How to calculate the thermal effect of a reaction. Calculations according to thermochemical equations

Task #6

Calculate the average heat capacity of the substance given in table. 6, in the temperature range from 298 to T TO.

Table 6

	Substance			Substance

Solution:

Consider the calculation of the average heat capacity of ammonia in the temperature range from 298 to 800 TO.

Heat capacity- this is the ratio of the amount of heat absorbed by the body during heating to the temperature increase that accompanies heating. For an individual substance, there are specific(one kilogram) and molar(one mole) heat capacity.

True heat capacity

, (21)

where δ Q is the infinitesimal amount of heat required to raise the temperature of a body by an infinitesimal amount dT .

Average heat capacity is the ratio of heat Q to temperature rise ∆ T = T 2 – T 1 ,

.

Since heat is not a state function and depends on the process path, it is necessary to specify the conditions for the heating process to occur. In isochoric and isobaric processes for an infinitesimal change δ Q V = dU and δ Q p = dH, that's why

and
. (22)

Connection between true isochoric(FROM V) and isobaric (C p) heat capacities substances and average isochoric
and isobaric
heat capacities in the temperature range from T 1 before T 2 is expressed by equations (23) and (24):

; (23)

. (24)

The dependences of the true heat capacity on temperature are expressed by the following empirical equations:

; (for inorganic substances) (25)

. (for organic substances) (26)

Let's use the reference book of physical and chemical quantities. Let us write the coefficients (a, b, c) of the equation for the dependence of the isobaric heat capacity of ammonia on temperature:

Table 7

Substance
		b·ten 3	c / ·ten –5

We write the equation for the dependence of the true heat capacity of ammonia on temperature:

.

We substitute this equation into formula (24) and calculate the average heat capacity of ammonia:

= 1/(800-298)
=

0.002 = 43.5 J/mol K.

Task #7

For the chemical reaction given in table. 2, plot the sum of the heat capacities of the reaction products as a function of temperature
and the sum of the heat capacities of the initial substances on temperature
. Dependency equations
take it from the handbook. Calculate the change in heat capacity during a chemical reaction (
) at temperatures of 298 K, 400 K and T K (Table 6).

Solution:

Let us calculate the change in heat capacity at temperatures of 298 K, 400 K and 600 K using the ammonia synthesis reaction as an example:

Let us write out the coefficients (a, b, c, c /) 1 of the equations for the dependence of the true heat capacity of ammonia on temperature for the initial substances and reaction products, taking into account the stoichiometric coefficients . Let's calculate the sum of the coefficients. For example, the sum of the coefficients a for starting materials is equal to

= 27.88 + 3 27.28 = 109.72.

The sum of the coefficients a for the reaction products is

= 2 29.8 = 59.6.

=
=59,6 – 109,72 = –50,12.

Table 8

Substance			b·ten 3	c / ·ten – 5	s 10 6
initial substances
( , , )
( , , )
, ,

Thus, the dependence equation

for reaction products has the following form:

\u003d 59.60 + 50.96 10 -3 T - 3.34 10 5 / T 2.

To plot the dependence of the sum of the heat capacity of the reaction products on temperature
calculate the sum of heat capacities at several temperatures:

At T = 298 K

\u003d 59.60 + 50.96 10 -3 298 - 3.34 10 5 / 298 2 \u003d 71.03 J / K;

At T = 400 K
= 77.89 J/K;

At T = 600 K
= 89.25 J/K.

Dependency equation
for starting materials has the form:

\u003d 109.72 + 14.05 10 -3 T + 1.50 10 -5 / T 2.

Similarly, we calculate
starting materials at several temperatures:

At T=298 K

\u003d 109.72 + 14.05 10 -3 298 + 1.50 10 5 / 298 2 \u003d 115.60 J / K;

At T = 400 K
= 116.28 J/K;

At T = 600 K
= 118.57 J/K.

Next, we calculate the change in isobaric heat capacity
during the reaction at several temperatures:

\u003d -50.12 + 36.91 10 -3 T - 4.84 10 5 / T 2,

= -44.57 J/K;

= -38.39 J/K;

= -29.32 J/K.

Based on the calculated values, we build graphs of the dependences of the sum of the heat capacities of the reaction products and the sum of the heat capacities of the starting substances on temperature.

Figure 2. Dependences of the total heat capacities of the initial substances and reaction products on temperature for the ammonia synthesis reaction

In this temperature range, the total heat capacity of the starting materials is higher than the total heat capacity of the products, therefore,
over the entire temperature range from 298 K to 600 K.

Task #8

Calculate the thermal effect of the reaction given in table. 2, at temperature T K (Table 6).

Solution:

Let us calculate the thermal effect of the ammonia synthesis reaction at a temperature of 800 TO.

Dependence of thermal effect
temperature response describes Kirchhoff's law

, (27)

where
- change in the heat capacity of the system during the reaction. Let's analyze the equation:

1) If
> 0, i.e. the sum of the heat capacities of the reaction products is greater than the sum of the heat capacities of the starting materials, then > 0,. addiction
increasing, and with increasing temperature, the thermal effect increases.

2) If
< 0, то< 0, т.е. зависимость убывающая, и с повышением температуры тепловой эффект уменьшается.

3) If
= 0, then = 0, the thermal effect does not depend on temperature.

In integral form, the Kirchhoff equation has the following form:

. (28)

a) if the heat capacity does not change during the process, i.e. the sum of the heat capacities of the reaction products is equal to the sum of the heat capacities of the starting materials (
), then the thermal effect does not depend on temperature

= const.

b) for approximate calculation we can neglect the dependence of heat capacities on temperature and use the values ​​of the average heat capacities of the reaction participants (
). In this case, the calculation is made according to the formula

c) for accurate calculation data are needed on the dependence of the heat capacity of all participants in the reaction on temperature
. In this case, the thermal effect is calculated by the formula

(30)

We write out the reference data (Table 9) and calculate the changes in the corresponding values ​​for each column by analogy with task No. 7). We use the obtained data to calculate:

Approximately:

\u003d -91880 + (-31.88) (800 - 298) \u003d -107883.8 J \u003d - 107.88 kJ.

\u003d -91880 + (-50.12) (800 - 298) + 1/2 36.91 10 -3 (800 2 - 298 2) +

- (-4.84 10 5) (1/800 - 1/298) \u003d - 107815 J \u003d - 107.82 kJ.

For the ammonia synthesis reaction, the change in heat capacity during the reaction
< 0 (см. задачу №7). Следовательно< 0, с повышением температуры тепловой эффект уменьшается.

Table 9

	Substance			Sum for reaction products	Amount for starting substances	Change in the course of a reaction
,				=	=	=
, J/(mol K)				=	=	=
				=	=	=
				=	=	=
				=	= 1,5	=
				= 0	= 0	= 0

As a result of studying this topic, you will learn: How do the usual equations of chemical reactions differ from their thermochemical equations. What factors affect the rate of chemical reactions? How does true (chemical) equilibrium differ from apparent equilibrium. In what direction does the equilibrium shift when external conditions change. What is the mechanism of homogeneous and heterogeneous catalysis. What are inhibitors and promoters. As a result of studying this topic, you will learn: Calculate the thermal effects of chemical reactions using the enthalpies of formation of substances. Carry out calculations using the mathematical expression of the van't Hoff principle. Determine the direction of shift in chemical equilibrium with changes in temperature and pressure. Study questions:

6.1. Energy of chemical processes

6.1.1. Internal energy and enthalpy

In any process, the law of conservation of energy is observed:

Q = ∆U + A.

This equality means that if heat Q is supplied to the system, then it is spent on changing the internal energy Δ U and on doing work A.

Internal energy system is its total reserve, including the energy of the translational and rotational motion of molecules, the energy of the motion of electrons in atoms, the energy of interaction of nuclei with electrons, nuclei with nuclei, etc., i.e. all types of energy, except for the kinetic and potential energy of the system as a whole.

The work performed by the system during the transition from state 1, characterized by volume V 1, to state 2 (volume V 2) at constant pressure (expansion work), is equal to:

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

At constant pressure (р=const), taking into account the expression for the expansion work, the energy conservation law will be written as follows:

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

The sum of the internal energy of a system and the product of its volume and pressure is called enthalpy H:

Since the exact value of the internal energy of the system is unknown, the absolute values ​​of the enthalpies cannot be obtained either. Changes in enthalpies Δ H are of scientific importance and practical application.

Internal energy U and enthalpy H are state functions systems. State functions are such characteristics of the system, changes in which are determined only by the final and initial state of the system, i.e. are independent of the process path.

6.1.2. Exo- and endothermic processes

The flow of chemical reactions is accompanied by the absorption or release of heat. exothermic called a reaction that proceeds with the release of heat into the environment, and endothermic- with the absorption of heat from the environment.

Many processes in industry and in laboratory practice proceed at constant pressure and temperature (T=const, p=const). The energy characteristic of these processes is the change in enthalpy:

Q P \u003d -Δ N.

For processes occurring at constant volume and temperature (T=const, V=const) Q V =-Δ U.

For exothermic reactions Δ H< 0, а в случае протекания эндотермической реакции Δ Н >0. For example,

N 2 (g) + SO 2 (g) \u003d N 2 O (g); ΔН 298 = +82kJ,

CH 4 (g) + 2O 2 (g) \u003d CO 2 (g) + 2H 2 O (g); ΔN 298 = -802kJ.

Chemical equations in which the thermal effect of the reaction is additionally indicated (the value of the DH process), as well as the state of aggregation of substances and temperature, are called thermochemical equations.

In thermochemical equations, the phase state and allotropic modifications of reagents and formed substances are noted: d - gaseous, g - liquid, k - crystalline; S (rhombus), S (monocle), C (graphite), C (diamond), etc.

6.1.3. Thermochemistry; Hess' law

Energy phenomena accompanying physical and chemical processes studies thermochemistry. The basic law of thermochemistry is the law formulated by the Russian scientist G.I. Hess in 1840.

Hess' law: the change in the enthalpy of the process depends on the type and state of the starting materials and reaction products, but does not depend on the path of the process.

When considering thermochemical effects, the expression “process enthalpy” is often used instead of the concept of “change in the enthalpy of the process”, meaning by this concept the value of Δ H. It is incorrect to use the concept of “heat effect of the process” when formulating the Hess law, since the value of Q in the general case is not a function of the state . As mentioned above, only at a constant pressure Q P =-Δ N (at a constant volume Q V =-Δ U).

So, the formation of PCl 5 can be considered as the result of the interaction of simple substances:

P (c, white) + 5/2Cl 2 (g) = PCl 5 (c); Δ H 1,

or as a result of a process that takes place in several stages:

P (k, white) + 3/2Cl 2 (g) = PCl 3 (g); Δ H 2,

PCl 3 (g) + Cl 2 (g) \u003d PCl 5 (c); Δ H 3,

or in total:

P (c, white) + 5/2Cl 2 (g) = PCl 5 (c); Δ H 1 \u003d Δ H 2 + Δ H 3.

6.1.4. Enthalpies of formation of substances

The enthalpy of formation is the enthalpy of the process of formation of a substance in a given state of aggregation from simple substances that are in stable modifications. The enthalpy of formation of sodium sulfate, for example, is the enthalpy of reaction:

2Na (c) + S (rhombus) + 2O 2 (g) \u003d Na 2 SO 4 (c).

The enthalpy of formation of simple substances is zero.

Since the thermal effect of a reaction depends on the state of substances, temperature, and pressure, it was agreed to use in thermochemical calculations standard enthalpies of formation are the enthalpies of formation of substances that are at a given temperature in standard condition. As a standard state for substances in a condensed state, the real state of the substance at a given temperature and pressure of 101.325 kPa (1 atm) is taken. Reference books usually give the standard enthalpies of formation of substances at a temperature of 25 o C (298K), referred to 1 mol of a substance (Δ H f o 298). Standard enthalpies of formation of some substances at T=298K are given in Table. 6.1.

Table 6.1.

Standard enthalpies of formation (Δ H f o 298) of some substances

Substance	Δ H f o 298, kJ/mol	Substance	Δ H f o 298, kJ/mol

The standard enthalpies of formation for most complex substances are negative values. For a small number of unstable substances, Δ H f o 298 > 0. Such substances, in particular, include nitric oxide (II) and nitric oxide (IV), Table 6.1.

6.1.5. Calculation of thermal effects of chemical reactions

To calculate the enthalpies of processes, a consequence of the Hess law is used: the enthalpy of reaction is equal to the sum of the enthalpies of formation of the reaction products minus the sum of the enthalpies of formation of the starting substances, taking into account the stoichiometric coefficients .

Calculate the enthalpy of decomposition of calcium carbonate. The process is described by the following equation:

CaCO 3 (c) \u003d CaO (c) + CO 2 (g).

The enthalpy of this reaction will be equal to the sum of the enthalpies of formation of calcium oxide and carbon dioxide minus the enthalpy of formation of calcium carbonate:

Δ H o 298 \u003d Δ H f o 298 (CaO (c)) + Δ H f o 298 (CO 2 (g)) - Δ H f o 298 (CaCO 3 (c)).

Using the data in Table 6.1. we get:

Δ H o 298 = - 635.1 -393.5 + 1206.8 = + 178.2 kJ.

It follows from the obtained data that the considered reaction is endothermic, i.e. proceeds with the absorption of heat.

CaO (c) + CO 2 (c) \u003d CaCO 3 (c)

Accompanied by the release of heat. Its enthalpy will be equal to

Δ H o 298 = -1206.8 + 635.1 + 393.5 = -178.2 kJ.

6.2. The rate of chemical reactions

6.2.1. The concept of reaction rate

The branch of chemistry that deals with the rate and mechanisms of chemical reactions is called chemical kinetics. One of the key concepts in chemical kinetics is the rate of a chemical reaction.

The rate of a chemical reaction is determined by the change in the concentration of the reacting substances per unit time at a constant volume of the system.

Consider the following process:

Let at some point in time t 1 the concentration of substance A be equal to the value c 1, and at the moment t 2 - the value c 2 . For a period of time from t 1 to t 2, the change in concentration will be Δ c \u003d c 2 - c 1. The average reaction rate is:

The minus sign is put because as the reaction proceeds (Δ t> 0), the concentration of the substance decreases (Δ c< 0), в то время, как скорость реакции является положительной величиной.

The rate of a chemical reaction depends on the nature of the reactants and on the reaction conditions: concentration, temperature, presence of a catalyst, pressure (for gas reactions) and some other factors. In particular, with an increase in the contact area of ​​substances, the reaction rate increases. The reaction rate also increases with an increase in the stirring rate of the reactants.

The numerical value of the reaction rate also depends on which component is used to calculate the reaction rate. For example, the speed of the process

H 2 + I 2 \u003d 2HI,

calculated from the change in the concentration of HI is twice the reaction rate calculated from the change in the concentration of the reagents H 2 or I 2 .

6.2.2. Dependence of reaction rate on concentration; order and molecularity of the reaction

The basic law of chemical kinetics is law of mass action- establishes the dependence of the reaction rate on the concentration of the reactants.

The reaction rate is proportional to the product of the concentrations of the reactants. For a reaction written in general form as

aA + bB = cC + dD,

the dependence of the reaction rate on concentration has the form:

v = k [A] α [B] β .

In this kinetic equation, k is the proportionality factor, called rate constant; [A] and [B] are the concentrations of substances A and B. The reaction rate constant k depends on the nature of the reacting substances and on the temperature, but does not depend on their concentrations. Coefficients α and β are found from experimental data.

The sum of the exponents in the kinetic equations is called the total in order reactions. There is also a particular order of the reaction in one of the components. For example, for the reaction

H 2 + C1 2 \u003d 2 HC1

The kinetic equation looks like this:

v = k 1/2,

those. the overall order is 1.5 and the reaction orders for the H 2 and C1 2 components are 1 and 0.5, respectively.

Molecularity reaction is determined by the number of particles, the simultaneous collision of which is the elementary act of chemical interaction. Elementary act (elementary stage)- a single act of interaction or transformation of particles (molecules, ions, radicals) into other particles. For elementary reactions, the molecularity and order of the reaction are the same. If the process is multi-stage and therefore the reaction equation does not reveal the mechanism of the process, the order of the reaction does not coincide with its molecularity.

Chemical reactions are divided into simple (single-stage) and complex, occurring in several stages.

Monomolecular reaction is a reaction in which the elementary act is a chemical transformation of one molecule. For example:

CH 3 CHO (g) \u003d CH 4 (g) + CO (g).

Bimolecular reaction- a reaction in which the elementary act is carried out when two particles collide. For example:

H 2 (g) + I 2 (g) \u003d 2 HI (g).

trimolecular reaction- a simple reaction, the elementary act of which is carried out with the simultaneous collision of three molecules. For example:

2NO (g) + O 2 (g) \u003d 2 NO 2 (g).

It has been established that the simultaneous collision of more than three molecules, leading to the formation of reaction products, is practically impossible.

The law of mass action does not apply to reactions involving solids, since their concentrations are constant and they react only on the surface. The rate of such reactions depends on the size of the contact surface between the reactants.

6.2.3. Temperature dependence of the reaction rate

The rate of chemical reactions increases with increasing temperature. This increase is caused by an increase in the kinetic energy of the molecules. In 1884, the Dutch chemist van't Hoff formulated the rule: for every 10 degrees increase in temperature, the rate of chemical reactions increases by 2-4 times.

Van't Hoff's rule is written as:

,

where V t 1 and V t 2 are the reaction rates at temperatures t 1 and t 2 ; γ - temperature coefficient of speed, equal to 2 - 4.

The van't Hoff rule is used to approximate the effect of temperature on the reaction rate. A more accurate equation describing the dependence of the reaction rate constant on temperature was proposed in 1889 by the Swedish scientist S. Arrhenius:

.

In the Arrhenius equation, A is a constant, E is the activation energy (J/mol); T is temperature, K.

According to Arrhenius, not all collisions of molecules lead to chemical transformations. Only molecules with some excess energy are able to react. This excess energy that colliding particles must have in order for a reaction to occur between them is called activation energy.

6.3. The concept of catalysis and catalysts

A catalyst is a substance that changes the rate of a chemical reaction but remains chemically unchanged at the end of the reaction.

Some catalysts speed up the reaction, while others, called inhibitors, slow it down. For example, adding a small amount of MnO 2 as a catalyst to hydrogen peroxide H2O2 causes rapid decomposition:

2 H 2 O 2 - (MnO 2) 2 H 2 O + O 2.

In the presence of small amounts of sulfuric acid, a decrease in the rate of decomposition of H 2 O 2 is observed. In this reaction, sulfuric acid acts as an inhibitor.

Depending on whether the catalyst is in the same phase as the reactants or forms an independent phase, there are homogeneous and heterogeneous catalysis.

homogeneous catalysis

In the case of homogeneous catalysis, the reactants and the catalyst are in the same phase, for example, gaseous. The mechanism of action of the catalyst is based on the fact that it interacts with reactants to form intermediate compounds.

Consider the mechanism of action of the catalyst. In the absence of a catalyst, the reaction

It flows very slowly. The catalyst forms with the starting materials (for example, with substance B) a reactive intermediate product:

which reacts vigorously with another starting material to form the final reaction product:

VK + A \u003d AB + K.

Homogeneous catalysis takes place, for example, in the process of oxidation of sulfur(IV) oxide to sulfur(VI) oxide, which occurs in the presence of nitrogen oxides.

homogeneous reaction

2 SO 2 + O 2 \u003d 2 SO 3

in the absence of a catalyst is very slow. But when a catalyst (NO) is introduced, an intermediate compound (NO2) is formed:

O 2 + 2 NO \u003d 2 NO 2,

which easily oxidizes SO 2:

NO 2 + SO 2 \u003d SO 3 + NO.

The activation energy of the latter process is very low, so the reaction proceeds at a high rate. Thus, the action of catalysts is reduced to a decrease in the activation energy of the reaction.

heterogeneous catalysis

In heterogeneous catalysis, the catalyst and reactants are in different phases. The catalyst is usually in the solid state and the reactants are in the liquid or gaseous state. In heterogeneous catalysis, the acceleration of the process is usually associated with the catalytic effect of the catalyst surface.

Catalysts differ in selectivity (selectivity) of action. So, for example, in the presence of an aluminum oxide catalyst Al 2 O 3 at 300 o C, water and ethylene are obtained from ethyl alcohol:

C 2 H 5 OH - (Al 2 O 3) C 2 H 4 + H 2 O.

At the same temperature, but in the presence of copper Cu as a catalyst, ethyl alcohol is dehydrogenated:

C 2 H 5 OH - (Cu) CH 3 CHO + H 2.

Small amounts of certain substances reduce or even completely destroy the activity of catalysts (catalyst poisoning). Such substances are called catalytic poisons. For example, oxygen causes reversible poisoning of the iron catalyst in the synthesis of NH 3 . The activity of the catalyst can be restored by passing a fresh mixture of nitrogen and hydrogen purified from oxygen. Sulfur causes irreversible poisoning of the catalyst in the synthesis of NH 3 . Its activity can no longer be restored by passing a fresh mixture of N 2 +H 2 .

Substances that enhance the action of catalysts are called promoters, or activators(promotion of platinum catalysts, for example, is carried out by adding iron or aluminum).

The mechanism of heterogeneous catalysis is more complex. To explain it, the adsorption theory of catalysis is used. The surface of the catalyst is heterogeneous, so it has the so-called active centers. Reacting substances are adsorbed on active sites. The latter process causes the approach of the reacting molecules and an increase in their chemical activity, since the bond between the atoms of the adsorbed molecules is weakened, the distance between the atoms increases.

On the other hand, it is believed that the accelerating effect of a catalyst in heterogeneous catalysis is due to the fact that the reactants form intermediate compounds (as in the case of homogeneous catalysis), which leads to a decrease in the activation energy.

6.4. Chemical equilibrium

Irreversible and reversible reactions

Reactions that proceed in only one direction and end with the complete transformation of the starting substances into final substances are called irreversible.

Irreversible, i.e. proceeding to the end are reactions in which

Chemical reactions that can go in opposite directions are called reversible. Typical reversible reactions are the reactions of ammonia synthesis and oxidation of sulfur(IV) oxide to sulfur(VI) oxide:

N 2 + 3 H 2 2 NH 3,

2 SO 2 + O 2 2 SO 3 .

When writing the equations of reversible reactions, instead of the equal sign, put two arrows pointing in opposite directions.

In reversible reactions, the rate of the direct reaction at the initial moment of time has a maximum value, which decreases as the concentration of the initial reagents decreases. On the contrary, the reverse reaction initially has a minimum rate, which increases as the concentration of the products increases. As a result, there comes a moment when the rates of the forward and reverse reactions become equal and chemical equilibrium is established in the system.

Chemical equilibrium

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

Chemical equilibrium is also called true equilibrium. In addition to the equality of the rates of forward and reverse reactions, true (chemical) equilibrium is characterized by the following features:

the immutability of the state of the system is caused by the flow of direct and reverse reactions, that is, the equilibrium state is dynamic;

the state of the system remains unchanged in time if there is no external influence on the system;

any external influence causes a shift in the equilibrium of the system; however, if the external influence is removed, then the system returns to its original state again;

• the state of the system is the same regardless of which side the system approaches equilibrium from - from the side of the starting substances or from the side of the reaction products.

must be distinguished from the real apparent equilibrium. So, for example, a mixture of oxygen and hydrogen in a closed vessel at room temperature can be stored for an arbitrarily long time. However, the initiation of the reaction (electric discharge, ultraviolet irradiation, temperature increase) causes the reaction of water formation to proceed irreversibly.

6.5. Le Chatelier's principle

The influence of changes in external conditions on the equilibrium position is determined by Le Chatel principle e (France, 1884): if a system in equilibrium is subjected to some external action, then the equilibrium in the system will shift in the direction of weakening this effect.

Le Chatelier's principle applies not only to chemical processes, but also to physical ones, such as boiling, crystallization, dissolution, etc.

Consider the influence of various factors on the chemical equilibrium using the ammonia synthesis reaction as an example:

N 2 + 3 H 2 2 NH 3; ΔH = -91.8 kJ.

Effect of concentration on chemical equilibrium.

In accordance with Le Chatelier's principle, an increase in the concentration of the initial substances shifts the equilibrium towards the formation of reaction products. An increase in the concentration of the reaction products shifts the equilibrium towards the formation of the starting substances.

In the process of ammonia synthesis considered above, the introduction of additional amounts of N 2 or H 2 into the equilibrium system causes a shift in the equilibrium in the direction in which the concentration of these substances decreases, therefore, the equilibrium shifts towards the formation of NH3. Increasing the concentration of ammonia shifts the equilibrium towards the starting materials.

A catalyst speeds up both the forward and reverse reactions equally, so the introduction of a catalyst does not affect the chemical equilibrium.

Effect of Temperature on Chemical Equilibrium

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

The degree of equilibrium shift is determined by the absolute value of the thermal effect: the greater the value of ΔH of the reaction, the greater the effect of temperature.

In the ammonia synthesis reaction under consideration, an increase in temperature will shift the equilibrium towards the starting materials.

Effect of pressure on chemical equilibrium

A change in pressure affects the chemical equilibrium with the participation of gaseous substances. According to Le Chatelier's principle, an increase in pressure shifts the equilibrium in the direction of a reaction proceeding with a decrease in the volume of gaseous substances, and a decrease in pressure shifts the equilibrium in the opposite direction. The ammonia synthesis reaction proceeds with a decrease in the volume of the system (there are four volumes on the left side of the equation, and two volumes on the right). Therefore, an increase in pressure shifts the equilibrium towards the formation of ammonia. A decrease in pressure will shift the equilibrium in the opposite direction. If in the reversible reaction equation the number of molecules of gaseous substances in the right and left parts are equal (the reaction proceeds without changing the volume of gaseous substances), then pressure does not affect the equilibrium position in this system.

Exercise 81.
Calculate the amount of heat that will be released during the reduction of Fe 2O3 metallic aluminum if 335.1 g of iron was obtained. Answer: 2543.1 kJ.
Solution:
Reaction equation:

\u003d (Al 2 O 3) - (Fe 2 O 3) \u003d -1669.8 - (-822.1) \u003d -847.7 kJ

Calculation of the amount of heat that is released upon receipt of 335.1 g of iron, we produce from the proportion:

(2 . 55,85) : -847,7 = 335,1 : X; x = (0847.7 . 335,1)/ (2 . 55.85) = 2543.1 kJ,

where 55.85 is the atomic mass of iron.

Answer: 2543.1 kJ.

Thermal effect of the reaction

Task 82.
Gaseous ethyl alcohol C2H5OH can be obtained by the interaction of ethylene C 2 H 4 (g) and water vapor. Write the thermochemical equation for this reaction, having previously calculated its thermal effect. Answer: -45.76 kJ.
Solution:
The reaction equation is:

C 2 H 4 (g) + H 2 O (g) \u003d C2H 5 OH (g); = ?

The values ​​of the standard heats of formation of substances are given in special tables. Considering that the heats of formation of simple substances are conditionally taken equal to zero. Calculate the thermal effect of the reaction, using the consequence of the Hess law, we get:

\u003d (C 2 H 5 OH) - [ (C 2 H 4) + (H 2 O)] \u003d
= -235.1 -[(52.28) + (-241.83)] = - 45.76 kJ

Reaction equations in which their states of aggregation or crystalline modification are indicated near the symbols of chemical compounds, as well as the numerical value of thermal effects, are called thermochemical. In thermochemical equations, unless otherwise specified, the values ​​of thermal effects at a constant pressure Q p are indicated equal to the change in the enthalpy of the system. The value is usually given on the right side of the equation, separated by a comma or semicolon. The following abbreviations for the aggregate state of matter are accepted: G- gaseous, and- liquid, to

If heat is released as a result of a reaction, then< О. Учитывая сказанное, составляем термохимическое уравнение данной в примере реакции:

C 2 H 4 (g) + H 2 O (g) \u003d C 2 H 5 OH (g); = - 45.76 kJ.

Answer:- 45.76 kJ.

Task 83.
Calculate the thermal effect of the reduction reaction of iron (II) oxide with hydrogen, based on the following thermochemical equations:

a) EEO (c) + CO (g) \u003d Fe (c) + CO 2 (g); = -13.18 kJ;
b) CO (g) + 1/2O 2 (g) = CO 2 (g); = -283.0 kJ;
c) H 2 (g) + 1/2O 2 (g) = H 2 O (g); = -241.83 kJ.
Answer: +27.99 kJ.

Solution:
The reaction equation for the reduction of iron oxide (II) with hydrogen has the form:

EeO (k) + H 2 (g) \u003d Fe (k) + H 2 O (g); = ?

\u003d (H2O) - [ (FeO)

The heat of formation of water is given by the equation

H 2 (g) + 1/2O 2 (g) = H 2 O (g); = -241.83 kJ,

and the heat of formation of iron oxide (II) can be calculated if equation (a) is subtracted from equation (b).

\u003d (c) - (b) - (a) \u003d -241.83 - [-283.o - (-13.18)] \u003d + 27.99 kJ.

Answer:+27.99 kJ.

Task 84.
During the interaction of gaseous hydrogen sulfide and carbon dioxide, water vapor and carbon disulfide СS 2 (g) are formed. Write the thermochemical equation for this reaction, preliminarily calculate its thermal effect. Answer: +65.43 kJ.
Solution:
G- gaseous, and- liquid, to- crystalline. These symbols are omitted if the aggregate state of substances is obvious, for example, O 2, H 2, etc.
The reaction equation is:

2H 2 S (g) + CO 2 (g) \u003d 2H 2 O (g) + CS 2 (g); = ?

The values ​​of the standard heats of formation of substances are given in special tables. Considering that the heats of formation of simple substances are conditionally taken equal to zero. The thermal effect of the reaction can be calculated using the corollary e from the Hess law:

\u003d (H 2 O) + (CS 2) - [(H 2 S) + (CO 2)];
= 2(-241.83) + 115.28 – = +65.43 kJ.

2H 2 S (g) + CO 2 (g) \u003d 2H 2 O (g) + CS 2 (g); = +65.43 kJ.

Answer:+65.43 kJ.

Thermochemical reaction equation

Task 85.
Write the thermochemical equation for the reaction between CO (g) and hydrogen, as a result of which CH 4 (g) and H 2 O (g) are formed. How much heat will be released during this reaction if 67.2 liters of methane were obtained in terms of normal conditions? Answer: 618.48 kJ.
Solution:
Reaction equations in which their states of aggregation or crystalline modification are indicated near the symbols of chemical compounds, as well as the numerical value of thermal effects, are called thermochemical. In thermochemical equations, unless it is specifically stated, the values ​​of thermal effects at constant pressure Q p are indicated equal to the change in the enthalpy of the system. The value is usually given on the right side of the equation, separated by a comma or semicolon. The following abbreviations for the aggregate state of matter are accepted: G- gaseous, and- something to- crystalline. These symbols are omitted if the aggregate state of substances is obvious, for example, O 2, H 2, etc.
The reaction equation is:

CO (g) + 3H 2 (g) \u003d CH 4 (g) + H 2 O (g); = ?

The values ​​of the standard heats of formation of substances are given in special tables. Considering that the heats of formation of simple substances are conditionally taken equal to zero. The thermal effect of the reaction can be calculated using the corollary e from the Hess law:

\u003d (H 2 O) + (CH 4) - (CO)];
\u003d (-241.83) + (-74.84) ​​- (-110.52) \u003d -206.16 kJ.

The thermochemical equation will look like:

22,4 : -206,16 = 67,2 : X; x \u003d 67.2 (-206.16) / 22? 4 \u003d -618.48 kJ; Q = 618.48 kJ.

Answer: 618.48 kJ.

Heat of Formation

Task 86.
The thermal effect of which reaction is equal to the heat of formation. Calculate the heat of formation of NO from the following thermochemical equations:
a) 4NH 3 (g) + 5O 2 (g) \u003d 4NO (g) + 6H 2 O (g); = -1168.80 kJ;
b) 4NH 3 (g) + 3O 2 (g) \u003d 2N 2 (g) + 6H 2 O (g); = -1530.28 kJ
Answer: 90.37 kJ.
Solution:
The standard heat of formation is equal to the heat of formation of 1 mol of this substance from simple substances under standard conditions (T = 298 K; p = 1.0325.105 Pa). The formation of NO from simple substances can be represented as follows:

1/2N 2 + 1/2O 2 = NO

Given the reaction (a) in which 4 moles of NO are formed and the reaction (b) is given in which 2 moles of N2 are formed. Both reactions involve oxygen. Therefore, to determine the standard heat of formation of NO, we compose the following Hess cycle, i.e., we need to subtract equation (a) from equation (b):

Thus, 1/2N 2 + 1/2O 2 = NO; = +90.37 kJ.

Answer: 618.48 kJ.

Task 87.
Crystalline ammonium chloride is formed by the interaction of gaseous ammonia and hydrogen chloride. Write the thermochemical equation for this reaction, having previously calculated its thermal effect. How much heat will be released if 10 liters of ammonia were consumed in the reaction in terms of normal conditions? Answer: 78.97 kJ.
Solution:
Reaction equations in which their states of aggregation or crystalline modification are indicated near the symbols of chemical compounds, as well as the numerical value of thermal effects, are called thermochemical. In thermochemical equations, unless it is specifically stated, the values ​​of thermal effects at constant pressure Q p are indicated equal to the change in the enthalpy of the system. The value is usually given on the right side of the equation, separated by a comma or semicolon. The following are accepted to- crystalline. These symbols are omitted if the aggregate state of substances is obvious, for example, O 2, H 2, etc.
The reaction equation is:

NH 3 (g) + HCl (g) \u003d NH 4 Cl (k). ; = ?

The values ​​of the standard heats of formation of substances are given in special tables. Considering that the heats of formation of simple substances are conditionally taken equal to zero. The thermal effect of the reaction can be calculated using the corollary e from the Hess law:

\u003d (NH4Cl) - [(NH 3) + (HCl)];
= -315.39 - [-46.19 + (-92.31) = -176.85 kJ.

The thermochemical equation will look like:

The heat released during the reaction of 10 liters of ammonia in this reaction is determined from the proportion:

22,4 : -176,85 = 10 : X; x \u003d 10 (-176.85) / 22.4 \u003d -78.97 kJ; Q = 78.97 kJ.

Answer: 78.97 kJ.

7. Calculate the thermal effect of the reaction under standard conditions: Fe 2 O 3 (t) + 3 CO (g) \u003d 2 Fe (t) + 3 CO 2 (g), if the heat of formation: Fe 2 O 3 (t) \u003d - 821.3 kJ / mol; CO (g ) = – 110.5 kJ/mol;

CO 2 (g) \u003d - 393.5 kJ / mol.

Fe 2 O 3 (t) + 3 CO (g) \u003d 2 Fe (t) + 3 CO 2 (g),

Knowing the standard thermal effects of combustion of the initial substances and reaction products, we calculate the thermal effect of the reaction under standard conditions:

16. Dependence of the rate of a chemical reaction on temperature. Van't Hoff's rule. Temperature coefficient of reaction.

Only collisions between active molecules lead to reactions, the average energy of which exceeds the average energy of the participants in the reaction.

When a certain activation energy E is communicated to molecules (excess energy above the average), the potential energy of interaction of atoms in molecules decreases, bonds within molecules weaken, molecules become reactive.

The activation energy is not necessarily supplied from the outside; it can be imparted to some part of the molecules by redistributing the energy during their collisions. According to Boltzmann, among N molecules there is the following number of active molecules N   with increased energy  :

N N e – E / RT

where E is the activation energy, showing the necessary excess of energy compared to the average level that molecules must have in order for the reaction to become possible; the rest of the designations are well known.

During thermal activation for two temperatures T 1 and T 2 the ratio of the rate constants will be:

, (2) , (3)

which allows you to determine the activation energy by measuring the reaction rate at two different temperatures T 1 and T 2 .

An increase in temperature by 10 0 increases the reaction rate by 2–4 times (approximate van't Hoff rule). The number showing how many times the reaction rate (and hence the rate constant) increases with an increase in temperature by 10 0 is called the temperature coefficient of the reaction:

 (4) .(5)

This means, for example, that with an increase in temperature by 100 0 for a conventionally accepted increase in the average rate by 2 times ( = 2), the reaction rate increases by 2 10 , i.e. approximately 1000 times, and when  = 4 - 4 10 , i.e. 1000000 times. The van't Hoff rule is applicable to reactions occurring at relatively low temperatures in a narrow range. The sharp increase in the reaction rate with increasing temperature is explained by the fact that the number of active molecules increases exponentially.

25. Van't Hoff chemical reaction isotherm equation.

In accordance with the law of mass action for an arbitrary reaction

and A + bB = cC + dD

The equation for the rate of a direct reaction can be written:

,

and for the rate of the reverse reaction:

.

As the reaction proceeds from left to right, the concentrations of substances A and B will decrease and the rate of the forward reaction will decrease. On the other hand, as reaction products C and D accumulate, the reaction rate will increase from right to left. There comes a moment when the speeds υ 1 and υ 2 become the same, the concentrations of all substances remain unchanged, therefore,

,

Where K c = k 1 / k 2 =

.

The constant value K c, equal to the ratio of the rate constants of the direct and reverse reactions, quantitatively describes the state of equilibrium through the equilibrium concentrations of the starting substances and the products of their interaction (in terms of their stoichiometric coefficients) and is called the equilibrium constant. The equilibrium constant is constant only for a given temperature, i.e.

K c \u003d f (T). The equilibrium constant of a chemical reaction is usually expressed as a ratio, the numerator of which is the product of the equilibrium molar concentrations of the reaction products, and the denominator is the product of the concentrations of the starting substances.

If the reaction components are a mixture of ideal gases, then the equilibrium constant (K p) is expressed in terms of the partial pressures of the components:

.

For the transition from K p to K with we use the equation of state P · V = n · R · T. Because the

, then P = C·R·T. .

It follows from the equation that K p = K s, provided that the reaction proceeds without changing the number of moles in the gas phase, i.e. when (c + d) = (a + b).

If the reaction proceeds spontaneously at constant P and T or V and T, then the valuesG and F of this reaction can be obtained from the equations:

,

where C A, C B, C C, C D are the nonequilibrium concentrations of the initial substances and reaction products.

,

where P A, P B, P C, P D are the partial pressures of the initial substances and reaction products.

The last two equations are called the van't Hoff chemical reaction isotherm equations. This relation makes it possible to calculate the values ​​of G and F of the reaction, to determine its direction at different concentrations of the initial substances.

It should be noted that both for gas systems and for solutions with participation of solids in the reaction (i.e. for heterogeneous systems), the concentration of the solid phase is not included in the expression for the equilibrium constant, since this concentration is practically constant. So for the reaction

2 CO (g) \u003d CO 2 (g) + C (t)

the equilibrium constant is written as

.

The dependence of the equilibrium constant on temperature (for temperature T 2 relative to temperature T 1) is expressed by the following van't Hoff equation:

,

where Н 0 is the thermal effect of the reaction.

For an endothermic reaction (the reaction proceeds with the absorption of heat), the equilibrium constant increases with increasing temperature, the system, as it were, resists heating.

34. Osmosis, osmotic pressure. Van't Hoff equation and osmotic coefficient.

Osmosis is the spontaneous movement of solvent molecules through a semipermeable membrane that separates solutions of different concentrations from a solution of a lower concentration to a solution of a higher concentration, which leads to the dilution of the latter. As a semi-permeable membrane, through small holes of which only small solvent molecules can selectively pass and large or solvated molecules or ions are retained, a cellophane film is often used - for high molecular weight substances, and for low molecular weight - a film of copper ferrocyanide. The process of solvent transfer (osmosis) can be prevented if an external hydrostatic pressure is applied to a solution with a higher concentration (under equilibrium conditions this will be the so-called osmotic pressure, denoted by the letter ). To calculate the value of  in solutions of non-electrolytes, the empirical Van't Hoff equation is used:

where C is the molar concentration of the substance, mol/kg;

R is the universal gas constant, J/mol K.

The value of osmotic pressure is proportional to the number of molecules (in the general case, the number of particles) of one or more substances dissolved in a given volume of solution, and does not depend on their nature and the nature of the solvent. In solutions of strong or weak electrolytes, the total number of individual particles increases due to the dissociation of molecules; therefore, it is necessary to introduce the appropriate proportionality coefficient, called the isotonic coefficient, into the equation for calculating the osmotic pressure.

i C R T,

where i is the isotonic coefficient, calculated as the ratio of the sum of the numbers of ions and undissociated electrolyte molecules to the initial number of molecules of this substance.

So, if the degree of electrolyte dissociation, i.e. the ratio of the number of molecules decomposed into ions to the total number of molecules of the solute is  and the electrolyte molecule decomposes into n ions, then the isotonic coefficient is calculated as follows:

i = 1 + (n – 1) ,(i > 1).

For strong electrolytes, you can take  = 1, then i = n, and the coefficient i (also greater than 1) is called the osmotic coefficient.

The phenomenon of osmosis is of great importance for plant and animal organisms, since the membranes of their cells in relation to solutions of many substances have the properties of a semipermeable membrane. In pure water, the cell swells strongly, in some cases up to the rupture of the shell, and in solutions with a high salt concentration, on the contrary, it decreases in size and shrinks due to the large loss of water. Therefore, when preserving food, a large amount of salt or sugar is added to them. Cells of microorganisms in such conditions lose a significant amount of water and die.

The heat of reaction (heat effect of the reaction) is the amount of heat released or absorbed Q. If heat is released during the reaction, such a reaction is called exothermic, if heat is absorbed, the reaction is called endothermic.

The heat of reaction is determined based on the first law (beginning) of thermodynamics, whose mathematical expression in its simplest form for chemical reactions is the equation:

Q = ΔU + рΔV (2.1)

where Q is the heat of reaction, ΔU is the change in internal energy, p is the pressure, ΔV is the change in volume.

Thermochemical calculation consists in determining the thermal effect of the reaction. In accordance with equation (2.1), the numerical value of the heat of reaction depends on the method of its implementation. In an isochoric process carried out at V=const, the heat of reaction Q V =Δ U, in isobaric process at p=const thermal effect Q P =Δ H. Thus, the thermochemical calculation is in determining the amount of change in either internal energy or enthalpy during a reaction. Since the vast majority of reactions proceed under isobaric conditions (for example, these are all reactions in open vessels proceeding at atmospheric pressure), when bringing thermochemical calculations, ΔН is almost always calculated . If aΔ H<0, то реакция экзотермическая, если же Δ H>0, then the reaction is endothermic.

Thermochemical calculations are made using either Hess's law, according to which the thermal effect of a process does not depend on its path, but is determined only by the nature and state of the initial substances and products of the process, or, most often, a consequence of Hess's law: the thermal effect of a reaction is equal to the sum of heats (enthalpies ) the formation of products minus the sum of the heats (enthalpies) of formation of the reactants.

In calculations according to the Hess law, the equations of auxiliary reactions are used, the thermal effects of which are known. The essence of operations in calculations according to the Hess law is that such algebraic operations are performed on the equations of auxiliary reactions that lead to a reaction equation with an unknown thermal effect.

Example 2.1. Determination of the heat of reaction: 2CO + O 2 \u003d 2CO 2 ΔH - ?

We use the reactions as auxiliary: 1) C + O 2 \u003d C0 2;Δ H 1 = -393.51 kJ and 2) 2C + O 2 = 2CO;Δ H 2 \u003d -220.1 kJ, whereΔ N/iΔ H 2 - thermal effects of auxiliary reactions. Using the equations of these reactions, one can obtain the equation for a given reaction if the auxiliary equation 1) is multiplied by two and equation 2) is subtracted from the result. Therefore, the unknown heat of a given reaction is:

Δ H = 2Δ H1-Δ H 2 \u003d 2 (-393.51) - (-220.1) \u003d -566.92 kJ.

If a consequence of the Hess law is used in the thermochemical calculation, then for the reaction expressed by the equation aA+bB=cC+dD, the relation is used:

ΔН =(сΔНоbr,с + dΔHobr D) - (аΔНоbr A + bΔН arr,c) (2.2)

where ΔН is the heat of reaction; ΔH o br - heat (enthalpy) of formation, respectively, of the reaction products C and D and reagents A and B; c, d, a, b - stoichiometric coefficients.

The heat (enthalpy) of formation of a compound is the heat effect of a reaction during which 1 mol of this compound is formed from simple substances that are in thermodynamically stable phases and modifications 1 *. For example , the heat of formation of water in the vapor state is equal to half the heat of reaction, expressed by the equation: 2H 2 (g)+ About 2 (d)= 2H 2 O(g).The unit of heat of formation is kJ/mol.

In thermochemical calculations, the heats of reactions are usually determined for standard conditions, for which formula (2.2) takes the form:

ΔН°298 = (сΔН° 298, arr, С + dΔH° 298, o 6 p, D) - (аΔН° 298, arr A + bΔН° 298, arr, c)(2.3)

where ΔH° 298 is the standard heat of reaction in kJ (the standard value is indicated by the superscript "0") at a temperature of 298K, and ΔH° 298,arr are the standard heats (enthalpies) of formation also at a temperature of 298K. ΔH° values ​​298 rev.are defined for all connections and are tabular data. 2 * - see application table.

Example 2.2. Calculation of standard heat p e shares expressed by the equation:

4NH 3 (r) + 5O 2 (g) \u003d 4NO (g) + 6H 2 O (g).

According to the corollary of Hess's law, we write 3*:

Δ H 0 298 = (4Δ H 0 298. o b p . No+6∆H0 298. code N20) - 4∆H0 298 arr. NH h. Substituting the tabular values ​​of the standard heats of formation of the compounds presented in the equation, we get:Δ H °298= (4(90.37) + 6(-241.84)) - 4(-46.19) = - 904.8 kJ.

The negative sign of the heat of reaction indicates that the process is exothermic.

In thermochemistry, it is customary to indicate thermal effects in reaction equations. Such equations with a designated thermal effect are called thermochemical. For example, the thermochemical equation of the reaction considered in example 2.2 is written:

4NH 3 (g) + 50 2 (g) \u003d 4NO (g) + 6H 2 0 (g);Δ H° 29 8 = - 904.8 kJ.

If the conditions differ from the standard ones, in practical thermochemical calculations it allows Xia approximation use:Δ H ≈Δ N° 298 (2.4) Expression (2.4) reflects the weak dependence of the heat of reaction on the conditions of its occurrence.