperature increases by 10 degrees
increases by 3 times. How many times will the rate of this reaction increase when
temperature increase by 50°C?

103. When the temperature increases by 10 degrees, the rate of some reaction
increases 4 times. At what temperature should this reaction be carried out?
so that the speed of a reaction occurring at 100 degrees is reduced by 16 times?

104. Calculate how many times the rate of the reaction occurring in
gas phase, with an increase in temperature from 30 to 70°C, if the temperature
Is the reaction coefficient equal to 2?

105. The temperature coefficient of the reaction rate is 2.8. How many times
will the reaction rate increase when the temperature increases from 20 to 75°C?

106. Calculate the equilibrium constant for a reversible reaction

2NO 2 ↔ 2NO + O 2,

knowing that the equilibrium concentrations are: = 0.056 mol/l; [O 2 ] = 0.028 mol/l;

107. The equilibrium of the reaction H 2 + I 2 ↔ 2HI was established at the following concentrations of substances: [H 2 ] = 0.05 mol/l;
= 0.09 mol/l;

= 0.15 mol/l. Determine the initial concentrations of iodine and hydrogen.
108. Equilibrium constant of the homogeneous system CO(g) + H 2 O(g) ↔ CO? +H 2 (g)
at 850°C is equal to 1. Calculate the concentrations of all substances at equilibrium,

if the initial concentrations are: [CO] ref. = 3 mol/l; [H 2 O] RI = 2 mol/l.
109. How will the reaction rate 2NO(r) + O 2 (g) = 2NO 2 (r) change if we decrease

the volume of the reaction vessel is 3 times?

110. In the system A(g) + 2B(g) = C(g) the equilibrium concentrations are equal to: - [A] = 0.06 mol/l; [B] - 0.12 mol/l; [WITH]

111. 0.216 mol/l. Find the equilibrium constant of the reaction and the initial concentrations of substances A and B.
The activation energy of a certain reaction in the absence of a catalyst is equal to

74.24 kJ/mol, and with a catalyst - 50.14 kJ/mol. How many times will it increase?

rate of reaction in the presence of a catalyst if the reaction occurs at 25°C?
112. The reaction proceeds according to the equation 3A + B ↔ C. Concentration of substance A

decreased by 0.3 mol/l. What is the change in the concentration of substance B?
113. 8 mol SO 2 and 4 mol O 2 are mixed in a closed vessel. The reaction proceeds
at constant temperature. By the time equilibrium occurs, the reaction
80% of the original quantity enters. Determine the pressure of the gas mixture

at equilibrium, if the initial pressure is 300 kPa.

114. At a certain temperature, the dissociation constant of hydrogen iodide into simple substances is 6.25x10 -2. What percentage of HI dissociates at this temperature?
115. The initial concentration of each substance in the mixture is 2.5 mol/l. = 3 After equilibrium is established [C]

mol/l. Calculate the equilibrium constant of the system A + B ↔5 C + D.

116. Determine how the reaction rate of ammonia synthesis will change:

N 2 (r) + 3H 2 (g) = 2NH 3 (r) with: a) increasing the concentration of the starting substances by 3 times; b) when the pressure in the reaction mixture decreases by 2 times?
117. The reaction proceeds according to the equation N 2 + O 2 = 2NO. Source concentrations
substances before the start of the reaction were: = 0.049 mol/l; [O 2 ] = 0.01 mol/l.

Calculate the concentration of these substances at the moment when = 0.005 mol/l.
118. The reaction proceeds according to the equation N 2 + 3H 2 = 2NH 3. Concentrations of participants = the substances in it were: = 0.80 mol/l; [H 2 ] = 1.5 mol/l;
0.70 mol/l.

Calculate the concentration of hydrogen and ammonia when = 0.5 mol/l.
119. In a homogeneous system CO + C1 2 ↔ COCl equilibrium concentrations - reactants: [CO] - 0.2 mol/l; [С1 2 ]
[COCl] = 1.2 mol/l. Calculate the equilibrium constant of the system and the initial
concentrations of chlorine and CO.

120. Rate constant for the decomposition reaction of N 2 O, proceeding according to the equation
2N 2 O = 2N 2 + O 2, equal to 5x10 -4. Initial concentration = 6 mol/l.
Calculate the initial rate of the reaction and its rate when 50% of N 2 O decomposes.

121. The equilibrium of the homogeneous system 4HC1(g) + O 2 ↔ 2H 2 O(g) + 2C1 2 (g) was established at the following concentrations of the reactants:

[H 2 O] p = 0.14 mol/l; [C1 2 ] p = 0.14 mol/l; [HC1] p = 0.20 mol/l; [O 2 ] p = 0.32 mol/l. Calculate the initial concentrations of hydrogen chloride and oxygen.

122. Calculate the equilibrium constant for the homogeneous system CO(g) + H?O(g) ↔ CO 2 (g) + H 2 (g), if the equilibrium concentrations of the reactants are: [CO] p = 0.004 mol/l; [H 2 O] p = 0.064 mol/l; [CO 2 ] p = 0.016 mol/l; [H 2 ] p = 0.016 mol/l. What are the initial concentrations of water and CO?

123. The equilibrium constant of the homogeneous system N 2 + ZN 2 ↔2NH 3 at a certain temperature is 0.1. Equilibrium concentrations of hydrogen and

ammonia are respectively equal to 0.2 and 0.8 mol/l. Calculate the equilibrium and initial nitrogen concentrations.

124. Initial concentrations and [C1 2 ] in a homogeneous system
2NO + C1 2 ↔ 2NOC1 are 0.5 and 0.2 mol/l, respectively. Calculate
equilibrium constant if, by the time equilibrium occurs, reaction has occurred
20% NO.

125. In a homogeneous gas system A + B ↔ C + D, equilibrium has been established
at concentrations: [B] = 0.05 mol/l and [C] = 0.02 mol/l. Equilibrium constant
system is 0.04. Calculate the initial concentrations of substances A and B.

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2CHEMICAL KINETICS AND CHEMICAL EQUILIBRIUM

2.1 KINETICS OF CHEMICAL REACTIONS

Chemical reactions occur at different rates. Some of them are completed completely in small fractions of a second (explosion), others are carried out in minutes, hours, days and long periods of time. In addition, the same reaction can proceed quickly under some conditions (for example, at elevated temperatures), and slowly under others (for example, upon cooling). Moreover, the difference in the speed of the same reaction can be very large.

When considering the issue of reaction rates, it is necessary to distinguish between homogeneous and heterogeneous reactions. Closely related to these concepts is the concept of phase.

Phase is a part of a system separated from its other parts by an interface, during the transition through which the properties change abruptly.

A homogeneous reaction occurs in the volume of the phase [example - the interaction of hydrogen and oxygen with the formation of water vapor: H 2 (g) + O 2 (g) → H 2 O(g)], and if the reaction is heterogeneous, then it occurs at the phase interface [for example, carbon combustion: C(s) + O2(g) → CO 2 (g)].

The rate of a homogeneous reaction is the amount of substance that reacts or is formed during the reaction per unit time per unit volume of the phase:

Where n- amount of substance, mol; V- phase volume, l;τ - time; WITH- concentration, mol/l.

The rate of a heterogeneous reaction is the amount of substance that reacts or is formed during a reaction per unit time per unit surface area of ​​the phase:

Where S- surface area of ​​the phase interface.

The most important factors influencing the rate of a homogeneous reaction are the following: the nature of the reactants, their concentrations, temperature, and the presence of catalysts.

Dependence of the reaction rate on the concentrations of the reactants. A reaction between molecules occurs when they collide. Therefore, the rate of a reaction is proportional to the number of collisions that the molecules of the reacting substances undergo. The higher the concentration of each of the starting substances, the greater the number of collisions. For example, reaction rate A + B→ Proportional to the product of concentrations A and B:

v = k · [A] · [B],

Where k- proportionality coefficient, called reaction rate constant. Meaningful value k equal to the reaction rate for the case when the concentrations of the reactants are 1 mol/l.

This ratio expresses law of mass action This law is also called the law existing wt. : At constant temperature, the rate of a chemical reaction is directly proportional to the product of the concentrations of the reactants.

Much less often, a reaction occurs as a result of the simultaneous collision of three reacting particles. For example, reaction

2A+B → A 2 B

can occur through a triple collision:

A+ A + B→ A 2 B

Then, in accordance with the law of mass action, the concentration of each of the reacting substances is included in the expression of the reaction rate to a degree equal to the coefficient in the reaction equation:

v = k · [A] · [A] · [B] = k · [A] 2 [B]

The sum of exponents in the equation of the law of mass action is called reaction order. For example, in the latter case, the reaction is third order (second - with respect to substance A and first - with respect to substance B.

Dependence of reaction rate on temperature. If we use the results of counting the number of collisions between molecules, the number of collisions will be so large that all reactions must occur instantly. This contradiction can be explained by the fact that only molecules with some energy enter into the reaction.

The excess energy that molecules must have in order for their collision to lead to the formation of a new substance is called activation energy (See Figure 2.1).

Figure 2.1 - Energy diagram for the reaction of formation of product AB from starting substances A and B. If the collision energy of molecules A and B is greater than or equal to activation energies E a , then the energy barrier is overcome, and movement occurs along the reaction coordinate r from starting materials to product. Otherwise, an elastic collision of molecules A and B takes place. The top of the energy barrier corresponds to the transition state (activated complex), in which the AB bond is partially formed.

As temperature increases, the number of active molecules increases Temperature is a measure of the average kinetic energy of molecules, so increasing the temperature leads to an increase in the average speed of their movement.. Therefore, the rate of a chemical reaction should increase with increasing temperature. The increase in reaction rate upon heating is usually characterized as temperature coefficient of reaction rate (γ ) - a number showing how many times the rate of a given reaction increases when the temperature increases by 10 degrees. Mathematically, this dependence is expressed rule van't Hoff :

,

Where v 1 - speed at temperature t 1 ; v 2 - speed at temperature t 2. For most reactions the temperature coefficientγ lies in the range from 2 to 4.

More strictly, the dependence of the reaction rate (or rather, the rate constant) on temperature is expressed Arrhenius equation :

,

Where A - pre-exponential a multiplier that depends only on the nature of the reactants; E a - activation energy, which is the height of the energy barrier separating the starting materials and reaction products (see Figure 2.1); R R=8.3144 J/(mol. K). In approximate calculations, R = 8.31 J/(mol. K) is often taken. - universal gas constant; T T - absolute temperature (in Kelvin scale). It is related to temperature in Celsius by the equation
T = t o C + 273.15.
In approximate calculations, the relation is used
T = t o C + 273. -