• Domestic science and medicine in the 19th - early 20th centuries Chemistry of the 19th century in medicine
• What are peptides? Types and types of peptides. All about peptides and anti-aging cosmetics with peptides Additional peptides are not formed
• Toxic effect on humans of hazardous chemicals
• Radioautography Method of autoradiography in cytology
• Identification of bacteria by antigenic structure
• Organic matter in wastewater What are organic compounds in water
• Hydrogen index (pH factor)
• What is the pH of water and why is it important to know it?
• Redox potential Redox systems
• Reversible redox system Reversible redox systems

Lecture 1 Chemical thermodynamics. Chemical kinetics and catalysis PLAN 1. Basic concepts of thermodynamics. 2. Thermochemistry. 3. Chemical balance. 4. The rate of chemical reactions. 5. Effect of temperature on the rate of reactions. 6. The phenomenon of catalysis. Prepared by: Ph.D., Assoc. Ivanets L.M., ass. Kozachok S.S. Lecturer Assistant of the Department of Pharmaceutical Chemistry Kozachok Solomeya Stepanovna

Thermodynamics - Thermodynamics is a branch of physics that studies the mutual transformations of various types of energy associated with the transfer of energy in the form of heat and work. The great practical significance of thermodynamics is that it makes it possible to calculate the thermal effects of a reaction, to indicate in advance the possibility or impossibility of a reaction, as well as the conditions for its passage.

Internal energy Internal energy is the kinetic energy of all particles of the system (molecules, atoms, electrons) and the potential energy of their interactions, except for the kinetic and potential energy of the system as a whole. Internal energy is a state function, i.e. its change is determined by the given initial and final states of the system and does not depend on the path of the process: U = U 2 - U 1

The first law of thermodynamics Energy does not disappear without a trace and does not arise from nothing, but only passes from one form to another in an equivalent amount. A perpetual motion machine of the first kind, that is, a periodically operating machine that gives work without wasting energy, is impossible. Q \u003d U + W In any isolated system, the total energy supply remains unchanged. Q=U+W

The thermal effect of a chemical reaction at constant V or p does not depend on the reaction path, but is determined by the nature and state of the starting materials and reaction products Hess's law H 1 H 2 H 3 H 4 4 H 1 \u003d H 2 + H 3 + H 4

The second law of thermodynamics, like the first, is the result of centuries of human experience. There are various formulations of the second law, but they all determine the direction of spontaneous processes: 1. Heat cannot spontaneously transfer from a cold body to a hot one (Clausius' postulate). 2. A process whose only result is the conversion of heat into work is impossible (Thomson's postulate). 3. It is impossible to build a batch machine that only cools the heat reservoir and performs work (Planck's first postulate). 4. Any form of energy can be completely converted into heat, but heat is only partially converted into other types of energy (Planck's second postulate).

Entropy is a thermodynamic function of state, therefore its change does not depend on the path of the process, but is determined only by the initial and final states of the system. then S 2 - S 1 = ΔS = S 2 - S 1 = ΔS = The physical meaning of entropy is the amount of bound energy, which is related to one degree: in isolated systems, the direction of the flow of spontaneous processes is determined by the change in entropy.

Characteristic functions U is a function of an isochoric-isoentropic process: dU = TdS – pdV. For an arbitrary process: U 0 H is a function of an isobaric isoentropic process: dH = TdS + Vdp For an arbitrary process: H 0 S is a function of an isolated system For an arbitrary process: S 0 For an arbitrary process: S 0 F is a function of an isochoric isothermal process dF = dU – TdS. For an arbitrary process: F 0 G is a function of an isobaric-isothermal process: dG = dH- TdS For an arbitrary process: G 0

Classification of chemical reactions according to the number of stages Simple ones proceed in one elementary chemical act Complex ones proceed in several stages Reverse reaction A B

The effect of temperature on the rate of reactions The effect of temperature on the rate of enzymatic reactions t t

Comparison of van't Hoff: Calculation of the shelf life of medicines according to the "accelerated aging" method of van't Hoff: at t 2 t 1 Temperature coefficient of speed:

Page 1

FOUNDATIONS OF CHEMICAL THERMODYNAMICS AND CHEMICAL KINETICS

Parameter	Designation, unit	semantic meaning
Internal energy	U, kJ/mol	The total energy of the system, equal to the sum of the kinetic, potential and other types of energy of all particles of this system. This is a state function whose increment is equal to the heat received by the system in an isochoric process.
Work	A, kJ/mol	An energy measure of directed forms of particle motion in the process of system interaction with the environment.
Heat	Q, kJ/mol	Energy measure of chaotic forms of particle motion in the process of system interaction with the environment.
First law of thermodynamics	Q=∆U+A	The heat supplied to a closed system is used to increase the internal energy of the system and to perform work by the system against the external forces of the environment.
Entropy	S, J. (mol∙K) ∆S=Q/T, ∆S° r-tion =∑v 1 S°(prod.r-tion)-∑v 1 (out.in-in)	A state function that characterizes the degree of system disorder, i.e. inhomogeneity of the location and movement of its particles, the increment of which is equal to the heat supplied to the system in a reversible isothermal process, divided by the absolute temperature at which the process is carried out.
Enthalpy	H, kJ/mol ∆H=∆U+p∆V	State function characterizing the energy state of the system under isobaric conditions.
Enthalpy of reaction	∆H solution, kJ/mol	The amount of heat that is released or absorbed during chemical reactions under isobaric conditions.
standard condition	-	The most stable form at a given temperature (usually 298 K) and a pressure of 1 atm.
Standard Conditions	s.u.	Pressure: 101 325 Pa = 1 atm = 760 mm Hg Temperature: 25⁰С≈298K. n(X)=1 mol.
Standard enthalpy of formation of simple substances		At s.u. is taken equal to zero for simple substances in their most thermodynamically stable aggregate and allotropic states.
Standard enthalpy of formation of complex substances	∆H° arr298 (substance, state of aggregation), kJ/mol	The enthalpy of the reaction of formation of 1 mol of this substance from simple substances in s.u.
Standard enthalpy of combustion	∆H° burn (X), kJ/mol	The enthalpy of combustion (oxidation) of 1 mol of a substance to higher oxides in an oxygen environment at s.u.
Enthalpy of dissolution	∆H° r-tion, kJ/mol Where is the heat capacity of the solution	Thermal effect of the dissolution of a solid under isobaric conditions.
Gibbs energy	G, kJ/mol ∆G°=∆H-T∆S, ∆G° r-tion =∑v 1 ∆G° 1 (prod.r-tion)-∑ v 1 ∆G° 1 (out.in-c)	Free energy, a generalized thermodynamic function of the state of the system, taking into account the energy and disorder of the system under isobaric conditions.
Equilibrium constant of a chemical reaction for equilibrium	K equals, (mol/l) ∆ v , where ∆v depends on the values ​​of the stoichiometric coefficients of the substances. For the reaction aA+bB=cC+dD	It is equal to the ratio of the product of the equilibrium concentration of the reaction products to the product of the equilibrium concentrations of the reactants in powers equal to the stoichiometric coefficients.
van't Hoff isotherm equation	For a reversible reaction aA+bB=cC+dD , ∆G° p-tion \u003d-RTlnK is equal,	Allows you to calculate the Gibbs energy at given concentrations of reactants and reaction products.
Mass action law for kinetics	V=kc(A) a c(B) b	The reaction rate is proportional to the product of the concentrations of the reactants in powers, which are called the reaction orders for the corresponding substances.
Substance reaction order	n i	The exponent to which the concentration of a reactant enters into the equation for the rate of a chemical reaction. The order can be any value: integer, fractional, positive, zero, negative, and even a variable depending on the depth of the reaction.
General reaction order	n=nλ+nβ+…	Sum of reaction orders over all reactants.
Average reaction rate by substance		The average speed over the substance for a given period of time
True reaction rate		Characterizes the reaction rate at a given time (∆τ→0); v 1 is the stoichiometric coefficient of the substance in the reaction.
True reaction rate by substance		It characterizes the speed through the substance at a given time (∆τ→0).
Reaction rate constant	k, c -1 - for reactions of the 1st order; l / (mol∙s) - for reactions of the 2nd order	The individual characteristic of the reaction is numerically equal to the reaction rate at reagent concentrations equal to 1 mol/l.
Activation energy	Еа, kJ/mol	The minimum excess energy of interacting particles sufficient for these particles to enter into a chemical reaction.
Half life	Τ1/2, s, min, h, day	The time it takes for the concentration of a reactant to decrease by half.
Half life	Τ1/2, s, min, h, day	The time it takes for the amount of radioactive material to decrease by half.
Kinetic equation for 1-round reactions (integral form)	c=c 0 e - kt	The equation is linear in the variables ln c and t; k is the rate constant of the 1st order reaction; с 0 is the concentration of the initial substance at the initial moment of time; c is the current concentration of the initial substance at time t; t is the time elapsed from the beginning of the reaction.
Van't Hoff's rule	where is the temperature coefficient of the reaction rate;

Topic 3. General laws of chemical processes.

Chemical thermodynamics and kinetics

Introduction

Central to chemistry is the doctrine of the transformation of substances, including energy and the kinetics of chemical reactions. The assimilation of this doctrine makes it possible to predict the possibility and direction of chemical processes, calculate the energy effects and energy costs, the rate of production and yield of products in the reaction, influence the rate of chemical processes, and also prevent undesirable reactions in certain devices, installations and devices.

3.1. Chemical thermodynamics and kinetics

The exchange of energy between the system under study and the externalenvironment describe the laws that it studiesthermodynamics. The application of the laws of thermodynamics in chemistry allows us to solve the problem of the fundamental possibility of various processes, the conditions for their implementation,divide the degree of conversion of reactants into chimic reactions and evaluate their energetics.

Chemical thermodynamics , examines the relationship between work and energy in relation to chemical transformations.

Thermal and mechanical energy - algebraicquantities. Signs of quantitiesQ and BUT in thermodynamicsviewed in relation to the system. Energy, receivedreceived by the system is indicated by the sign “+”, given to the systemstem - sign "-".

Variables that determine the state of the SIstems are calledstate parameters. Among them in chemistry, the most commonly used pressure, temperature, volume, composition of the system. System status and prooutgoing changes in it are also characterized with the help ofstate functions, depending on the state parameters and not depending on the path of the system transition fromone state to another. These include internalenergy, enthalpy, entropy, isobaric-isothermal potential, etc.

Processes occurring at constant pressure -isobaric, at constant volumeisochoric, at constant temperature -isothermal. Majority chemical reactions take place in open vessels,i.e. at a constant pressure equal to atmospheric.

Chemical kineticsstudies the characteristics of a chemical process, such as the rate of a reaction and its dependence on external conditions.

3.2. Energy of chemical processes

Breakdown occurs during a chemical reactionsome chemical bonds and the formation of new ones. This process is accompanied by the release or absorption of heat.you, light or some other kind of energy. Energy effReaction effects are studied by the science of thermochemistry. In thermochemistryuse thermochemical reaction equations, whichwhich take into account:

aggregate state of matter;

thermal effect of the reaction (Q).

These equations often use fractional coefficients. So, the reaction equations for the formation of 1 mol of gasfigurative water is written as follows:

H 2 (g) + 1 / 2O 2 (g) \u003d H 2 O (g) + 242 kJ (*)

The symbol (d) indicates that hydrogen, oxygen andwater is in the gas phase. "+242 kJ" - means thatAs a result of this reaction, so much heat is released atformation of 1 mole of water.

The importance of taking into account the state of aggregation is related to the fact thatthat the heat of formation of liquid water is greater byheat released during the condensation of steam:

H 2 (g) + 1 / 2O 2 (g) \u003d H 2 O (g) + 286 kJ (**)

Condensation process:

H 2 O (g) \u003d H 2 O (g) + 44 kJ (***)

In addition to the thermal effect, thermodynamics usesyut the concept of "change in heat content" - enthalpy(reserve of internal energy) during the reaction ( H)

Exothermic reactions: heat is released Q > 0

internal energy is decreasing H<0

Endothermic reactions: heat is absorbed Q< 0

internal energy increases H>0.

Thus, the reaction (*) of water formation is exothermic.Thermal effect of the reaction:Q = 242 kJ,  H = -242 kJ.

Enthalpy of formation of chemical compounds

Standard enthalpy (heat) of formation chemical compound  H 0 f,V,298 is the change in enthalpy in the process of formation of one mole of this compound, which is in the standard state (p = 1 atm; T = 25 0 C), from simple substances, also in standard states and phases and modifications that are thermodynamically stable at a given temperature .

The standard enthalpies of formation of simple substances are taken equal to zero if their states of aggregation and modifications are stable under standard conditions.

The standard enthalpies of formation of substances are collected and summarized in reference books.

3.2. 1. Thermochemical calculations

The independence of the heat of a chemical reaction from the process path at p=const was established in the first half of the 19th century. Russian scientist G.I. Hess: the thermal effect of a chemical reaction does not depend on the path of its occurrence, but depends only on the nature and physical state of the initial substances and reaction products.

For most reactions, the change in the thermal effect within the temperature limits of practical importance is small. Therefore, in the future we will use  H 0 f,B,298 and are assumed to be constant in calculations.

Consequence from Hess' law– the heat effect of a chemical reaction is equal to the sum of the heats (enthalpies) of the formation of the reaction products minus the sum of the heats (enthalpies) of the formation of the starting substances.

When using the corollary from the Hess law in thermochemical calculations, it should be borne in mind that stoichiometric coefficients in the reaction equation should be taken into account in algebraic summation.

So, for the reaction equation aA + bB = cC + dD, the thermal effect  H is equal to

H=(s  N ex.C +d N ex.D) – (а N ex.A +v N ex.B) (*)

Equation (*) makes it possible to determine both the thermal effect of the reaction from the known enthalpies of formation of the substances participating in the reaction, and one of the enthalpies of formation, if the thermal effect of the reaction and all other enthalpies of formation are known.

Fuel combustion heat

The thermal effect of the oxidation reaction with oxygen of the elements that make up the substance to the formation of higher oxides is called the heat of combustion of this substance
.

Example: determine the heat of combustion of ethanol C 2 H 5 OH (g)

If a calculation conducted for
with the formation of liquid water, then the heat of combustion is called higher, if with the formation of gaseous water, then lower. By default, they usually mean the higher calorific value.

In technical calculations, the specific heat of combustion Q T is used, which is equal to the amount of heat released during the combustion of 1 kg of a liquid or solid substance or 1 m 3 of a gaseous substance, then

Q T = -  N ST  1000/M (for w, tv.)

Q T = -  N ST  1000/22.4 (for city),

where M is the mass of a mole of a substance, 22.4 l is the volume of a mole of gas.

3.3. Chemical and phase equilibrium

3.3.1. Chemical equilibrium

Reversible reactions - chemical reactions occurring simultaneously in two opposite directions.

Chemical equilibrium - the state of the system in which the rate of the direct reaction (V 1 ) is equal to the rate of the reverse reaction (V 2 ). In chemical equilibrium, the concentrations of substances remain unchanged. Chemical equilibrium has a dynamic character: forward and reverse reactions do not stop at equilibrium.

The state of chemical equilibrium is quantitatively characterized by the equilibrium constant, which is the ratio of the constants of the straight line (K 1 ) and reverse (K 2 ) reactions.

For the reaction mA + nB « pC + dD the equilibrium constant is

K=K 1 /k 2 = ([C] p[D] d) / ([A] m[B] n)

The equilibrium constant depends on the temperature and the nature of the reactants. The larger the equilibrium constant, the more the equilibrium is shifted towards the formation of direct reaction products.

Ways to shift the balance

Le Chatelier's principle. If an external influence is made on a system in equilibrium (concentration, temperature, pressure change), then it favors the flow of one of the two opposite reactions that weakens this effect.

V 1

A+B

V 2

Pressure. An increase in pressure (for gases) shifts the equilibrium towards a reaction leading to a decrease in volume (i.e., to the formation of a smaller number of molecules).

V 1

A+B

; an increase in P leads toV 1 >V 2

V 2

An increase in temperature shifts the equilibrium position towards an endothermic reaction (i.e. towards a reaction proceeding with the absorption of heat)

V 1

A+B

B + Q, then increase t° C leads to V 2 > V 1

V 2

V 1

A+B

B - Q, then increase t° C leads to V 1 > V 2

V 2

Increasing the concentration of starting materials and removing products from the reaction sphere shifts the equilibrium towards the direct reaction. Increasing concentrations of starting materials [A] or [B] or [A] and [B]: V 1 > V 2 .

Catalysts do not affect the equilibrium position.

3.3.2. Phase equilibria

The equilibrium of the process of transition of a substance from one phase to another without changing the chemical composition is called phase equilibrium.

Examples of phase equilibrium:

Solid.............Liquid

Liquid .......... Steam

3.3.3. Reaction rate and methods of its regulation

Speed ​​reaction is determined by the change in the molar concentration of one of the reactants:

V = ± (С 2 - С 1) / (t 2 - t 1) \u003d ± D FROM / D t

where C 1 and C 2 - molar concentrations of substances at time t 1 and t2 respectively (sign (+) - if the rate is determined by the reaction product, sign (-) - by the starting material).

Reactions occur when molecules of reactants collide. Its speed is determined by the number of collisions and the likelihood that they will lead to a transformation. The number of collisions is determined by the concentrations of the reacting substances, and the probability of a reaction is determined by the energy of the colliding molecules.

Factors affecting the rate of chemical reactions

The nature of the reactants. An important role is played by the nature of chemical bonds and the structure of the molecules of the reagents. Reactions proceed in the direction of the destruction of less strong bonds and the formation of substances with stronger bonds. For example, to break bonds in H 2 and N 2 high energies are required; such molecules are not very reactive. To break bonds in highly polar molecules (HCl, H 2 O) less energy is required and the reaction rate is much faster. Reactions between ions in electrolyte solutions proceed almost instantaneously.

Examples: Fluorine reacts explosively with hydrogen at room temperature; bromine reacts with hydrogen slowly even when heated.

Calcium oxide reacts vigorously with water, releasing heat; copper oxide - does not react.

Concentration. With an increase in concentration (the number of particles per unit volume), collisions of reactant molecules occur more often - the reaction rate increases.

The law of active masses (K. Guldberg, P. Waage, 1867)

The rate of a chemical reaction is directly proportional to the product of the concentrations of the reactants.

aA + bB + . . .® . . .

V=k[A] a[B] b . . .

The reaction rate constant k depends on the nature of the reactants, temperature, and catalyst, but does not depend on the concentrations of the reactants.

The physical meaning of the rate constant is that it is equal to the reaction rate at unit concentrations of the reactants.

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

Temperature. With an increase in temperature for every 10° C, the reaction rate increases by 2-4 times (Van't Hoff's Rule). As the temperature increases from t 1 to t 2 the change in reaction rate can be calculated by the formula:

(t 2 - t 1 ) / 10

Vt 2 / Vt 1

= g

(where Vt 2 and Vt 1 - reaction rates at temperatures t 2 and t1 respectively;gis the temperature coefficient of this reaction).

Van't Hoff's rule is applicable only in a narrow temperature range. More accurate is the Arrhenius equation:

k \u003d Ae -Ea / RT

where

A is a constant depending on the nature of the reactants;

R is the universal gas constant;

Ea is the activation energy, i.e. the energy that colliding molecules must have in order for the collision to result in a chemical transformation.

Energy diagram of a chemical reaction.

exothermic reaction

Endothermic reaction

A - reagents, B - activated complex (transition state), C - products.

The higher the activation energy Ea, the more the reaction rate increases with increasing temperature.

1. The contact surface of the reactants. For heterogeneous systems (when substances are in different states of aggregation), the larger the contact surface, the faster the reaction proceeds. The surface of solids can be increased by grinding them, and for soluble substances by dissolving them.

3.3.4. Mechanisms of chemical reactions, oscillatory reactions

Classification of chemical reactions

I . According to the number and composition of the starting materials and reaction products:

1) Reactions connections are reactions in which two or more substances form one substance of a more complex composition. Reactions of the combination of simple substances are always redox reactions. Complex substances can also participate in compound reactions.

2) Reactions decomposition Reactions in the course of which two or more simpler substances are formed from one complex substance.
Decomposition products of the initial substance can be both simple and complex substances.

Decomposition reactions usually proceed when substances are heated and are endothermic reactions. Like compound reactions, decomposition reactions can proceed with or without changing the oxidation states of the elements;

3) Reactions substitution - these are reactions between simple and complex substances, during which the atoms of a simple substance replace the atoms of one of the elements in the molecule of a complex substance, as a result of the substitution reaction, a new simple and a new complex substance are formed.
These reactions are almost always redox reactions.

4) Reactions exchange - these are reactions between two complex substances, the molecules of which exchange their constituent parts.
Exchange reactions always proceed without electron transfer, that is, they are not redox reactions.

II . On the basis of changes in the degree of oxidation

1) Reactions that go without changing the oxidation state - neutralization reactions

2) With a change in the degree of oxidation

III . Depending on the presence of a catalyst

1) Non-catalytic (go without the presence of a catalyst);

2) catalytic (comes with a catalyst)

IV . According to the thermal effect

1) exothermic (with heat release):

2) Endothermic (with heat absorption):

V . On the basis of reversibility

1) irreversible (flow in one direction only):

2) reversible (flowing simultaneously in the forward and reverse directions):

VI . On the basis of homogeneity

1) homogeneous (flowing in a homogeneous system):

2) Heterogeneous (flowing in an inhomogeneous system):

According to the flow mechanism All reactions can be divided into simple and complex. Simple reactions proceed in one stage and are called one-stage.

Complex reactions proceed either sequentially (multi-stage reactions), or in parallel, or in series-parallel.

Each reaction step can involve one molecule (monomolecular reactions), two molecules (bimolecular reactions), and three molecules (trimolecular reactions).

Vibrational reactions - a class of chemical reactions occurring in an oscillatory mode, in which some reaction parameters (color, concentration of components, temperature, etc.) change periodically, forming a complex spatio-temporal structure of the reaction medium.

(System bromate-malonic acid-cerium Belousov-Zhabotinsky reaction)

3.4. Catalysis

Substances that participate in reactions and increase its rate, remaining unchanged at the end of the reaction, are calledcatalysts .

The mechanism of action of catalysts is associated with a decrease in the activation energy of the reaction due to the formation of intermediate compounds.

At homogeneous catalysis the reactants and the catalyst constitute one phase (they are in the same state of aggregation).

At heterogeneous catalysis - different phases (they are in different states of aggregation).

In some cases, the course of undesirable chemical processes can be drastically slowed down by adding to the reaction mediuminhibitors(phenomenon " negative catalysis ").

Solving problems in the section

The topic "Chemical thermodynamics and kinetics", which involves the study of conditions affecting the rate of a chemical reaction, occurs twice in the school chemistry course - in the 9th and 11th grades. However, it is this topic that is one of the most difficult and difficult enough not only for understanding by the “average” student, but even for presentation by some teachers, especially non-specialists working in rural areas, for whom chemistry is an additional subject, taking into account the hours of which the teacher gains rate, and hence the hope for a more or less decent salary.
In conditions of a sharp decrease in the number of students in rural schools, for well-known reasons, the teacher is forced to be a generalist. Having attended 2-3 courses, he begins teaching subjects, often very far from his main specialty.
This development is focused primarily on novice teachers and subject teachers who are forced to teach chemistry in a market economy. The material contains tasks to find the rates of heterogeneous and homogeneous reactions and increase the reaction rate with increasing temperature. Despite the fact that these tasks are based on school material, albeit difficult for the “average” student to master, it is advisable to solve several of them at a chemistry lesson in
11th grade, and the rest to offer in a circle or optional class for students who plan to connect their future fate with chemistry.
In addition to the tasks analyzed in detail and provided with answers, this development contains theoretical material that will help a chemistry teacher, primarily a non-specialist, to understand the essence of this complex topic of a general chemistry course.
Based on the proposed material, you can create your own version of the lesson-lecture, depending on the abilities of the students in the class, and you can use the proposed theoretical part when studying this topic both in the 9th and 11th grades.
Finally, the material contained in this development will not be superfluous to analyze on their own for a graduate who is preparing to enter a university, including one in which chemistry is a major subject.

Theoretical part on the topic
"Chemical thermodynamics and kinetics"

Conditions affecting the rate of a chemical reaction

1. The rate of a chemical reaction depends on the nature of the reactants.

EXAMPLES.

Metallic sodium, having an alkaline nature, reacts violently with water with the release of a large amount of heat, in contrast to zinc, which has an amphoteric nature, which reacts with water slowly and when heated:

Powdered iron reacts more vigorously with strong mineral hydrochloric acid than with weak organic acetic acid:

2. The rate of a chemical reaction depends on the concentration of reacting substances in a dissolved or gaseous state.

EXAMPLES.

Sulfur burns more vigorously in pure oxygen than in air:

With a 30% solution of hydrochloric acid, powdered magnesium reacts more vigorously than with a 1% solution of it:

3. The rate of a chemical reaction is directly proportional to the surface area of ​​the reacting substances in the solid state of aggregation.

EXAMPLES.

A piece of charcoal (carbon) is very difficult to light with a match, but charcoal dust burns with an explosion:

C + O 2 \u003d CO 2.

Aluminum in the form of a granule does not quantitatively react with an iodine crystal, but crushed iodine combines vigorously with aluminum in the form of powder:

4. The rate of a chemical reaction depends on the temperature at which the process occurs.

EXAMPLE

For every 10°C increase in temperature, the rate of most chemical reactions increases by a factor of 2–4. A specific increase in the rate of a chemical reaction is determined by a specific temperature coefficient (gamma).

Calculate how many times the reaction rate will increase:

2NO + O 2 \u003d 2NO 2,

if the temperature coefficient is 3 and the process temperature has increased from 10 °C to 50 °C.

The temperature change is:

t= 50 °С - 10 °С = 40 °С.

We use the formula:

where is the rate of a chemical reaction at an elevated temperature, is the rate of a chemical reaction at an initial temperature.

Consequently, the rate of a chemical reaction with an increase in temperature from 10 °C to 50 °C will increase 81 times.

5. The rate of a chemical reaction depends on the presence of certain substances.

Catalyst A substance that speeds up the course of a chemical reaction, but is not itself consumed during the reaction. The catalyst lowers the activation barrier of a chemical reaction.

Inhibitor A substance that slows down the course of a chemical reaction, but is not itself consumed during the reaction.

EXAMPLES.

The catalyst that accelerates the course of this chemical reaction is manganese (IV) oxide.

Red phosphorus is the catalyst that speeds up the course of this chemical reaction.

An inhibitor that slows down the course of this chemical reaction is an organic substance - urotropine (hexamethylenetetramine).

The rate of a homogeneous chemical reaction is measured by the number of moles of the substance that has entered into the reaction or formed as a result of the reaction per unit of time per unit volume:

where homog is the rate of a chemical reaction in a homogeneous system, is the number of moles of one of the reactants or one of the substances formed as a result of the reaction, V- volume,
t- time, - change in the number of moles of a substance during the reaction t.

Since the ratio of the number of moles of a substance to the volume of the system is the concentration With, then

Consequently:

The rate of a homogeneous chemical reaction is measured in mol/(l s).

With this in mind, the following definition can be given:

the rate of a homogeneous chemical reaction is equal to the change in the concentration of one of the reactants or one of the substances formed as a result of the reaction per unit time.

If the reaction proceeds between substances in a heterogeneous system, then the reacting substances do not come into contact with each other in the entire volume, but only on the surface of the solid. So, for example, when a piece of crystalline sulfur burns, oxygen molecules react only with those sulfur atoms that are on the surface of the piece. When grinding a piece of sulfur, the area of ​​the reacting surface increases, and the burning rate of sulfur increases.

In this regard, the definition of the rate of a heterogeneous chemical reaction is as follows:

the rate of a heterogeneous chemical reaction is measured by the number of moles of the substance that has entered into the reaction or formed as a result of the reaction per unit time per unit surface:

where S is the surface area.

The rate of a heterogeneous chemical reaction is measured in mol / (cm 2 s).

Related tasks
"Chemical thermodynamics and kinetics"

1. 4 mol of nitric oxide (II) and an excess of oxygen were introduced into the vessel for carrying out chemical reactions. After 10 s, the amount of nitric oxide (II) substance turned out to be 1.5 mol. Find the rate of this chemical reaction if it is known that the volume of the vessel is 50 liters.

2. The amount of methane substance in the vessel for chemical reactions is 7 mol. An excess of oxygen was introduced into the vessel and the mixture was exploded. It was experimentally established that after 5 s the amount of methane substance decreased by 2 times. Find the rate of this chemical reaction if it is known that the volume of the vessel is 20 liters.

3. The initial concentration of hydrogen sulfide in the gas combustion vessel was 3.5 mol/l. An excess of oxygen was introduced into the vessel and the mixture was exploded. After 15 seconds, the concentration of hydrogen sulfide was 1.5 mol/l. Find the rate of this chemical reaction.

4. The initial concentration of ethane in the gas combustion vessel was 5 mol/L. An excess of oxygen was introduced into the vessel and the mixture was exploded. After 12 seconds, the ethane concentration was 1.4 mol/L. Find the rate of this chemical reaction.

5. The initial ammonia concentration in the gas combustion vessel was 4 mol/l. An excess of oxygen was introduced into the vessel and the mixture was exploded. After 3 seconds, the ammonia concentration was 1 mol/L. Find the rate of this chemical reaction.

6. The initial concentration of carbon monoxide(II) in the gas combustion vessel was 6 mol/l. An excess of oxygen was introduced into the vessel and the mixture was exploded. After 5 s, the concentration of carbon monoxide(II) decreased by half. Find the rate of this chemical reaction.

7. A piece of sulfur with a reacting surface area of ​​7 cm 2 was burned in oxygen with the formation of sulfur oxide (IV). For 10 s, the amount of sulfur matter decreased from 3 mol to 1 mol. Find the rate of this chemical reaction.

8. A piece of carbon with a reacting surface area of ​​10 cm 2 was burned in oxygen to form carbon monoxide(IV). In 15 seconds, the amount of carbon matter decreased from 5 mol to 1.5 mol. Find the rate of this chemical reaction.

9. A magnesium cube with a total reacting surface area of ​​15 cm2 and an amount of substance
6 mol burned in excess oxygen. At the same time, 7 s after the start of the reaction, the amount of magnesium substance turned out to be equal to 2 mol. Find the rate of this chemical reaction.

10. A bar of calcium with a total reacting surface area of ​​12 cm 2 and an amount of substance of 7 mol was burned in an excess of oxygen. At the same time, 10 s after the start of the reaction, the amount of calcium substance turned out to be 2 times less. Find the rate of this chemical reaction.

Solutions and answers

1 (NO) = 4 mol,

O 2 - excess,

t 2 = 10 s,

t 1 = 0 s,

2 (NO) = 1.5 mol,

Find:

Solution

2NO + O 2 \u003d 2NO 2.

Using the formula:

R-tion \u003d (4 - 1.5) / (50 (10 - 0)) \u003d 0.005 mol / (l s).

Answer. p-tion \u003d 0.005 mol / (l s).

2.

1 (CH 4) \u003d 7 mol,

O 2 - excess,

t 2 = 5 s

t 1 = 0 s,

2 (CH 4) \u003d 3.5 mol,

Find:

Solution

CH 4 + 2O 2 \u003d CO 2 + 2H 2 O.

Using the formula:

find the rate of this chemical reaction:

R-tion \u003d (7 - 3.5) / (20 (5 - 0)) \u003d 0.035 mol / (l s).

Answer. p-tion \u003d 0.035 mol / (l s).

3.

s 1 (H 2 S) = 3.5 mol / l,

O 2 - excess,

t 2 = 15 s,

t 1 = 0 s,

With 2 (H 2 S) \u003d 1.5 mol / l.

Find:

Solution

2H 2 S + 3O 2 \u003d 2SO 2 + 2H 2 O.

Using the formula:

find the rate of this chemical reaction:

R-tions \u003d (3.5 - 1.5) / (15 - 0) \u003d 0.133 mol / (l s).

Answer. p-tion \u003d 0.133 mol / (l s).

4.

s 1 (C 2 H 6) = 5 mol / l,

O 2 - excess,

t 2 = 12 s,

t 1 = 0 s,

c 2 (C 2 H 6) \u003d 1.4 mol / l.

Find:

Solution

2C 2 H 6 + 7O 2 \u003d 4CO 2 + 6H 2 O.

find the rate of this chemical reaction:

P-tions \u003d (6 - 2) / (15 (7 - 0)) \u003d 0.0381 mol / (cm 2 s).

Answer. p-tion \u003d 0.0381 mol / (cm 2 s).

10. Answer. p-tion \u003d 0.0292 mol / (cm 2 s).

Literature

Glinka N.L. General Chemistry, 27th ed. Ed. V.A.Rabinovich. L.: Chemistry, 1988; Akhmetov N.S. General and inorganic chemistry. M.: Higher. school, 1981; Zaitsev O.S. General chemistry. M.: Higher. school, 1983; Karapetyants M.Kh., Drakin S.I. General and inorganic chemistry. M.: Higher. school, 1981; Korolkov D.V. Fundamentals of inorganic chemistry. Moscow: Education, 1982; Nekrasov B.V. Fundamentals of General Chemistry. 3rd ed., M.: Chemistry, 1973; Novikov G.I. Introduction to inorganic chemistry. Ch. 1, 2. Minsk: Highest. school, 1973–1974; Schukarev S.A.. Inorganic chemistry. T. 1, 2. M.: Higher. school, 1970–1974; Schroeter W., Lautenschläger K.-H., Bibrak H. et al. Chemistry. Reference ed. Per. with him. Moscow: Chemistry, 1989; Feldman F.G., Rudzitis G.E. Chemistry-9. Textbook for the 9th grade of high school. M.: Education, 1990; Feldman F.G., Rudzitis G.E. Chemistry-9. Textbook for the 9th grade of high school. M.: Enlightenment, 1992.

The rate of chemical reactions. Concept definition. Factors affecting the rate of a chemical reaction: reagent concentration, pressure, temperature, presence of a catalyst. The law of mass action (LMA) as the basic law of chemical kinetics. The rate constant, its physical meaning. Influence on the reaction rate constant of the nature of the reactants, temperature and the presence of a catalyst.

1. With. 102-105; 2. With. 163-166; 3. With. 196-207, p. 210-213; 4. With. 185-188; 5. With. 48-50; 6. With. 198-201; 8. With. 14-19

Homogeneous reaction rate - this is a value numerically equal to the change in the concentration of any participant in the reaction per unit time.

Average reaction rate v cf in the time interval from t 1 to t 2 is determined by the ratio:

The main factors affecting the rate of a homogeneous chemical reaction :

- the nature of the reactants;

- reagent concentration;

- pressure (if gases are involved in the reaction);

- temperature;

- the presence of a catalyst.

Heterogeneous reaction rate - this is a value numerically equal to the change in the concentration of any participant in the reaction per unit time per unit surface: .

According to the stages of chemical reactions are divided into elementary and complex. Most chemical reactions are complex processes that occur in several stages, i.e. consisting of several elementary processes.

For elementary reactions, law of mass action: the rate of an elementary chemical reaction at a given temperature is directly proportional to the product of the concentrations of reactants in powers equal to the stoichiometric coefficients of the reaction equation.

For an elementary reaction aA + bB → ... the reaction rate, according to the law of mass action, is expressed by the ratio:

wheres (A) and With (AT) - molar concentrations of reactants BUT and AT; a and b- corresponding stoichiometric coefficients; k- rate constant of this reaction .

For heterogeneous reactions, the equation of the law of mass action does not include the concentrations of all reagents, but only gaseous or dissolved ones. So, for the combustion reaction of carbon:

C (c) + O 2 (g) → CO 2 (g)

the velocity equation has the form .

The physical meaning of the rate constant is it is numerically equal to the rate of a chemical reaction at concentrations of reactants equal to 1 mol/dm 3 .

The value of the rate constant of a homogeneous reaction depends on the nature of the reactants, temperature and catalyst.

Effect of temperature on the rate of a chemical reaction. Temperature coefficient of the rate of a chemical reaction. active molecules. Distribution curve of molecules according to their kinetic energy. Activation energy. Ratio of activation energy and chemical bond energy in initial molecules. Transition state, or activated complex. Activation energy and thermal effect of the reaction (energy scheme). Dependence of the temperature coefficient of the reaction rate on the value of the activation energy.

1. With. 106-108; 2. With. 166-170; 3. With. 210-217; 4. With. 188-191; 5. With. 50-51; 6. With. 202-207; 8 . With. 19-21.

As the temperature increases, the rate of a chemical reaction usually increases.

The value showing how many times the reaction rate increases with an increase in temperature by 10 degrees (or, what is the same, by 10 K), is called temperature coefficient of chemical reaction rate (γ):

where are the reaction rates, respectively, at temperatures T 2 and T 1 ; γ is the temperature coefficient of the reaction rate.

The dependence of the reaction rate on temperature is approximately determined by the empirical van't Hoff's rule: for every 10 degrees increase in temperature, the rate of a chemical reaction increases by 2-4 times.

A more accurate description of the dependence of the reaction rate on temperature is feasible within the framework of the Arrhenius activation theory. According to this theory, a chemical reaction can only occur when active particles collide. active Particles are called that have a certain energy characteristic of a given reaction, which is necessary to overcome the repulsive forces that arise between the electron shells of the reacting particles.

The proportion of active particles increases with increasing temperature.

Activated complex - this is an intermediate unstable grouping, which is formed during the collision of active particles and is in a state of redistribution of bonds. The reaction products are formed during the decomposition of the activated complex.

Activation energy and E a is equal to the difference between the average energy of the reacting particles and the energy of the activated complex.

For most chemical reactions, the activation energy is less than the dissociation energy of the weakest bond in the molecules of the reactants.

In activation theory, the influence temperature on the rate of a chemical reaction is described by the Arrhenius equation for the rate constant of a chemical reaction:

where BUT is a constant factor that does not depend on temperature and is determined by the nature of the reactants; e is the base of the natural logarithm; E a is the activation energy; R is the molar gas constant.

As follows from the Arrhenius equation, the higher the rate constant of the reaction, the lower the activation energy. Even a slight decrease in the activation energy (for example, when a catalyst is introduced) leads to a noticeable increase in the reaction rate.

According to the Arrhenius equation, an increase in temperature leads to an increase in the rate constant of a chemical reaction. The larger the value E a, the more noticeable the effect of temperature on the reaction rate and, therefore, the greater the temperature coefficient of the reaction rate.

Effect of a catalyst on the rate of a chemical reaction. Homogeneous and heterogeneous catalysis. Elements of the theory of homogeneous catalysis. Theory of intermediate compounds. Elements of the theory of heterogeneous catalysis. Active centers and their role in heterogeneous catalysis. The concept of adsorption. Influence of a catalyst on the activation energy of a chemical reaction. Catalysis in nature, industry, technology. biochemical catalysis. Enzymes.

1. With. 108-109; 2. With. 170-173; 3. With. 218-223; 4 . With. 197-199; 6. With. 213-222; 7. With. 197-202.; 8. With. 21-22.

catalysis called the change in the rate of a chemical reaction under the influence of substances, the number and nature of which after the completion of the reaction remain the same as before the reaction.

Catalyst - This is a substance that changes the rate of a chemical reaction and remains chemically unchanged after it.

positive catalyst speeds up the reaction negative catalyst, or inhibitor slows down the reaction.

In most cases, the effect of a catalyst is explained by the fact that it reduces the activation energy of the reaction. Each of the intermediate processes involving a catalyst proceeds with a lower activation energy than the non-catalyzed reaction.

At homogeneous catalysis the catalyst and reactants form one phase (solution). At heterogeneous catalysis the catalyst (usually a solid) and the reactants are in different phases.

In the course of homogeneous catalysis, the catalyst forms an intermediate compound with the reagent, which reacts with the second reagent at a high rate or rapidly decomposes with the release of the reaction product.

An example of homogeneous catalysis: the oxidation of sulfur oxide (IV) to sulfur oxide (VI) with oxygen in the nitrous method for producing sulfuric acid (here the catalyst is nitrogen oxide (II), which easily reacts with oxygen).

In heterogeneous catalysis, the reaction proceeds on the surface of the catalyst. The initial stages are the diffusion of reactant particles to the catalyst and their adsorption(i.e. absorption) by the catalyst surface. Reagent molecules interact with atoms or groups of atoms located on the surfaces of the catalyst, forming intermediate surface connections. The redistribution of electron density that occurs in such intermediate compounds leads to the formation of new substances, which desorbed, i.e., are removed from the surface.

The process of formation of intermediate surface compounds occurs on active centers catalyst - on surface areas characterized by a special distribution of electron density.

An example of heterogeneous catalysis: the oxidation of sulfur oxide (IV) to sulfur oxide (VI) with oxygen in the contact method for producing sulfuric acid (vanadium oxide (V) with additives can be a catalyst here).

Examples of catalytic processes in industry and technology: the synthesis of ammonia, the synthesis of nitric and sulfuric acids, the cracking and reforming of oil, the afterburning of products of incomplete combustion of gasoline in cars, etc.

Examples of catalytic processes in nature are numerous, since most biochemical reactions- chemical reactions occurring in living organisms - are among the catalytic reactions. These reactions are catalyzed by proteins called enzymes. There are about 30 thousand enzymes in the human body, each of which catalyses the passage of only one process or one type of processes (for example, ptyalin in saliva catalyzes the conversion of starch into sugar).

chemical balance. Reversible and irreversible chemical reactions. state of chemical equilibrium. Chemical equilibrium constant. Factors that determine the value of the equilibrium constant: the nature of the reactants and temperature. Shift in chemical equilibrium. Influence of changes in concentration, pressure and temperature on the position of chemical equilibrium.

1. With. 109-115; 2. With. 176-182; 3 . With. 184-195, p. 207-209; 4. pp.172-176, p. 187-188; 5. With. 51-54; 8 . With. 24-31.

Chemical reactions, as a result of which the initial substances are completely converted into reaction products, are called irreversible. Reactions that occur simultaneously in two opposite directions (forward and reverse) are calledreversible.

In reversible reactions, the state of the system in which the rates of the forward and reverse reactions are equal () is called state of chemical equilibrium. The chemical equilibrium is dynamic, i.e., its establishment does not mean the termination of the reaction. In the general case, for any reversible reaction аА + bB ↔ dD + eE, regardless of its mechanism, the relation is fulfilled:

At steady equilibrium, the product of the concentrations of the reaction products, referred to the product of the concentrations of the starting materials, for a given reaction at a given temperature is a constant value called equilibrium constant(To).

The value of the equilibrium constant depends on the nature of the reactants and temperature, but does not depend on the concentrations of the components of the equilibrium mixture.

Changing the conditions (temperature, pressure, concentration), under which the system is in a state of chemical equilibrium (), causes an imbalance. As a result of unequal changes in the rates of direct and reverse reactions () over time, a new chemical equilibrium () is established in the system, corresponding to new conditions. The transition from one equilibrium state to another is called a shift, or displacement, of the equilibrium position.

If, during the transition from one equilibrium state to another, the concentrations of substances recorded on the right side of the reaction equation increase, they say that balance shifts to the right. If, during the transition from one equilibrium state to another, the concentrations of substances recorded on the left side of the reaction equation increase, they say that balance shifts to the left.

The direction of shift of chemical equilibrium as a result of changes in external conditions is determined by Le Chatelier's principle: If an external influence is exerted on a system that is in a state of chemical equilibrium, then it will favor the flow of one of the two opposite processes that weakens this influence.

According to Le Chatelier's principle,

An increase in the concentration of the component written on the left side of the equation leads to a shift in equilibrium to the right; an increase in the concentration of the component written on the right side of the equation leads to a shift in equilibrium to the left;

With an increase in temperature, the equilibrium shifts towards the occurrence of an endothermic reaction, and with a decrease in temperature, in the direction of an exothermic reaction;

With an increase in pressure, the equilibrium shifts towards a reaction that reduces the number of molecules of gaseous substances in the system, and with a decrease in pressure, towards a reaction that increases the number of molecules of gaseous substances.

Photochemical and chain reactions. Features of the course of photochemical reactions. Photochemical reactions and wildlife. Unbranched and branched chemical reactions (on the example of the reactions of the formation of hydrogen chloride and water from simple substances). Conditions for the initiation and termination of chains.

2. With. 173-176; 3. With. 224-226; 4. 193-196; 6. With. 207-210; 8. With. 49-50.

Photochemical reactions - These are reactions that take place under the influence of light. A photochemical reaction proceeds if the reagent absorbs radiation quanta, which are characterized by an energy that is quite specific for this reaction.

In the case of some photochemical reactions, by absorbing energy, the reactant molecules pass into an excited state, i.e. become active.

In other cases, a photochemical reaction proceeds if quanta of such high energy are absorbed that chemical bonds are broken and the molecules dissociate into atoms or groups of atoms.

The rate of the photochemical reaction is the greater, the greater the intensity of irradiation.

An example of a photochemical reaction in wildlife: photosynthesis, i.e. the formation by organisms of organic substances of cells due to the energy of light. In most organisms, photosynthesis takes place with the participation of chlorophyll; in the case of higher plants, photosynthesis is summarized by the equation:

CO 2 + H 2 O organic matter + O 2

The functioning of vision is also based on photochemical processes.

Chain reaction - a reaction, which is a chain of elementary acts of interaction, and the possibility of the occurrence of each act of interaction depends on the success of the passage of the previous act.

stages chain reaction:

The origin of the chain

chain development,

Chain break.

The origin of the chain occurs when, due to an external source of energy (quantum of electromagnetic radiation, heating, electric discharge), active particles with unpaired electrons (atoms, free radicals) are formed.

In the course of chain development, the radicals interact with the initial molecules, and new radicals are formed in each act of interaction.

Chain termination occurs if two radicals collide and transfer the energy released in this case to a third body (a molecule resistant to decay, or a vessel wall). The chain can also be terminated if an inactive radical is formed.

Two types chain reactions: unbranched and branched.

AT unbranched reactions at the stage of chain development, one new radical is formed from one reacting radical.

AT branched reactions at the chain development stage, more than one new radical is formed from one reacting radical.

6. Factors that determine the direction of a chemical reaction. Elements of chemical thermodynamics. Concepts: phase, system, environment, macro- and microstates. Basic thermodynamic characteristics. The internal energy of the system and its change in the course of chemical transformations. Enthalpy. The ratio of enthalpy and internal energy of the system. The standard enthalpy of a substance. Enthalpy change in systems during chemical transformations. Thermal effect (enthalpy) of a chemical reaction. Exo- and endothermic processes.

1. With. 89-97; 2. With. 158-163, p. 187-194; 3. With. 162-170; 4. With. 156-165; 5. With. 39-41; 6. With. 174-185; 8. With. 32-37.

Thermodynamics studies the patterns of energy exchange between the system and the environment, the possibility, direction and limits of the spontaneous flow of chemical processes.

Thermodynamic system(or simply system) – a body or a group of interacting bodies mentally distinguished in space. The rest of the space outside the system is called environment(or simply environment). The system is separated from the environment by a real or imaginary surface .

homogeneous system consists of one phase heterogeneous system- from two or more phases.

phasea –this is a part of the system, homogeneous at all its points in chemical composition and properties and separated from other phases of the system by the interface.

State system is characterized by the totality of its physical and chemical properties. macro state is determined by the averaged parameters of the entire set of particles of the system, and microstate- the parameters of each individual particle.

Independent variables that determine the macrostate of the system are called thermodynamic variables, or state parameters. Temperature is usually chosen as the state parameter. T, pressure R, volume V, chemical quantity n, concentration With etc.

A physical quantity, the value of which depends only on the state parameters and does not depend on the transition path to a given state, is called state function. The state functions are, in particular:

U- internal energy;

H- enthalpy;

S- entropy;

G- Gibbs energy (or free energy, or isobaric-isothermal potential).

Internal energy of the system U – this is its total energy, consisting of the kinetic and potential energy of all particles of the system (molecules, atoms, nuclei, electrons) without taking into account the kinetic and potential energy of the system as a whole. Since a full account of all these components is impossible, then in the thermodynamic study of the system, we consider change its internal energy during the transition from one state ( U 1) to another ( U 2):

U 1 U 2 DU = U 2 - U 1

The change in the internal energy of the system can be determined experimentally.

The system can exchange energy (heat Q) with the environment and do work BUT, or, conversely, work can be done on the system. According to first law of thermodynamics, which is a consequence of the law of conservation of energy, the heat received by the system can only be used to increase the internal energy of the system and to perform work by the system:

In the future, we will consider the properties of such systems, which are not affected by any other forces, except for the forces of external pressure.

If the process in the system proceeds at a constant volume (i.e., there is no work against the forces of external pressure), then A = 0. Then thermal effectprocess at constant volume, Q v is equal to the change in the internal energy of the system:

Q v = ΔU

Most chemical reactions encountered in everyday life take place at constant pressure ( isobaric processes). If no other forces act on the system, except for constant external pressure, then:

A \u003d p (V 2 -V 1) \u003d pDV

Therefore, in our case ( R= const):

Q p \u003d U 2 - U 1 + p (V 2 - V 1), whence

Q p \u003d (U 2 + pV 2) - (U 1 + pV 1)

Function U+PV, is called enthalpy; it is denoted by the letter H . Enthalpy is a state function and has the dimension of energy (J).

Q p \u003d H 2 - H 1 \u003d DH

Thermal effect of a reaction at constant pressure and temperature T is equal to the change in the enthalpy of the system during the reaction. It depends on the nature of the reactants and products, their physical state, conditions ( T, r) carrying out the reaction, as well as the amount of substances involved in the reaction.

Enthalpy of reactioncalled the change in the enthalpy of the system in which the reactants interact in amounts equal to the stoichiometric coefficients of the reaction equation.

The enthalpy of reaction is called standard, if the reactants and reaction products are in standard states.

The standard states are:

For a solid, an individual crystalline substance at 101.32 kPa,

For a liquid substance, the individual liquid substance at 101.32 kPa,

For a gaseous substance - gas at a partial pressure of 101.32 kPa,

For a solute, a substance in solution at a molality of 1 mol/kg, the solution being assumed to have the properties of an infinitely dilute solution.

The standard enthalpy of the reaction of formation of 1 mole of a given substance from simple substances is called standard enthalpy of formation this substance.

Recording example: D f H o 298(CO 2) \u003d -393.5 kJ / mol.

The standard enthalpy of formation of a simple substance that is in the most stable (for given p and T) state of aggregation is taken to be 0. If an element forms several allotropic modifications, then only the most stable one has zero standard enthalpy of formation (for given R and T) modification.

Usually, thermodynamic quantities are determined at standard conditions:

R= 101.32 kPa and T\u003d 298 K (25 ° C).

Chemical equations that indicate changes in enthalpy (heat effects of reactions) are called thermochemical equations. There are two forms of writing thermochemical equations in the literature.

Thermodynamic form of the thermochemical equation:

C (graphite) + O 2 (g) ® CO 2 (g); DH o 298= -393.5 kJ

The thermochemical form of the thermochemical equation for the same process:

C (graphite) + O 2 (g) ® CO 2 (g) + 393.5 kJ.

In thermodynamics, the thermal effects of processes are considered from the point of view of the system, therefore, if the system releases heat, then Q<0, а энтальпия системы уменьшается (ΔH< 0).

In classical thermochemistry, thermal effects are considered from the standpoint of the environment, therefore, if the system releases heat, then it is assumed that Q>0.

exothermic is a process that proceeds with the release of heat (ΔH<0).

endothermic a process proceeding with the absorption of heat (ΔH>0) is called.

The basic law of thermochemistry is Hess' law: the thermal effect of a reaction is determined only by the initial and final states of the system and does not depend on the path of the system's transition from one state to another.

Consequence from Hess' law : the standard thermal effect of the reaction is equal to the sum of the standard heats of formation of the reaction products minus the sum of the standard heats of formation of the starting substances, taking into account the stoichiometric coefficients:

DH o 298 (p-tion) = åD f H o 298 (cont.) –åD f H o 298 (outgoing)

7. The concept of entropy. Entropy change during phase transformations and chemical processes. The concept of the isobaric-isothermal potential of the system (Gibbs energy, free energy). The ratio between the magnitude of the change in the Gibbs energy and the magnitude of the change in the enthalpy and entropy of the reaction (basic thermodynamic relationship). Thermodynamic analysis of the possibility and conditions for the occurrence of chemical reactions. Features of the course of chemical processes in living organisms.

1. With. 97-102; 2. With. 189-196; 3. With. 170-183; 4. With. 165-171; 5. With. 42-44; 6. With. 186-197; 8. With. 37-46.

Entropy S- is a value proportional to the logarithm of the number of equiprobable microstates through which a given macrostate can be realized:

The unit of entropy is J/mol·K.

Entropy is a quantitative measure of the degree of disorder in a system.

Entropy increases during the transition of a substance from a crystalline state to a liquid and from a liquid to a gaseous state, during the dissolution of crystals, during the expansion of gases, during chemical interactions leading to an increase in the number of particles, and above all particles in the gaseous state. On the contrary, all processes, as a result of which the ordering of the system increases (condensation, polymerization, compression, reduction in the number of particles), are accompanied by a decrease in entropy.

There are methods for calculating the absolute value of the entropy of a substance, therefore, in the tables of thermodynamic characteristics of individual substances, data are given for S0, not for Δ S0.

The standard entropy of a simple substance, unlike the enthalpy of formation of a simple substance, is not equal to zero.

For entropy, a statement similar to that considered above for DH: the change in the entropy of the system as a result of a chemical reaction (DS) is equal to the sum of the entropies of the reaction products minus the sum of the entropies of the initial substances. As in the calculation of enthalpy, summation is carried out taking into account stoichiometric coefficients.

The direction in which a chemical reaction proceeds spontaneously is determined by the combined action of two factors: 1) the tendency for the system to transition to a state with the lowest internal energy (in the case of isobaric processes-with the lowest enthalpy) 2) the tendency to achieve the most probable state, i.e., the state that can be realized in the largest number of equiprobable ways (microstates):

Δ H → min,Δ S→max

The state function, which simultaneously reflects the influence of both tendencies mentioned above on the direction of chemical processes, is Gibbs energy (free energy , or isobaric-isothermal potential) , related to enthalpy and entropy by the relation

G=H-TS,

where T is the absolute temperature.

As you can see, the Gibbs energy has the same dimension as the enthalpy, and therefore is usually expressed in J or kJ.

For isobaric-isothermal processes, (i.e. processes occurring at constant temperature and pressure) the change in the Gibbs energy is equal to:

As in case D H and D S, Gibbs energy change D G as a result of a chemical reaction(Gibbs energy of the reaction) is equal to the sum of the Gibbs energies of the formation of reaction products minus the sum of the Gibbs energies of the formation of the initial substances; summation is carried out taking into account the number of moles of the substances involved in the reaction.

The Gibbs energy of formation of a substance is related to 1 mole of this substance and is usually expressed in kJ/mol; while D G 0 of the formation of the most stable modification of a simple substance is taken equal to zero.

At a constant temperature and pressure, chemical reactions can spontaneously proceed only in such a direction, in which the Gibbs energy of the system decreases ( D G<0).This is a condition for the fundamental possibility of implementing this process.

The table below shows the possibility and conditions for the reaction to proceed with various combinations of signs D H and D S.

By sign D G one can judge the possibility (impossibility) spontaneous leaks individual process. If the system is provided impact, then it is possible to carry out a transition from one substance to another, characterized by an increase in free energy (D G>0). For example, in the cells of living organisms reactions of formation of complex organic compounds proceed; the driving force of such processes are solar radiation and oxidation reactions in the cell.

• Activities and pedagogical views of Jan Amos Comenius
• The rate constant is calculated by the formula Rate constant value
• Phase diagrams of two-component systems "polymer - solvent" Phase diagrams of two-component systems "polymer - solvent"
• Fine and hyperfine structure of optical spectra
• How to learn chemistry on your own from scratch: effective ways
• Properties and characteristics of alkenes
• Characteristics of a student from the place of internship
• Characteristics for a student who had an internship at an enterprise: writing rules
• Biography of Faraday. Great scientists. Michael Faraday. discovery of electromagnetic rotation
• Modern innovations in education

• Substituents on the benzene ring are divided into two groups Reactions of the benzene ring
• Types of chemical reactions in organic chemistry Electrophilic addition reactions
• Methods for separating heterogeneous mixtures
• Polymethacrylic acid
• General characteristics of the elements of the VI A subgroup Elements 6a of the subgroup
• Theoretical foundations of physics
• Graph theory in chemistry. graph theory. Basic definitions and theorems of graph theory
• Algebra and harmony in chemical applications
• Tests on the course "Ecology and nature management
• WWF Russia Director Igor Chestin: Yakutia is among the leaders I know that you travel a lot… Why do you need this