1. Speed chemical reactions. Definition of the concept. 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. Rate constant, its physical meaning. The influence of the nature of the reactants, temperature and the presence of a catalyst on the reaction rate constant.
The rate of a homogeneous reaction is a value numerically equal to the change in the molar concentration of any reaction participant per unit time.
The average reaction speed v avg in the time interval from t 1 to t 2 is determined by the relation:
The main factors influencing the rate of a homogeneous chemical reaction:
- - the nature of the reacting substances;
- - molar concentrations of reagents;
- - pressure (if gases are involved in the reaction);
- - temperature;
- - presence of a catalyst.
The rate of a heterogeneous reaction is a value numerically equal to the change in the chemical amount of any reaction participant per unit time per unit interface area: .
According to the stages, chemical reactions are divided into simple (elementary) and complex. Most chemical reactions are complex processes occurring in several stages, i.e. consisting of several elementary processes.
For elementary reactions, the law of mass action is valid: the rate of an elementary chemical reaction is directly proportional to the product of the concentrations of the reactants in powers equal to the stoichiometric coefficients in the reaction equation.
For the elementary reaction aA + bB > ... the reaction rate, according to the law of mass action, is expressed by the relation:
where c(A) and c(B) are the molar concentrations of reactants A and B; a and b are the corresponding stoichiometric coefficients; k is the rate constant of this reaction.
For heterogeneous reactions, the equation of the law of mass action does not include the concentrations of all reactants, but only gaseous or dissolved ones. So, for the carbon combustion reaction:
C (k) + O 2 (g) > CO 2 (g)
The speed equation has the form: .
Physical meaning rate constant - 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 for a homogeneous reaction depends on the nature of the reactants, temperature and catalyst.
2. The influence 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. The relationship between activation energy and energy chemical bond in the original molecules. Transition state, or activated complex. Activation energy and thermal effect of the reaction (energy diagram). Dependence of the temperature coefficient of the reaction rate on the activation energy.
As 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, which is the same, by 10 K) is called the temperature coefficient of the rate of a chemical reaction (g):
where are the reaction rates at temperatures T 2 and T 1, respectively; r is the temperature coefficient of the reaction rate.
The dependence of the reaction rate on temperature is approximately determined by Van't Hoff's empirical rule: with every 10 degree 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 possible within the framework of the Arrhenius activation theory. According to this theory, a chemical reaction can occur when only active particles collide. Active particles are those that have a certain energy characteristic of a given reaction that 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.
An activated complex is an intermediate unstable group formed during the collision of active particles and is in a state of redistribution of bonds. When the activated complex decomposes, reaction products are formed.
The activation energy 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 bonds in the molecules of the reacting substances.
In activation theory, the effect of temperature on the rate of a chemical reaction is described by the Arrhenius equation for the rate constant of a chemical reaction:
where A is a constant factor independent of temperature, determined by the nature of the reacting substances; e is the base of the natural logarithm; E a - activation energy; R is the molar gas constant.
As follows from the Arrhenius equation, the lower the activation energy, the greater the reaction rate constant. Even a slight decrease in activation energy (for example, when adding a catalyst) 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 smaller the value of E a, the more noticeable the effect of temperature on the reaction rate and, therefore, the greater the temperature coefficient of the reaction rate.
3. The influence 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. The influence of a catalyst on the activation energy of a chemical reaction. Catalysis in nature, industry, technology. Biochemical catalysis. Enzymes.
Catalysis is a change in the rate of a chemical reaction under the influence of substances, the quantity and nature of which after the completion of the reaction remain the same as before the reaction.
A catalyst is a substance that changes the rate of a chemical reaction but remains chemically unchanged.
A positive catalyst speeds up the reaction; a 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 a reaction. Each of the intermediate processes involving a catalyst occurs with a lower activation energy than a non-catalyzed reaction.
In homogeneous catalysis, the catalyst and reactants form one phase (solution). In heterogeneous catalysis, the catalyst (usually a solid) and the reactants are in different phases.
During homogeneous catalysis, the catalyst forms an intermediate compound with the reagent, with high speed reacting with a second reagent or rapidly decomposing to release a reaction product.
An example of homogeneous catalysis: oxidation of sulfur(IV) oxide to sulfur(VI) oxide with oxygen using the nitrous method for producing sulfuric acid (here the catalyst is nitrogen(II) oxide, which easily reacts with oxygen).
In heterogeneous catalysis, the reaction occurs on the surface of the catalyst. The initial stages are the diffusion of reagent particles to the catalyst and their adsorption (i.e. absorption) by the surface of the catalyst. Reactant molecules interact with atoms or groups of atoms located on the surface of the catalyst, forming intermediate surface compounds. The redistribution of electron density that occurs in such intermediate compounds leads to the formation of new substances that are desorbed, i.e., removed from the surface.
The process of formation of intermediate surface compounds occurs at the active centers of the catalyst.
An example of heterogeneous catalysis is an increase in the rate of oxidation of sulfur(IV) oxide to sulfur(VI) oxide with oxygen in the presence of vanadium(V) oxide.
Examples of catalytic processes in industry and technology: ammonia synthesis, synthesis of nitric and sulfuric acids, cracking and reforming of oil, afterburning of products of incomplete combustion of gasoline in cars, etc.
Examples of catalytic processes in nature are numerous, since most biochemical reactions occurring in living organisms are classified as catalytic reactions. The catalysts for such reactions are protein substances called enzymes. There are about 30,000 enzymes in the human body, each of which catalyzes only one type of process (for example, salivary ptyalin catalyzes only the conversion of starch into glucose).
4. Chemical equilibrium. 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. The influence of changes in concentration, pressure and temperature on the position of chemical equilibrium.
Chemical reactions, as a result of which the starting substances are completely converted into reaction products, are called irreversible. Reactions occurring simultaneously in two opposite directions(direct and inverse) are called reversible.
In reversible reactions, the state of the system in which the rates of the forward and reverse reactions are equal () is called the state of chemical equilibrium. Chemical equilibrium is dynamic, i.e. its establishment does not mean the cessation of the reaction. In general, for any reversible reaction aA + bB - dD + eE, regardless of its mechanism, the following relation holds:
At established equilibrium, the product of the concentrations of the reaction products divided by the product of the concentrations of the starting substances for a given reaction at a given temperature is a constant value called the equilibrium constant (K).
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.
A change in 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 forward 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 written on the right side of the reaction equation increase, the equilibrium is said to shift to the right. If, during the transition from one equilibrium state to another, the concentrations of substances written on the left side of the reaction equation increase, the equilibrium is said to shift to the left.
The direction of the shift in 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 in a state of chemical equilibrium (change temperature, pressure or concentrations of substances), then it will favor the occurrence of whichever of the two opposite processes occurs. weakens this effect.
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 of equilibrium to the right; an increase in the concentration of the component written on the right side of the equation leads to a shift of equilibrium to the left;
When the temperature increases, the equilibrium shifts towards the endothermic reaction, and when the temperature decreases, towards the exothermic reaction;
- - As the pressure increases, the equilibrium shifts towards a reaction that reduces the number of molecules of gaseous substances in the system, and as the pressure decreases, towards a reaction that increases the number of molecules of gaseous substances.
- 5. Photochemical and chain reactions. Features of the course of photochemical reactions. Photochemical reactions and Live nature. Unbranched and branched chemical reactions (using the example of reactions of the formation of hydrogen chloride and water from simple substances). Conditions for the initiation and termination of chains.
Photochemical reactions are reactions that occur under the influence of light. A photochemical reaction occurs if the reagent absorbs radiation quanta characterized by an energy quite specific for a given reaction.
In the case of some photochemical reactions, absorbing energy, the molecules of the reagent pass into an excited state, i.e. become active.
In other cases, a photochemical reaction occurs if quanta of such high energy are absorbed that chemical bonds are broken and molecules dissociate into atoms or groups of atoms.
The greater the irradiation intensity, the greater the speed of the photochemical reaction.
An example of a photochemical reaction in living nature is photosynthesis, i.e. formation of organic cell substances due to light energy. In most organisms, photosynthesis occurs 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 processes is also based on photochemical processes.
A chain reaction is a reaction that is a chain of elementary acts of interaction, and the possibility of each act of interaction depends on the success of the previous act.
The stages of a chain reaction are chain initiation, chain development and chain termination.
The initiation of a chain occurs when, due to an external source of energy (quanta of electromagnetic radiation, heating, electrical discharge), active particles with unpaired electrons (atoms, free radicals) are formed.
During the development of the chain, radicals interact with the original molecules, and new radicals are formed in each act of interaction.
Chain termination occurs when two radicals collide and transfer the energy released in the process to a third body (a molecule resistant to decay or the wall of a vessel). The chain can also terminate if a low-active radical is formed.
There are two types of chain reactions - unbranched and branched.
In unbranched reactions at the stage of chain development, one new radical is formed from each reacting radical.
In branched reactions, at the stage of chain development, 2 or more new radicals are formed from one reacting radical.
6. Factors that determine the direction of a chemical reaction. Elements chemical thermodynamics. Concepts: phase, system, environment, macro- and microstates. Basic thermodynamic characteristics. Internal energy of the system and its change during chemical transformations. Enthalpy. The relationship between enthalpy and internal energy of a system. Standard enthalpy of a substance. Changes in enthalpy in systems during chemical transformations. Thermal effect (enthalpy) of a chemical reaction. Exo- and endothermic processes. Thermochemistry. Hess's law. Thermochemical calculations.
Thermodynamics studies the patterns of energy exchange between the system and the external environment, the possibility, direction and limits of the spontaneous occurrence of chemical processes.
A thermodynamic system (or simply a system) is a body or a group of interacting bodies, mentally isolated in space. The rest of the space outside the system is called environment(or just the environment). The system is separated from the environment by a real or imaginary surface.
A homogeneous system consists of one phase, a heterogeneous system consists of two or more phases.
A phase is a part of a system that is homogeneous at all its points along chemical composition and properties and separated from other parts of the system by an interface.
The state of a system is characterized by the totality of its physical and chemical properties. The macrostate is determined by the averaged parameters of the entire set of particles in the system, and the microstate by the parameters of each individual particle.
The independent variables that determine the macrostate of the system are called thermodynamic variables, or state parameters. Temperature T, pressure p, volume V, chemical quantity n, concentration c, etc. are usually chosen as state parameters.
A physical quantity, the value of which depends only on the parameters of the state and does not depend on the path of transition to a given state, is called a state function. The functions of the state are, in particular:
U - internal energy;
H - enthalpy;
S - entropy;
G - Gibbs energy (free energy or isobaric-isothermal potential).
The internal energy of a system U 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, when studying a system thermodynamically, we consider the change in its internal energy during the transition from one state (U 1) to another (U 2):
U 1 U 2 U = U 2 - U1
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 A, or, conversely, work can be done on the system. According to the 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:
Q= U+A
In the future, we will consider the properties of such systems that are not affected by any forces other than external pressure forces.
If a process occurs in a system at a constant volume (i.e., there is no work against external pressure forces), then A = 0. Then the thermal effect of a process occurring at a constant volume, Q v, is equal to the change in the internal energy of the system:
Most chemical reactions encountered in everyday life occur at constant pressure (isobaric processes). If no forces other than constant external pressure act on the system, then:
A = p(V2 - V 1 ) = pV
Therefore, in our case (p = const):
Qp=U + pV
Q р = U 2 -U 1 + p(V 2 - V 1 ), where
Q p = (U 2 +pV 2 ) - (U 1 +pV 1 ).
The function U + pV is called enthalpy; it is denoted by the letter N. Enthalpy is a function of state and has the dimension of energy (J).
Qp= H 2 - H 1 = H,
i.e., the thermal effect of a reaction at constant pressure and temperature T is equal to the change in enthalpy of the system during the reaction. It depends on the nature of the reagents and products, their physical state, conditions (T, p) for the reaction, as well as on the amount of substances participating in the reaction.
The enthalpy of a reaction is the change in enthalpy of a system in which the reactants react in quantities equal to the stoichiometric coefficients in the reaction equation.
The enthalpy of a reaction is called standard if the reactants and products of the reaction are in standard states.
The standard state of a substance is the aggregate state or crystalline form of a substance in which it is thermodynamically most stable under standard conditions (T = 25 o C or 298 K; p = 101.325 kPa).
The standard state of a substance existing at 298 K in solid form is considered to be its pure crystal under a pressure of 101.325 kPa; in liquid form - pure liquid under a pressure of 101.325 kPa; in gaseous form - gas with its own pressure of 101.325 kPa.
For a solute, the standard state is considered to be its state in solution at a molality of 1 mol/kg, and it is assumed that the solution has 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 in their standard states is called the standard enthalpy of formation of this substance.
Example entry: (CO 2) = - 393.5 kJ/mol.
Standard enthalpy of formation simple substance, which is in the most stable (given p and T) state of aggregation, is assumed to be equal to 0. If an element forms several allotropic modifications, then only the most stable (given p and T) modification has a zero standard enthalpy of formation.
Typically, thermodynamic quantities are determined under standard conditions:
p = 101.32 kPa and T = 298 K (25 o C).
Chemical equations that specify enthalpy changes (heat effects of reactions) are called thermochemical equations. In the literature you can find two forms of writing thermochemical equations.
Thermodynamic form of writing the thermochemical equation:
C (graphite) + O 2 (g) CO 2 (g); = - 393.5 kJ.
Thermochemical form of writing the thermochemical equation of 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 perspective of the system. Therefore, if the system releases heat, then Q< 0, а энтальпия системы уменьшается (ДH < 0).
In classical thermochemistry, thermal effects are considered from an environmental perspective. Therefore, if the system releases heat, then it is assumed that Q > 0.
Exothermic is a process that occurs with the release of heat (DH< 0).
Endothermic is a process that occurs with the absorption of heat (DH > 0).
The basic law of thermochemistry is Hess's 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 transition of the system from one state to another.”
Corollary from Hess's law: The standard thermal effect of a 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 stoichiometric coefficients:
- (reactions) = (cont.) -(out.)
- 7. The concept of entropy. Changes in entropy during phase transformations and chemical processes. The concept of the isobaric-isothermal potential of a system (Gibbs energy, free energy). The relationship between the magnitude of the change in the Gibbs energy and the magnitude of the change in enthalpy and entropy of the reaction (basic thermodynamic relationship). Thermodynamic analysis of the possibility and conditions of chemical reactions. Features of the flow of chemical processes in living organisms.
Entropy S is a value proportional to the logarithm of the number of equally probable microstates (W) through which a given macrostate can be realized:
S = k lnW
The unit of entropy is J/mol?K.
Entropy is a quantitative measure of the degree of disorder of a system.
Entropy increases when a substance moves from crystalline state into liquid and from liquid into gaseous, 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 order 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 the thermodynamic characteristics of individual substances, data are given for S 0, and not for DS 0.
Standard entropy of a simple substance, as opposed to enthalpy of formation simple substance is not equal to zero.
For entropy, a statement similar to that discussed above for H is valid: the change in the entropy of a system as a result of a chemical reaction (S) is equal to the sum of the entropies of the reaction products minus the sum of the entropies of the starting substances. As with the calculation of enthalpy, the summation is carried out taking into account the stoichiometric coefficients.
The direction in which a chemical reaction spontaneously occurs in an isolated system 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) a tendency to achieve the most probable state, i.e. a state that can be realized in the largest number of equally probable ways (microstates), i.e.:
DH > min, DS > max.
The state function, which simultaneously reflects the influence of both of the above-mentioned trends on the direction of the flow of chemical processes, is the Gibbs energy (free energy, or isobaric-isothermal potential), related to enthalpy and entropy by the relation
where T is the absolute temperature.
As can be seen, the Gibbs energy has the same dimension as enthalpy and is therefore usually expressed in J or kJ.
For isobaric-isothermal processes (i.e. processes occurring at constant temperature and pressure), the change in Gibbs energy is equal to:
G= H-TS
As in the case of H and S, the change in Gibbs energy G as a result of a chemical reaction (Gibbs energy of the reaction) is equal to the sum of the Gibbs energies of formation of the reaction products minus the sum of the Gibbs energies of formation of the starting materials; the summation is made taking into account the number of moles of substances participating in the reaction.
The Gibbs energy of formation of a substance is referred to 1 mole of this substance and is usually expressed in kJ/mol; in this case, G 0 of the formation of the most stable modification of a simple substance is taken equal to zero.
At constant temperature and pressure, chemical reactions can spontaneously proceed only in a direction in which the Gibbs energy of the system decreases (G0). This is a condition for the fundamental possibility of carrying out this process.
The table below shows the possibility and conditions for the reaction to occur with various combinations of H and S signs:
By the sign of G one can judge the possibility (impossibility) of the spontaneous occurrence of a particular process. If the system is influenced, then it is possible to carry out a transition from one substance to another, characterized by an increase in free energy (G>0). For example, in the cells of living organisms reactions occur to form complex organic compounds; The driving force behind such processes is solar radiation and oxidation reactions in the cell.
Methodical advice
(L.1, pp. 168-210)
Thermochemistry studies the thermal effects of chemical reactions. Thermochemical calculations are based on the application of Hess's law. Based on this law, it is possible to calculate the thermal effects of reactions using tabular data (Appendix, Table 3). It should be noted that thermochemical tables are usually constructed on the basis of data for simple substances, the heats of formation of which are assumed to be zero.
Thermodynamics develops the general laws of the occurrence of chemical reactions. These patterns can be quantified by the following thermodynamic quantities: internal energy of the system (U), enthalpy (H), entropy (S) and isobaric-isothermal potential (G - Gibbs free energy).
The study of the rate of chemical reactions is called chemical kinetics. The central issues of this topic are the law of mass action and chemical equilibrium. Pay attention to the fact that the study of the rate of chemical reactions and chemical equilibrium is of great importance, since it allows you to control the flow of chemical reactions.
Theoretical aspects
4.1 Chemical thermodynamics
Chemical thermodynamics - the science of the dependence of the direction and limits of transformations of substances on the conditions in which these substances are located.
Unlike other branches of physical chemistry (structure of matter and chemical kinetics), chemical thermodynamics can be applied without knowing anything about the molecular structure of matter. Such a description requires significantly less initial data.
Example:
The enthalpy of glucose formation cannot be determined by direct experiment:
6 C + 6 H 2 + 3 O 2 = C 6 H 12 O 6 (H x -?) such a reaction is impossible
6 CO 2 + 6 H 2 O = C 6 H 12 O 6 + 6 O 2 (H y - ?) the reaction occurs in green leaves, but together with other processes.
Using Hess's law, it is enough to combine three combustion equations:
1) C + O 2 = CO 2 H 1 = -394 kJ
2) H 2 + 1/2 O 2 = H 2 O (steam) H 2 = -242 kJ
3) C 6 H 12 O 6 + 6 O 2 = 6 CO 2 + 6 H 2 O H 3 = -2816 kJ
We add the equations, “expanding” the third, then
H x = 6 H 1 + 6 H 2 - H 3 = 6(-394) + 6(-242) -(-2816) = -1000 kJ/mol
The solution did not use any data on the structure of glucose; The mechanism of its combustion was also not considered.
The isobaric potential is expressed in kJ/mol. Its change during a chemical reaction does not depend on the path of the reaction, but is determined only by the initial and final states of the reacting substances (Hess’s law):
ΔG reaction = Σ ΔG final product - Σ ΔG starting materials
Specific object of thermodynamic research called a thermodynamic system, isolated from the surrounding world by real or imaginary surfaces. A system can be a gas in a vessel, a solution of reagents in a flask, a crystal of a substance, or even a mentally isolated part of these objects.
If the system has real interface, separating parts of the system from each other that differ in properties, then the system is called heterogeneous(saturated solution with sediment), if there are no such surfaces, the system is called homogeneous(true solution). Heterogeneous systems contain at least two phases.
Phase– a set of all homogeneous parts of the system, identical in composition and in all physical and chemical properties(independent of the amount of substance) and delimited from other parts of the system by an interface. Within one phase, properties can change continuously, but at the interface between phases, properties change abruptly.
Components name the substances that are minimally necessary to compose a given system (at least one). The number of components in a system is equal to the number of substances present in it, minus the number of independent equations connecting these substances.
According to the levels of interaction with the environment, thermodynamic systems are usually divided into:
– open – exchange matter and energy with the environment (for example, living objects);
– closed – exchange only energy (for example, a reaction in a closed flask or a flask with reflux), the most common object of chemical thermodynamics;
– isolated – do not exchange either matter or energy and maintain a constant volume (approximation – reaction in a thermostat).
The properties of the system are divided into extensive (summing) - for example, total volume, mass, and intensive (levelling) - pressure, temperature, concentration, etc. The set of properties of a system determines its state. Many properties are interrelated, therefore, for a homogeneous one-component system with a known amount of substance n, it is enough to choose the state to characterize two out of three properties: temperature T, pressure p and volume V. The equation connecting the properties is called the equation of state; for an ideal gas it is:
Laws of thermodynamics
First law of thermodynamics:Energy is neither created nor destroyed. A perpetual motion machine (perpetuum mobile) of the first kind is impossible. In any isolated system the total amount of energy is constant.
In general, the work done by a chemical reaction at constant pressure (isobaric process) consists of a change in internal energy and work of expansion:
For most chemical reactions carried out in open vessels, it is convenient to use state function, the increment of which is equal to the heat received by the system in an isobaric process. This function is called enthalpy(from the Greek “enthalpo” - heat):
Another definition: the enthalpy difference in the two states of the system is equal to the thermal effect of the isobaric process.
There are tables containing data on the standard enthalpies of formation of substances H o 298. The indices mean that for chemical compounds the enthalpy of formation of 1 mole of them from simple substances taken in the most stable modification (except for white phosphorus - not the most stable, but the most reproducible form of phosphorus) at 1 atm (1.01325∙10 5 Pa or 760 mmHg) and 298.15 K (25 o C). If we are talking about ions in solution, then the standard concentration is 1M (1 mol/l).
The sign of enthalpy is determined “from the point of view” of the system itself: when heat is released, the change in enthalpy is negative, when heat is absorbed, the change in enthalpy is positive.
Second law of thermodynamics
Change entropy equal (by definition) to the minimum heat supplied to the system in a reversible (all intermediate states are in equilibrium) isothermal process, divided by the absolute temperature of the process:
S = Q min. /T
At this stage of studying thermodynamics, it should be accepted as a postulate that there is some extensive property of the system S, called entropy, the change of which is so connected with the processes in the system:
In a spontaneous process S > Q min. /T
In the equilibrium process S = Q min. /T
< Q мин. /T
For an isolated system, where dQ = 0, we obtain:
In a spontaneous process S > 0
In the equilibrium process S = 0
In a non-spontaneous process S< 0
In general the entropy of an isolated system either increases or remains constant:
The concept of entropy arose from earlier formulations of the second law (beginning) of thermodynamics. Entropy is a property of the system as a whole, and not of an individual particle.
Third law of thermodynamics (Planck's postulate)
The entropy of a properly formed crystal of a pure substance at absolute zero is zero(Max Planck, 1911). This postulate can be explained by statistical thermodynamics, according to which entropy is a measure of the disorder of a system at the micro level:
S = k b lnW - Boltzmann equation
W is the number of different states of the system available to it under given conditions, or the thermodynamic probability of the macrostate of the system.
k b = R/N A = 1.38. 10 -16 erg/deg – Boltzmann constant
In 1872, L. Boltzmann proposed a statistical formulation of the second law of thermodynamics: an isolated system evolves predominantly in the direction of a higher thermodynamic probability.
The introduction of entropy made it possible to establish criteria that make it possible to determine the direction and depth of any chemical process (for large number particles in equilibrium).
Macroscopic systems reach equilibrium when the change in energy is compensated by the entropy component:
At constant volume and temperature:
U v = TS v or (U-TS) =F = 0 - Helmholtz energy or isochoric-isothermal potential
At constant pressure and temperature:
H p = TS p or (H-TS) =G = 0 - Gibbs energy or Gibbs free energy or isobaric-isothermal potential.
Change in Gibbs energy as a criterion for the possibility of a chemical reaction: G =H - TS
At G< 0 реакция возможна;
at G > 0 the reaction is impossible;
at G = 0 the system is in equilibrium.
The possibility of a spontaneous reaction in an isolated system is determined by a combination of the signs of the energy (enthalpy) and entropy factors:
Extensive tabular data are available for standard values of G 0 and S 0 to allow calculation of the G 0 reaction.
If the temperature differs from 298 K and the concentration of reagents from 1M, for the process in general form:
G = G 0 + RT ln([C] c [D] d /[A] a [B] b)
In the equilibrium position G = 0 and G 0 = -RTlnK р, where
K р = [C] c equals [D] d equals /[A] a equals [B] b equals equilibrium constant
K р = exp (-G˚/RT)
Using the above formulas, it is possible to determine the temperature from which the endothermic reaction, at which entropy increases, becomes easily feasible. The temperature is determined from the condition.
1 . What does chemical thermodynamics study:
1) the rate of chemical transformations and the mechanisms of these transformations;
2) energy characteristics of physical and chemical processes and the ability of chemical systems to perform useful work;
3) conditions for shifting chemical equilibrium;
4) the influence of catalysts on the rate of biochemical processes.
2. An open system is a system that:
2) exchanges both matter and energy with the environment;
3. A closed system is a system that:
1) does not exchange either matter or energy with the environment;
3) exchanges energy with the environment, but does not exchange matter;
4) exchanges matter with the environment, but does not exchange energy.
4. An isolated system is a system that:
1) does not exchange either matter or energy with the environment;
2) exchanges both matter and energy with the environment;
3) exchanges energy with the environment, but does not exchange matter;
4) exchanges matter with the environment, but does not exchange energy.
5. To what type of thermodynamic systems does the solution located in a sealed ampoule placed in a thermostat belong?
1) isolated;
2) open;
3) closed;
4) stationary.
6. What type of thermodynamic systems does the solution in the sealed ampoule belong to?
1) isolated;
2) open;
3) closed;
4) stationary.
7. What type of thermodynamic systems does a living cell belong to?
1) open;
2) closed;
3) isolated;
4) equilibrium.
8 . What parameters of a thermodynamic system are called extensive?
1) the magnitude of which does not depend on the number of particles in the system;
2) whose value depends on the number of particles in the system;
3) the value of which depends on the state of aggregation of the system;
9. What parameters of a thermodynamic system are called intensive?
!) the magnitude of which does not depend on the number of particles in the system;
2) the magnitude of which depends on the number of particles in the system;
3) the value of which depends on the state of aggregation;
4) the magnitude of which depends on time.
10 . Functions of the state of a thermodynamic system are quantities that:
1) depend only on the initial and final state of the system;
2) depend on the process path;
3) depend only on the initial state of the system;
4) depend only on the final state of the system.
11 . What quantities are functions of the state of the system: a) internal energy; b) work; c) warmth; d) enthalpy; d) entropy.
1) a, d, e;
3) all quantities;
4) a, b, c, d.
12 . Which of the following properties are intensive: a) density; b) pressure; c) mass; d) temperature; e) enthalpy; e) volume?
1) a, b, d;
3) b, c, d, f;
13. Which of the following properties are extensive: a) density; b) pressure; c) mass; d) temperature; e) enthalpy; e) volume?
1) c, d, f;
3) b, c, d, f;
14 . What forms of energy exchange between the system and the environment are considered by thermodynamics: a) heat; b) work; c) chemical; d) electric; e) mechanical; e) nuclear and solar?
1)a, b;
2) c, d, e, f;
3) a, c, d, e, f;
4) a, c, d, e.
15. Processes occurring at a constant temperature are called:
1) isobaric;
2) isothermal;
3) isochoric;
4) adiabatic.
16 . Processes occurring at constant volume are called:
1) isobaric;
2) isothermal;
3) isochoric;
4) adiabatic.
17 . Processes occurring at constant pressure are called:
1) isobaric;
2) isothermal;
3) isochoric;
4) adiabatic.
18 . The internal energy of the system is: 1) the entire energy reserve of the system, except for the potential energy of its position andkinetic energysystems as a whole;
2) the entire energy reserve of the system;
3) the entire energy reserve of the system, except for the potential energy of its position;
4) a quantity characterizing the degree of disorder in the arrangement of particles of the system.
19 . What law reflects the relationship between work, heat and internal energy of a system?
1) second law of thermodynamics;
2) Hess's law;
3) the first law of thermodynamics;
4) van't Hoff's law.
20 . The first law of thermodynamics reflects the relationship between:
1) work, heat and internal energy;
2) free energy Gibbs, enthalpy and entropy of the system;
3) work and heat of the system;
4) work and internal energy.
21 . Which equation is the mathematical expression of the first law of thermodynamics for isolated systems?
l)AU=0 2)AU=Q-p-AV 3)AG = AH-TAS
22 . Which equation is the mathematical expression of the first law of thermodynamics for closed systems?
2)AU=Q-p-AV;
3) AG = AH - T*AS;
23 . Is the internal energy of an isolated system a constant or variable quantity?
1) constant;
2) variable.
24 . In an isolated system, the reaction of hydrogen combustion occurs with the formation of liquid water. Does the internal energy and enthalpy of the system change?
1) internal energy will not change, enthalpy will change;
2) internal energy will change, enthalpy will not change;
3) internal energy will not change, enthalpy will not change;
4) internal energy will change, enthalpy will change.
25 . Under what conditions is the change in internal energy equal to the heat received by the system from the environment?
1) at constant volume;
3) at constant pressure;
4) under no circumstances.
26 . The thermal effect of a reaction occurring at constant volume is called a change:
1) enthalpy;
2) internal energy;
3) entropy;
4) Gibbs free energy.
27 . The enthalpy of a reaction is:
1) the amount of heat that is released or absorbed during a chemical reaction under isobaric-isothermal conditions;
4) a quantity characterizing the degree of disorder in the arrangement and movement of particles in the system.
28. Chemical processes during which the enthalpy of the system decreases and heat is released into the external environment are called:
1) endothermic;
2) exothermic;
3) exergonic;
4) endergonic.
29 . Under what conditions is the change in enthalpy equal to the heat received by the system from the environment?
1) at constant volume;
2) at constant temperature;
3) at constant pressure;
4) under no circumstances.
30 . The thermal effect of a reaction occurring at constant pressure is called a change:
1) internal energy;
2) none of the previous definitions are correct;
3) enthalpy;
4) entropy.
31. What processes are called endothermic?
1) for which AN is negative;
3) for whichANpositively;
32 . What processes are called exothermic?
1) for whichANnegative;
2) for which AG is negative;
3) for which AN is positive;
4) for which AG is positive.
33 . Specify the formulation of Hess's law:
1) the thermal effect of the reaction depends only on the initial and final state of the system and does not depend on the reaction path;
2) the heat absorbed by the system at a constant volume is equal to the change in the internal energy of the system;
3) the heat absorbed by the system at constant pressure is equal to the change in enthalpy of the system;
4) the thermal effect of the reaction does not depend on the initial and final state of the system, but depends on the reaction path.
34. What law underlies the calculation of caloric content of food?
1) van't Hoff;
2) Hess;
3) Sechenov;
35. When oxidizing which substances under body conditions, more energy is released?
1) proteins;
2) fat;
3) carbohydrates;
4) carbohydrates and proteins.
36 . A spontaneous process is a process that:
1) carried out without the help of a catalyst;
2) accompanied by the release of heat;
3) carried out without external energy consumption;
4) proceeds quickly.
37 . The entropy of a reaction is:
1) the amount of heat that is released or absorbed during a chemical reaction under isobaric-isothermal conditions;
2) the amount of heat that is released or absorbed during a chemical reaction under isochoric-isothermal conditions;
3) a value characterizing the possibility of spontaneous occurrence of the process;
4) a quantity characterizing the degree of disorder in the arrangement and movement of particles in a system.
38 . What state function characterizes the tendency of a system to achieve a probable state that corresponds to the maximum randomness of the distribution of particles?
1) enthalpy;
2) entropy;
3) Gibbs energy;
4) internal energy.
39 . What is the relationship between the entropies of three aggregate states of one substance: gas, liquid, solid:
I) S(d) >S(g) >S(TV); 2) S(solid)>S(g)>S(g); 3)S(g)>S(g)>S(TB); 4) the state of aggregation does not affect the entropy value.
40 . Which of the following processes should exhibit the greatest positive change in entropy:
1) CH3OH (s) --> CH,OH (g);
2) CH3OH (s) --> CH 3 OH (l);
3) CH,OH (g) -> CH3OH (s);
4) CH,OH (l) -> CH3OH (sol).
41 . Choose the correct statement: the entropy of the system increases when:
1) increase in pressure;
2) transition from liquid to solid state of aggregation
3) temperature increase;
4) transition from gaseous to liquid state.
42. What thermodynamic function can be used to predict whether a reaction will occur spontaneously in an isolated system?
1) enthalpy;
2) internal energy;
3) entropy;
4) potential energy of the system.
43 . Which equation is the mathematical expression of the 2nd law of thermodynamics for isolated systems?
2)AS>Q\T
44 . If the system reversibly receives a quantity of heat Q at temperature T, then about T;
2) increases by the amountQ/ T;
3) increases by an amount greater than Q/T;
4) increases by an amount less than Q/T.
45 . In an isolated system, a chemical reaction occurs spontaneously to form a certain amount of product. How does the entropy of such a system change?
1) increases
2) decreases
3) does not change
4) reaches the minimum value
46 . Indicate in which processes and under what conditions the change in entropy can be equal to the work of the process?
1) in isobaric conditions, at constant P and T;
2) in isochoric, at constant Vi and T;
H) the change in entropy is never equal to work;
4) in isothermal conditions, at constant P and 47 . How will the bound energy of the system TS change when heated and when it condenses?
Basic concepts and laws of chemistry. Chemical bond. Structure and properties of matter
1. What substances are called simple? Difficult? From the given substances, select simple ones: CO, O 3, CaO, K, H 2, H 2 O.
2. What substances are called oxides? Acids? Reasons? Salts?
3. From the given oxides - SO 2, CaO, ZnO, Cr 2 O 3, CrO, P 2 O 5, CO 2, Cl 2 O 3, Al 2 O 3 - select basic, acidic and amphoteric.
4. Which salts are classified as acidic, basic, medium, double, mixed, complex?
5. Name the following compounds: ZnOHCl, KHSO 3, NaAl(SO 4) 2. What class of compounds do they belong to?
6. What is the basicity of an acid called?
7. From the given hydroxides, select amphoteric ones: Fe(OH) 2 , KOH, Al(OH) 3 , Ca(OH) 2 , Fe(OH) 3 , Pb(OH) 2 .
8. What is called a reaction scheme? Reaction equation?
9. What are the numbers in the reaction equation called? What do they show?
10. How to move from a reaction diagram to an equation?
11. What substances do basic oxides interact with? Amphoteric oxides? Acidic oxides?
12. What substances do bases interact with?
13. What substances do acids interact with?
14. What substances do salts interact with?
15. Determine the mass fractions of elements in nitric acid HNO 3.
16. What metals react with alkalis?
17. What metals react with solutions of sulfuric and hydrochloric acids?
18. What products are formed when metals react with nitric acid of varying concentrations?
19. What reactions are called decomposition reactions? Connections? Substitutions? Redox?
20. Write down the reaction equations: CrCl 3 + NaOH→; CrCl 3 + 2NaOH→; CrCl 3 + 3NaOH→; CrCl 3 + NaOH (excess) →.
21. Write down the reaction equations: Al + KOH →; Al + KOH + H 2 O →.
22. What is called an atom? Chemical element? A molecule?
23. What elements are classified as metals? Non-metals? Why?
24. What is the chemical formula of a substance? What does it show?
25. What is the structural formula of a substance called? What does it show?
26. What is called the amount of a substance?
27. What is called a mole? What does it show? How many structural units are contained in a mole of the substance?
28. What masses of elements are indicated in Periodic table?
29. What are called relative atomic, molecular weight? How are they determined? What are their units of measurement?
30. What is the molar mass of a substance called? How is it defined? What is its unit of measurement?
31. What conditions are called normal conditions?
32. What volume does 1 mole of gas occupy at zero level? 5 moles of gas at ambient conditions?
33. What does an atom consist of?
34. What does the nucleus of an atom consist of? What charge does the nucleus of an atom have? What determines the charge of an atomic nucleus? What determines the mass of the nucleus of an atom?
35. What is called mass number?
36. What is called the energy level? How many electrons are there in a particular energy level?
37. What is called an atomic orbital? How is she portrayed?
38. What characterizes the principal quantum number? Orbital quantum number? Magnetic quantum number? Spin quantum number?
39. What is the relationship between the principal and orbital quantum numbers? Between orbital and magnetic quantum numbers?
40. What are electrons with = 0 called? = 1? = 2? = 3? How many orbitals correspond to each of these electron states?
41. What state of the atom is called the ground state? Excited?
42. How many electrons can be located in one atomic orbital? What is the difference?
44. How many and what sublevels can be located at the first energy level? On the second? On the third? On the fourth?
45. Formulate the principle of least energy, Klechkovsky’s rules, Pauli’s principle, Hund’s rule, periodic law.
46. What changes periodically for atoms of elements?
47. What do elements of one subgroup have in common? One period?
48. How do the elements of the main subgroups differ from the elements side subgroups?
49. Make up electronic formulas for the ions Cr +3, Ca +2, N -3. How many unpaired electrons do these ions have?
50. What energy is called ionization energy? Electron affinity? Electronegativity?
51. How do the radii of atoms and ions change in a group and in a period of the Periodic System D.I. Mendeleev?
52. How do the electronegativity of atoms change in a group and in a period of the Periodic System D.I. Mendeleev?
53. How do the metallic properties of elements and the properties of their compounds change in a group and in a period of the Periodic System D.I. Mendeleev?
54. Make up formulas for higher oxides of aluminum, phosphorus, bromine, and manganese.
55. How is the number of protons, neutrons and electrons in an atom determined?
56. How many protons, neutrons and electrons are there in a zinc atom?
57. How many electrons and protons are contained in the ions Cr +3, Ca +2, N -3?
58. Formulate the law of conservation of mass? What remains constant during any chemical reaction?
59. Which parameter remains constant in isobaric chemical reactions?
60. Formulate the law of constancy of composition. For substances of what structure is this true?
61. Formulate Avogadro’s law and consequences from it.
62. If the gas density for nitrogen is 0.8, then what is the molar mass of the gas?
63. If what external parameters change, does the molar volume of the gas change?
64. Formulate the combined gas law.
65. For equal volumes of different gases under the same conditions, will the masses of the gases be equal?
66. Formulate Dalton's law. If the total pressure of a mixture of nitrogen and hydrogen is 6 atm, and the volumetric content of hydrogen is 20%, then what are the partial pressures of the components?
67. Write down the Mendeleev-Clapeyron equation (state of an ideal gas).
68. What is the mass of a mixture of gases consisting of 11.2 liters of nitrogen and 11.2 liters of fluorine (n.s.)?
69. What is called a chemical equivalent? Molar mass equivalent?
70. How are the molar masses of equivalents of simple and complex substances determined?
71. Determine the molar masses of equivalents of the following substances: O 2, H 2 O, CaCl 2, Ca(OH) 2, H 2 S.
72. Determine the equivalent of Bi(OH) 3 in the reaction Bi(OH) 3 + HNO 3 = Bi(OH) 2 (NO 3) + H 2 O.
73. Formulate the law of equivalents.
74. What is the molar volume of an equivalent substance called? How is it defined?
75. Formulate the law of volumetric relations.
76. What volume of oxygen will be required for the oxidation of 8 m 3 of hydrogen (n.s.) according to the reaction 2H 2 + O 2 ↔ 2H 2 O?
77. What volume of hydrogen chloride is formed by the interaction of 15 liters of chlorine and 20 liters of hydrogen?
78. What is meant by a chemical bond? Specify the characteristics of a chemical bond.
79. What is a measure of the strength of a chemical bond?
80. What affects the distribution of electron density?
81. What determines the shape of a molecule?
82. What is called valency?
83. Determine the valency of nitrogen in the following compounds: N 2, NH 3, N 2 H 4, NH 4 Cl, NaNO 3.
84. What is the oxidation state called?
85. What bond is called covalent?
86. Specify the properties of a covalent bond.
87. How does the polarity of a bond change in the series KI, KBr, KCl, KF?
88. The molecules of which substance are non-polar: oxygen, hydrogen chloride, ammonia, acetic acid.
89. What is meant by hybridization of valence orbitals?
90. Determine the types of hybridization of central atoms in the following substances: beryllium fluoride, aluminum chloride, methane.
91. How does the type of hybridization affect the spatial structure of molecules?
92. What bond is called ionic? Under the influence of what forces does it arise?
93. What kind of bond is called metallic?
94. What properties do substances with a metal type of chemical bond have?
95. What is the maximum number of -bonds that can form between two atoms in a molecule?
96. How is the absolute electronegativity of an atom of an element determined?
97. Arrange the elements in increasing order of their electronegativity: Fe, C, Ag, H, Cl.
98. What is called the coupling dipole moment? How is it calculated?
99. What features do substances with an atomic crystal lattice have? With a molecular crystal lattice?
100.What kind of bond is called a hydrogen bond? What does its strength depend on? Between the molecules of what inorganic substances does it occur?
Thermodynamics and kinetics of chemical reactions
1. What does thermodynamics study?
2. What is called a thermodynamic system? What types of systems exist?
3. What are called state parameters? What parameters are called intensive, extensive? Name the main parameters of a chemical system.
4. What is a process called? Spontaneous process? Cycle? An equilibrium process? A nonequilibrium process? A reversible process?
5. What is a phase called? Homogeneous, heterogeneous system?
6. What is a state function called?
7. What characterizes the internal energy U? What does internal energy depend on?
8. What is heat Q called? Which reactions are exothermic and endothermic? How do heat and enthalpy change during their course?
9. What is work p∆V called?
10. Formulate the first law of thermodynamics. Write it down mathematically.
11. Formulate the first law of thermodynamics for isothermal, isochoric and isobaric processes.
12. What is called enthalpy?
13. What is called the thermal effect of a reaction? What determines the thermal effect of a reaction?
14. Which equation is called thermodynamic? Thermochemical?
15. What conditions are called standard?
16. What is called the enthalpy of a reaction? Standard enthalpy of reaction?
17. What is the enthalpy of formation of a substance called? Standard enthalpy of formation of a substance?
18. What state of matter is standard? What is the enthalpy of formation of a simple substance in the standard state?
19. The enthalpy of formation of H 2 SO 3 is equal in magnitude to the thermal effect of the reaction: H 2 (g) + S (solid) + 1.5O 2 (g) H 2 SO 3 (l); H 2 (g) + SO 2 (g) + 0.5 O 2 (g) H 2 SO 3 (l); H 2 O (g) + SO 2 (g) H 2 SO 3 (l); 2H (g) + S (s) + 3O (g) H 2 SO 3 (l).
20. When 1 mole of hydrogen and 1 mole of bromine react, 500 kJ of heat is released. What is ∆Н arr., HBr?
21. When 5 moles of substance A x B y are formed, 500 kJ of heat is absorbed. What is ∆Н arr of this substance?
22. What is the enthalpy of combustion called? Standard enthalpy of combustion? Calorific value?
23. Formulate Hess’s law, the first and second consequences from it.
24. What expression is applicable for calculating ∆Н р reaction 2A + 3B 2C according to Hess's law:
∆Н р = 2∆Н arr, C + 2∆Н arr, A + 3∆Н arr, B; ∆Н р = 2∆Н arr, C – (2∆Н arr, A + 3∆Н arr, B);
∆Н р = 2∆Н arr, A + 3∆Н arr, B –2∆Н arr, C; ∆Н р = – 2∆Н arr, C – (2∆Н arr, A + 3∆Н arr, B)?
25. The standard enthalpy of combustion (∆H 0 combustion) of methanol CH 4 O (l) (M = 32 g/mol) is equal to -726.6 kJ/mol. How much heat will be released during the combustion of 2.5 kg of substance?
26. In what case is the standard enthalpy of combustion of one substance equal to the standard enthalpy of formation of another substance?
27. For what substances does the standard enthalpy of combustion equal zero: CO, CO 2, H 2, O 2?
28. For the reaction 2Cl 2 (g) + 2H 2 O (l) 4HCl (g) + O 2 (g), calculate the standard enthalpy (kJ), if the standard enthalpies of formation of substances are known:
29. ∆Н = -1410.97 kJ/mol; ∆H = -2877.13 kJ/mol. How much heat will be released when 2 moles of ethylene and 4 moles of butane are burned together?
30. ∆Н = -1410.97 kJ/mol; ∆H = -2877.13 kJ/mol. What amount of heat will be released when burning 0.7 kg of a gas mixture consisting of 20% ethylene and 80% butane?
31. The standard enthalpy of the reaction MgCO 3 (s) → MgO (s) + CO 2 (g) is 101.6 kJ; standard enthalpies of formation of MgO (s) and CO 2 (g): -601.0 and -393.5 kJ/mol, respectively. What is the standard enthalpy of formation of magnesium carbonate MgCO 3?
32. What is called the thermodynamic probability of a system? What is entropy called? How is entropy expressed in terms of thermodynamic probability?
33. Formulate the second law of thermodynamics.
34. What is called the standard entropy of matter?
35. Formulate the third law of thermodynamics (Planck’s postulate).
36. What is called the entropy of a reaction? Standard entropy of a reaction?
37. Which expression is applicable to calculate ∆S p of the reaction CH 4 + CO 2 2CO + 2H 2:
∆S р = S + S + S + S ; ∆S р = S + S + 2S + 2S ;
∆S р = 2S + 2S – S + S ; ∆S р = 2S + 2S – S – S ?
38. For the reaction 2Cl 2 (u) + 2H 2 O (l) 4HCl (g) + O 2 (g), calculate the standard entropy (J/K), if the standard entropies of formation of substances are known:
39. What is called Gibbs free energy? What is its relationship with other thermodynamic functions?
40. How do we determine the direction of a reaction based on the sign of the Gibbs energy of a reaction?
41. At what temperatures is a reaction possible if ∆H<0, ∆S>0; ∆H<0, ∆S<0; ∆H>0, ∆S>0; ∆H>0, ∆S<0.
42. How is the equilibrium temperature of a process determined?
43. What is the Gibbs energy of the reaction ∆G р called? Standard Gibbs energy of a reaction?
44. Which expression is applicable to calculate ∆G p of the reaction 4NH 3 (g) + 5O 2 (g) 4NO (g) + 6H 2 O (l)
∆G р = ∆G 4 + ∆G 5 + ∆G 4 + ∆G 6 ; ∆G р = ∆G + ∆G + ∆G + ∆G ;
∆G р = 4∆G + 5∆G - 4∆G - 6∆G ; ∆G р = 4∆G + 6∆G - 4∆G - 5∆G ?
45. For the reaction HNO 3 (l) + HNO 2 (l) 2NO 2 (g) + H 2 O (l), calculate the standard Gibbs energy (kJ), if the standard Gibbs energies of formation of substances are known:
46. For the reaction Fe (s) + Al 2 O 3 (s) → Al (s) + Fe 2 O 3 (s), determine the equilibrium temperature and the possibility of the process occurring at 125 0 C, if ∆H = 853.8 kJ/ mole; ∆S = 37.68 J/mol·K.
47. What is meant by the rate of a chemical reaction?
48. Formulate the law of mass action.
49. In 40 s, as a result of two reactions Zn + 2HCl = ZnCl 2 + H 2 (1) and Zn + 2HBr = ZnBr 2 + H 2 (2), 8 g of zinc chloride and bromide were formed. Compare reaction rates.
50. If in the reaction 3Fe(NO 3) 2(solution) + 4HNO 3 = 3Fe(NO 3) 3(solution) + NO (g) + 2H 2 O (l) the concentration of Fe(NO 3) 2 increase by 7 times, and the concentration of HNO 3 by 4 times, how will the reaction rate change?
51. Make up the kinetic equation for the reaction Sb 2 S 3 (s) + 3H 2 (g) 2Sb (s) + 3H 2 S (g).
52. How is the rate of a multistage reaction determined?
53. How will the rate of the direct reaction CO (g) + 3H 2 (g) CH 4 (g) + H 2 O (g) change when the pressure of the system increases by 3 times?
54. What is the speed constant called? What does it depend on?
55. What is called activation energy? What does it depend on?
56. The rate constant of a certain reaction at a temperature of 310 K is equal to 4.6∙10 -5 l·mol -1 ·s -1 , and at a temperature of 330 K it is 6.8∙10 -5 l·mol -1 ·s -1 . What is the activation energy equal to?
57. The activation energy of a certain reaction is 250 kJ/mol. How will the rate constant change when the reaction temperature changes from 320 K to 340 K?
58. Write down the Arrhenius equation and van’t Hoff’s rule.
59. The activation energy of reaction (1) is 150 kJ/mol, the activation energy of reaction (2) is 176 kJ/mol. Compare the rate constants k 1 and k 2 .
60. How can we explain the increase in reaction rate with increasing temperature?
61. What is called the temperature coefficient of the reaction?
62. What is the temperature coefficient of the reaction if the rate constant of a certain reaction at 283 and 308 K is 1.77 and 7.56 l mol -1 s -1, respectively?
63. At a temperature of 350 K, the reaction ended in 3 s, and at a temperature of 330 K - in 28 s. How long will it take for it to end at a temperature of 310 K?
64. How does activation energy affect the temperature coefficient of a reaction?
65. What is called a catalyst? An inhibitor? Promoter? Catalytic poison?
66. What is called chemical equilibrium? How long does the system remain in equilibrium?
67. How are the rates of forward and reverse reactions related at the moment of equilibrium?
68. What is called the equilibrium constant? What does it depend on?
69. Express the equilibrium constant for the reactions 2NO + O 2 ↔ 2NO 2 ; Sb 2 S 3 (tv) + 3H 2 ↔ 2Sb (tv) + 3H 2 S (g).
70. At a certain temperature, the equilibrium constant of the reaction N 2 O 4 ↔ 2NO 2 is 0.16. In the initial state there was no NO 2, and the equilibrium concentration of NO 2 was 0.08 mol/l. What will be the equilibrium and initial concentrations of N 2 O 4?
71. Formulate Le Chatelier’s principle. How do changes in temperature, concentration, and total pressure affect equilibrium mixing?
72. Chemical dynamic equilibrium in the system was established at 1000 K and a pressure of 1 atm, when, as a result of the reaction Fe (sol) + CO 2 (g) ↔ FeO (sol) + CO (g), the partial pressure of carbon dioxide became equal to 0.54 atm. What is the equilibrium constant K p of this reaction?
73. Equilibrium concentrations (mol/l) of the components of the gas-phase system in which the reaction took place
3N 2 H 4 ↔ 4NH 3 + N 2, equal: =0.2; =0.4; =0.25. What is the equilibrium constant of the reversible
74. Equilibrium concentrations (mol/l) of the components of the gas-phase system in which the reaction occurs
N 2 + 3H 2 ↔ 2NH 3, equal: =0.12; =0.14; =0.1. Determine the initial concentrations of N 2 and H 2.
75. Equilibrium concentrations of the components of the gas phase of the system in which the reaction occurs
C (tv) + CO 2 ↔ 2CO at 1000 K and P total = 1 atm., equal to CO 2 - 17% vol. and CO - 83% vol. What is the constant?
reaction equilibrium?
76. The equilibrium constant Kc of the reversible gas-phase reaction CH 4 + H 2 O ↔ CO + 3H 2 at a certain temperature is 9.54 mol 2 l -2. The equilibrium concentrations of methane and water are 0.2 mol/L and 0.4 mol/L, respectively. Determine the equilibrium concentrations of CO and H 2.
77. Write down the relationship between the equilibrium constant K p and the Gibbs energy ∆G of a reversible reaction occurring under isothermal conditions.
78. Determine the equilibrium constant K p of the gas-phase reversible reaction COCl 2 ↔ CO + Cl 2 ; ∆H 0 = 109.78 kJ,
∆S 0 = 136.62 J/K at 900 K.
79. Equilibrium constant K p of the gas-phase reaction PCl 3 + Cl 2 ↔ PCl 5; ∆Н 0 = -87.87 kJ at 450 K is equal to 40.29 atm -1. Determine the Gibbs energy of this process (J/K).
80. Write down the relationship between K p and K c of the reversible gas-phase reaction 2CO + 2H 2 ↔ CH 4 + CO 2.
Related information.
Thermodynamics – the science of converting one form of energy into another based on the law of conservation of energy. Thermodynamics establishes the direction of the spontaneous flow of chemical reactions under given conditions. During chemical reactions, bonds in the starting substances are broken and new bonds are formed in the final products. The sum of binding energies after the reaction is not equal to the sum of binding energies before the reaction, i.e. the occurrence of a chemical reaction is accompanied by the release or absorption of energy, and its forms are different.
Thermochemistry is a branch of thermodynamics devoted to the study of the thermal effects of reactions. The thermal effect of a reaction measured at constant temperature and pressure is called enthalpy of reaction and are expressed in joules (J) and kilojoules (kJ).
For exothermic reactions, for endothermic reactions -. The enthalpy of formation of 1 mol of a given substance from simple substances, measured at a temperature of 298 K (25 ° C) and a pressure of 101.825 kPa (1 atm), is called standard (kJ/mol). The enthalpies of simple substances are conventionally assumed to be zero.
Thermochemical calculations are based on Hess’s law: t The thermal effect of a reaction depends only on the nature and physical state of the starting materials and final products, but does not depend on the transition path. Often in thermochemical calculations a corollary from Hess’s law is used: thermal effect of a chemical reaction equal to the sum of the heats of formation reaction products minus the sum of the heats of formation of the starting substances, taking into account the coefficients in front of the formulas of these substances in the reaction equation:
Thermochemical equations indicate the enthalpy of a chemical reaction. At the same time, the formula of each substance indicates its physical state: gaseous (g), liquid (l), solid crystalline (c).
In thermochemical equations, the thermal effects of reactions are given per 1 mole of the starting or final substance. Therefore, fractional odds are allowed here. In chemical reactions, the dialectical law of unity and struggle of opposites is manifested. On the one hand, the system strives for ordering (aggregation) - reducing N, and on the other hand – to disorder (disaggregation). The first trend increases with decreasing temperature, and the second - with increasing temperature. The tendency towards disorder is characterized by a quantity called entropy S[J/(mol. K)]. It is a measure of the disorder of the system. Entropy is proportional to the amount of matter and increases with increasing movement of particles during heating, evaporation, melting, gas expansion, weakening or breaking of bonds between atoms, etc. Processes associated with the orderliness of the system: condensation, crystallization, compression, strengthening of bonds, polymerization, etc. – lead to a decrease in entropy. Entropy is a function of state, i.e.
The general driving force of the process consists of two forces: the desire for order and the desire for disorder. With p = const and T = const, the overall driving force of the process can be represented as follows:
Gibbs energy, or isobaric-isothermal potential, also obeys a corollary of Hess's law:
Processes proceed spontaneously in the direction of decreasing any potential and, in particular, in the direction of decreasing . At equilibrium, the temperature at which the equilibrium reaction begins is equal to:
Table 5
Standard enthalpies of formation , entropy and Gibbs energy formation some substances at 298 K (25°C)
| Substance | , kJ/mol | , J/mol | , kJ/mol |
| CaO(k) | -635,5 | 39,7 | -604,2 |
| CaCO 3 (k) | -1207,0 | 88,7 | -1127,7 |
| Ca(OH) 2 (k) | -986,6 | 76,1 | -896,8 |
| H 2 O (l) | -285,8 | 70,1 | -237,3 |
| H2O (g) | -241,8 | 188,7 | -228,6 |
| Na 2 O (k) | -430,6 | 71,1 | -376,6 |
| NaOH (k) | -426,6 | 64,18 | -377,0 |
| H2S (g) | -21,0 | 205,7 | -33,8 |
| SO2 (g) | -296,9 | 248,1 | -300,2 |
| SO 3 (g) | -395,8 | 256,7 | -371,2 |
| C 6 H 12 O 6 (k) | -1273,0 | - | -919,5 |
| C 2 H 5 OH (l) | -277,6 | 160,7 | -174,8 |
| CO 2 (g) | -393,5 | 213,7 | -394,4 |
| CO(g) | -110,5 | 197,5 | -137,1 |
| C 2 H 4 (g) | 52,3 | 219,4 | 68,1 |
| CH 4 (g) | -74,9 | 186,2 | -50,8 |
| Fe 2 O 3 (k) | -822,2 | 87,4 | -740,3 |
| FeO(k) | -264,8 | 60,8 | -244,3 |
| Fe 3 O 4 (k) | -1117,1 | 146,2 | -1014,2 |
| CS 2 (g) | 115,3 | 65,1 | 237,8 |
| P 2 O 5 (k) | -1492 | 114,5 | -1348,8 |
| NH 4 Cl (k) | -315,39 | 94,56 | -343,64 |
| HCl (g) | -92,3 | 186,8 | -95,2 |
| NH 3 (g) | -46,2 | 192,6 | -16,7 |
| N2O (g) | 82,0 | 219,9 | 104,1 |
| NO (g) | 90,3 | 210,6 | 86,6 |
| NO 2 (g) | 33,5 | 240,2 | 51,5 |
| N2O4 (g) | 9,6 | 303,8 | 98,4 |
| CuO(k) | -162,0 | 42,6 | -129,9 |
| H2(g) | 130,5 | ||
| C (graphite) | 5,7 | ||
| O2 (g) | 205,0 | ||
| N 2 (g) | 181,5 | ||
| Fe(k) | 27,15 | ||
| Cl 2 (g) | 222,9 | ||
| KNO 3 (k) | -429,71 | 132,93 | -393,13 |
| KNO 2 (k) | -370,28 | 117,15 | -281,58 |
| K 2 O (k) | -361,5 | 87,0 | -193,3 |
| ZnO(k) | -350,6 | 43,6 | -320,7 |
| Al 2 O 3 (k) | -1676,0 | 50,9 | -1582,0 |
| PCl 5 (g) | -369,45 | 362,9 | -324,55 |
| PCl 3 (g) | -277,0 | 311,7 | -286,27 |
| H 2 O 2 (l) | -187,36 | 105,86 | -117,57 |
Speed reaction determined by the nature and concentration of the reacting substances and depends on temperature and catalyst.
Law of mass action: At constant temperature, the rate of a chemical reaction is proportional to the product of the concentration of the reactants to the power of their stoichiometric coefficients.
For the reaction aA + bB = cC + dD, the rate of the direct reaction is:
,
reverse reaction speed: 
, where are the concentrations of dissolved or gaseous compounds, mol/l;
a, b, c, d – stoichiometric coefficients in the equation;
K is the rate constant.
The expression for the reaction rate does not include solid concentrations.
The effect of temperature on the reaction rate is described by Van't Hoff's rule: for every 10 degrees heated, the reaction rate increases by 2-4 times.
Reaction rate at temperatures t 1 and t 2;
Temperature coefficient of reaction.
Most chemical reactions are reversible:
aA + bB cC + dD
the ratio of rate constants is a constant quantity called equilibrium constant
K p = const at T = const.
Le-Chatelier's principle: If a system in a state of chemical equilibrium is subject to any impact (change in temperature, pressure or concentration), the system will react in such a way as to reduce the applied impact:
a) when the temperature in equilibrium systems increases, the equilibrium shifts towards the endothermic reaction, and when the temperature decreases, towards the exothermic reaction;
b) when the pressure increases, the equilibrium shifts towards smaller volumes, and when the pressure decreases, towards larger volumes;
c) as the concentration increases, the equilibrium shifts towards its decrease.
Example 1. Determine the standard enthalpy change of the reaction:
Is this reaction exo- or endothermic?
Solution: The standard enthalpy change of a chemical reaction is equal to the sum of the standard enthalpies of formation of the reaction products minus the sum of the standard enthalpies of formation of the starting substances
With each summation, the number of moles of substances participating in the reaction should be taken into account in accordance with the reaction equation. The standard enthalpies of formation of simple substances are zero:
According to the tabular data:
Reactions that are accompanied by the release of heat are called exothermic, and those that are accompanied by the absorption of heat are called endothermic. At constant temperature and pressure, the change in enthalpy of a chemical reaction is equal in magnitude, but opposite in sign to its thermal effect. Since the standard enthalpy change for a given chemical reaction is , we conclude that this reaction is exothermic.
Example 2. The reduction reaction of Fe 2 O 3 with hydrogen proceeds according to the equation:
Fe 2 O 3 (K) + 3H 2 (G) = 2Fe (K) + 3H 2 O (G)
Is this reaction possible under standard conditions?
Solution: To answer this question of the problem, you need to calculate the standard change in the Gibbs energy of the reaction. Under standard conditions:
The summation is carried out taking into account the number of models involved in the reaction of substances; the formation of the most stable modification of a simple substance is taken equal to zero.
Taking into account the above
According to the tabular data:
Spontaneously occurring processes are decreasing. If< 0, процесс принципиально осуществим, если >0, the process cannot proceed spontaneously.
Therefore, this reaction is impossible under standard conditions.
Example 3. Write expressions for the law of mass action for reactions:
a) 2NO (G) + Cl 2 (G) = 2NOCl (G)
b) CaCO 3 (K) = CaO (K) + CO 2 (G)
Solution: According to the law of mass action, the reaction rate is directly proportional to the product of the concentrations of the reacting substances in powers equal to the stoichiometric coefficients:
a) V = k 2.
b) Since calcium carbonate is a solid whose concentration does not change during the reaction, the desired expression will look like:
V = k, i.e. V in this case the reaction rate at a certain temperature is constant.
Example 4. The endothermic reaction of phosphorus pentachloride decomposition proceeds according to the equation:
PCl 5(G) = PCl 3(G) + Cl 2(G); 
How to change: a) temperature; b) pressure; c) concentration to shift the equilibrium towards the direct reaction - decomposition of PCl 5? Write a mathematical expression for the rates of forward and reverse reactions, as well as the equilibrium constant.
Solution: A displacement or a shift in chemical equilibrium is a change in the equilibrium concentrations of reactants as a result of a change in one of the reaction conditions.
A shift in chemical equilibrium is subject to Le Chatelier's principle, according to which a change in one of the conditions under which a system is in equilibrium causes a shift in the equilibrium in the direction of the reaction that counteracts the resulting change.
a) Since the decomposition reaction of PCl 5 is endothermic, to shift the equilibrium towards the direct reaction, the temperature must be increased.
b) Since in this system the decomposition of PCl 5 leads to an increase in volume (two gaseous molecules are formed from one gas molecule), then to shift the equilibrium towards the direct reaction it is necessary to reduce the pressure.
c) A shift in equilibrium in the indicated direction can be achieved either by increasing the concentration of PCl 5 or by decreasing the concentration of PCl 3 or Cl 2 .
According to the law of mass action, the rates of direct (V 1) and reverse (V 2) reactions are expressed by the equations:
V 2 = k
The equilibrium constant of this reaction is expressed by the equation:
Test tasks:
81 - 100. a) calculate the standard change in enthalpy of the forward reaction and determine whether the reaction is exo- or endothermic;
b) determine the change in the Gibbs energy of the direct reaction and draw a conclusion about the possibility of its implementation under standard conditions;
c) write a mathematical expression for the rate of forward and reverse reactions, as well as the equilibrium constant;
d) how should the conditions be changed to shift the equilibrium of the process to the right?
81. CH 4 (g) + CO 2 (g) = 2CO (g) + 2H 2 (g)
82. FeO (K) + CO (g) = Fe (K) + CO 2 (g)
83. C 2 H 4 (g) + O 2 (g) = CO 2 (g) + H 2 O (g)
84. N 2(g) + 3H 2(g) = 2NH 3(g)
85. H 2 O (g) + CO (g) = CO 2 (g) + H 2 (g)
86. 4HCl (g) + O 2 (g) = 2H 2 O (g) + 2Cl 2 (g)
87. Fe 2 O 3 (K) + 3H 2 (g) = 2Fe (K) + 3H 2 O (g)
88. 2SO 2 (g) + O 2 (g) = 2SO 3 (g)
89. PCl 5(g) = PCl 3(g) + Cl 2(g)
90. CO 2 (g) + C (graphite) = 2CO (g)
91. 2H 2 S (g) + 3O 2 (g) = 2SO 2 (g) + H 2 O (g)
92. Fe 2 O 3 (K) + CO (g) = 2FeO (K) + CO 2 (g)
93. 4NH 3 (g) + 5O 2 (g) = 4NO (g) + 6H 2 O (g)
94. NH 4 Cl (K) = NH 3 (g) + HCl (g)
95. CH 4 (g) + 2O 2 (g) = CO 2 (g) + 2H 2 O (g)
96. CS 2(g) + 3O 2(g) = CO 2(g) + 2SO 2(g)
97. 4HCl (g) + O 2 (g) = 2Cl 2 (g) + 2H 2 O (g)
98. 2NO (g) + O 2 (g) = N 2 O 4 (g)
99. NH 3 (g) + HCl (g) = NH 4 Cl (K)
100. CS 2(g) + 3O 2(g) = 2Cl 2(g) + 2SO 2(g)
Topic 6: Solutions. Methods of expressing the concentration of solutions
Solutions are homogeneous systems consisting of a solvent, solutes and possible products of their interaction. The concentration of a solution is the content of dissolved substance in a certain mass or a known volume of solution or solvent.
Ways to express the concentration of solutions:
Mass fraction() shows the number of grams of solute in 100 g of solution:
Where T– mass of solute (g), T 1 – mass of solution (g).
Molar concentration shows the number of moles of solute contained in 1 liter of solution:
where M is the molar mass of the substance (g/mol), V is the volume of the solution (l).
Molal concentration shows the number of moles of solute contained in 1000 g of solvent: n 101-120.
| Find the mass fraction, molar concentration, molal concentration for the following solutions: | Option | Substance (x) | Mass of substance (x) | Water volume |
| Solution density | CuSO4 | 320 g | 1,019 | |
| 10 l | NaCl | 0.6 g | 1,071 | |
| 50 ml | H2SO4 | 2 g | 1,012 | |
| 100 ml | Na2SO4 | 2 g | 1,111 | |
| 13 g | HNO3 | 2 g | 1,066 | |
| 12.6 g | HCl | 3.6 kg | 1,098 | |
| 10 kg | NaOH | 8 g | 1,043 | |
| 200 g | MgCl2 | 190 g | 1,037 | |
| 810 g | KOH | 224 g | 1,206 | |
| 776 g | CuCl2 | 13.5 g | 1,012 | |
| 12.6 g | 800 ml | 8 g | 1,149 | |
| Solution density | NaOH | 10.8 g | 1,040 | |
| 10 l | 200 ml | 6.1 g | 1,005 | |
| Na2SO3 | 4.2 g | 500 ml | 1,082 | |
| 50 ml | 98 g | 1000 ml | 1,066 | |
| ZnCl2 | 13.6 g | 2 g | 1,052 | |
| H3PO4 | 9.8 g | 1000 ml | 1,012 | |
| Ba(OH)2 | 100 g | 900 g | 1,085 | |
| H3PO4 | 29.4 g | 6.1 g | 1,023 | |
| 10 kg | 28 g | 72 g | 1,309 |
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