Reversible and irreversible chemical reactions. Chemical balance. Shift of equilibrium under the influence of various factors
Chemical equilibrium
Chemical reactions proceeding in one direction are called irreversible.
Most chemical processes are reversible. This means that under the same conditions both forward and reverse reactions occur (especially if we are talking about closed systems).
For example:
a) reaction
$CaCO_3(→)↖(t)CaO+CO_2$
in an open system is irreversible;
b) the same reaction
$CaCO_3⇄CaO+CO_2$
in a closed system is reversible.
Let us consider in more detail the processes occurring during reversible reactions, for example, for a conditional reaction:
Based on the law of mass action, the rate of direct reaction
$(υ)↖(→)=k_(1) C_(A)^(α) C_(B)^(β)$
Since the concentrations of substances $A$ and $B$ decrease over time, the rate of the direct reaction also decreases.
The appearance of reaction products means the possibility of a reverse reaction, and over time the concentrations of substances $C$ and $D$ increase, which means the rate of the reverse reaction also increases:
$(υ)↖(→)=k_(2) C_(C)^(γ) C_(D)^(δ)$
Sooner or later a state will be reached in which the rates of forward and reverse reactions become equal
${υ}↖{→}={υ}↖{←}$
The state of the system in which the rate of the forward reaction is equal to the rate of the reverse reaction is called chemical equilibrium.
In this case, the concentrations of reactants and reaction products remain unchanged. They are called equilibrium concentrations. At the macro level, it seems that overall nothing is changing. But in fact, both the forward and reverse processes continue to occur, but at the same speed. Therefore, such equilibrium in the system is called mobile And dynamic.
Equilibrium constant
Let us denote the equilibrium concentrations of substances as $[A], [B], [C], [D]$.
Then since $(υ)↖(→)=(υ)↖(←), k_(1)·[A]^(α)·[B]^(β)=k_(2)·[C]^ (γ)·[D]^(δ)$, whence
$([C]^(γ)·[D]^(δ))/([A]^(α)·[B]^(β))=(k_1)/(k_2)=K_(equal) $
where $γ, δ, α, β$ are exponents equal to the coefficients in the reversible reaction; $K_(equal)$ is the chemical equilibrium constant.
The resulting expression quantitatively describes the state of equilibrium and is a mathematical expression of the law of mass action for equilibrium systems.
At a constant temperature, the equilibrium constant is a constant value for a given reversible reaction. It shows the relationship between the concentrations of reaction products (numerator) and starting substances (denominator), which is established at equilibrium.
Equilibrium constants are calculated from experimental data, determining the equilibrium concentrations of the starting substances and reaction products at a certain temperature.
The value of the equilibrium constant characterizes the yield of reaction products and the completeness of its progress. If we get $K_(equal) >> 1$, this means that at equilibrium $[C]^(γ)·[D]^(δ) >> [A]^(α)·[B]^( β)$, i.e. the concentrations of reaction products prevail over the concentrations of the starting substances, and the yield of reaction products is high.
At $K_(equal)
$CH_3COOC_2H_5+H_2O⇄CH_3COOH+C_2H_5OH$
equilibrium constant
$K_(equal)=(·)/(·)$
at $20°С$ the value is $0.28$ (i.e. less than $1$). This means that a significant portion of the ester was not hydrolyzed.
In the case of heterogeneous reactions, the expression of the equilibrium constant includes the concentrations of only those substances that are in the gas or liquid phase. For example, for the reaction
the equilibrium constant is expressed as follows:
$K_(equal)=(^2)/()$
The value of the equilibrium constant depends on the nature of the reactants and temperature.
The constant does not depend on the presence of a catalyst, since it changes the activation energy of both forward and reverse reactions by the same amount. The catalyst can only accelerate the onset of equilibrium without affecting the value of the equilibrium constant.
Shift of equilibrium under the influence of various factors
The state of equilibrium is maintained indefinitely under constant external conditions: temperature, concentration of starting substances, pressure (if gases participate in the reaction or are formed).
By changing these conditions, it is possible to transfer the system from one equilibrium state to another that meets the new conditions. This transition is called displacement or shift in equilibrium.
Let's consider different ways to shift the equilibrium using the example of the reaction between nitrogen and hydrogen to form ammonia:
$N_2+3H_2⇄2HN_3+Q$
$K_(equal)=(^2)/(·^3)$
Effect of changing the concentration of substances
When nitrogen $N_2$ and hydrogen $H_2$ are added to the reaction mixture, the concentration of these gases increases, which means the rate of the direct reaction increases. The equilibrium shifts to the right, towards the reaction product, i.e. towards ammonia $NH_3$.
The same conclusion can be drawn by analyzing the expression for the equilibrium constant. As the concentration of nitrogen and hydrogen increases, the denominator increases, and since $K_(equal)$ is a constant value, the numerator must increase. Thus, the amount of the reaction product $NH_3$ in the reaction mixture will increase.
An increase in the concentration of the ammonia reaction product $NH_3$ will lead to a shift of equilibrium to the left, towards the formation of the starting substances. This conclusion can be drawn based on similar reasoning.
Effect of Pressure Change
A change in pressure affects only those systems where at least one of the substances is in a gaseous state. As pressure increases, the volume of gases decreases, which means their concentration increases.
Let's assume that the pressure in a closed system is increased, for example, $2$ times. This means that the concentrations of all gaseous substances ($N_2, H_2, NH_3$) in the reaction we are considering will increase by $2$ times. In this case, the numerator in the expression for $K_(equal)$ will increase by 4 times, and the denominator by $16$ times, i.e. the balance will be disrupted. To restore it, the concentration of ammonia must increase and the concentrations of nitrogen and hydrogen must decrease. The balance will shift to the right. A change in pressure has virtually no effect on the volume of liquids and solids, i.e. does not change their concentration. Consequently, the state of chemical equilibrium of reactions that do not involve gases does not depend on pressure.
Effect of temperature change
As the temperature increases, as you know, the rates of all reactions (exo- and endothermic) increase. Moreover, an increase in temperature has a greater effect on the rate of those reactions that have a high activation energy, and therefore are endothermic.
Thus, the rate of the reverse reaction (endothermic in our example) increases more than the rate of the forward reaction. The equilibrium will shift towards the process accompanied by the absorption of energy.
The direction of the equilibrium shift can be predicted using Le Chatelier's principle (1884):
If an external influence is exerted on a system that is in equilibrium (concentration, pressure, temperature changes), then the equilibrium shifts to the side that weakens this influence.
Let's draw conclusions:
- with an increase in the concentration of reactants, the chemical equilibrium of the system shifts towards the formation of reaction products;
- with an increase in the concentration of reaction products, the chemical equilibrium of the system shifts towards the formation of the starting substances;
- with increasing pressure, the chemical equilibrium of the system shifts towards the reaction in which the volume of gaseous substances formed is smaller;
- with increasing temperature, the chemical equilibrium of the system shifts towards the endothermic reaction;
- with decreasing temperature - towards an exothermic process.
Le Chatelier's principle is applicable not only to chemical reactions, but also to many other processes: evaporation, condensation, melting, crystallization, etc. In the production of the most important chemical products, Le Chatelier's principle and calculations arising from the law of mass action make it possible to find such conditions for carrying out chemical processes that provide maximum yield of the desired substance.
Chemical reactions proceeding in one direction are called irreversible.
Most chemical processes are reversible. This means that under the same conditions both forward and reverse reactions occur (especially if we are talking about closed systems). 
For example:
a) reaction
in an open system irreversible;
b) the same reaction
in a closed system reversible.
Chemical equilibrium
Let us consider in more detail the processes occurring during reversible reactions, for example, for a conditional reaction:
Based on the law of mass action rate of forward reaction:
Since the concentrations of substances A and B decrease over time, the rate of the direct reaction also decreases.
The appearance of reaction products means the possibility of a reverse reaction, and over time the concentrations of substances C and D increase, which means that the reverse reaction speed.
Sooner or later a state will be reached in which the rates of forward and reverse reactions become equal = .
The state of the system in which the rate of the forward reaction is equal to the rate of the reverse reaction is called chemical equilibrium.
In this case, the concentrations of reactants and reaction products remain unchanged. They are called equilibrium concentrations. At the macro level, it seems that overall nothing is changing. But in fact, both the forward and reverse processes continue to occur, but at the same speed. Therefore, such equilibrium in the system is called mobile and dynamic.
Let us denote the equilibrium concentrations of substances [A], [B], [C], [D]. Then since = , k 1 [A] α [B] β = k 2 [C] γ [D] δ , where
where α, β, γ, δ are exponents, equal to the coefficients in the reversible reaction; K equal - chemical equilibrium constant.
The resulting expression quantitatively describes state of equilibrium and is a mathematical expression of the law of mass action for equilibrium systems.
At a constant temperature, the equilibrium constant is constant value for a given reversible reaction. It shows the relationship between the concentrations of reaction products (numerator) and starting substances (denominator), which is established at equilibrium.
Equilibrium constants are calculated from experimental data, determining the equilibrium concentrations of starting substances and reaction products at a certain temperature.
The value of the equilibrium constant characterizes the yield of reaction products and the completeness of its progress. If we get K » 1, this means that at equilibrium [C] γ [D] δ "[A] α [B] β , i.e., the concentrations of reaction products prevail over the concentrations of the starting substances, and the yield of reaction products is high.
At K equal to « 1, the yield of reaction products is correspondingly low. For example, for the hydrolysis reaction of acetic acid ethyl ester
equilibrium constant:
at 20 °C it has a value of 0.28 (that is, less than 1).
This means that a significant portion of the ester was not hydrolyzed.
In the case of heterogeneous reactions, the expression of the equilibrium constant includes the concentrations of only those substances that are in the gas or liquid phase. For example, for the reaction
The equilibrium constant is expressed as follows:
The value of the equilibrium constant depends on the nature of the reactants and temperature.
The constant does not depend on the presence of a catalyst, since it changes the activation energy of both the forward and reverse reactions by the same amount. The catalyst can only accelerate the onset of equilibrium without affecting the value of the equilibrium constant.
The state of equilibrium is maintained indefinitely under constant external conditions: temperature, concentration of starting substances, pressure (if gases participate in the reaction or are formed).
By changing these conditions, it is possible to transfer the system from one equilibrium state to another that meets the new conditions. This transition is called displacement or shift in equilibrium.
Let's consider different ways to shift the equilibrium using the example of the reaction between nitrogen and hydrogen to form ammonia:
Effect of changing the concentration of substances
When nitrogen N2 and hydrogen H2 are added to the reaction mixture, the concentration of these gases increases, which means the rate of forward reaction increases. The equilibrium shifts to the right, towards the reaction product, that is, towards ammonia NH 3.
N 2 +3H 2 → 2NH 3
The same conclusion can be drawn by analyzing the expression for the equilibrium constant. As the concentration of nitrogen and hydrogen increases, the denominator increases, and since K is equal. - the value is constant, the numerator must increase. Thus, the amount of the reaction product NH 3 in the reaction mixture will increase.
An increase in the concentration of the ammonia reaction product NH 3 will lead to a shift of equilibrium to the left, towards the formation of the starting substances. This conclusion can be drawn based on similar reasoning.
Effect of Pressure Change
A change in pressure affects only those systems where at least one of the substances is in a gaseous state. As pressure increases, the volume of gases decreases, which means their concentration increases.
Let's assume that the pressure in a closed system is increased, for example, by 2 times. This means that the concentrations of all gaseous substances (N 2, H 2, NH 3) in the reaction under consideration will increase by 2 times. In this case, the numerator in the expression for K equal will increase by 4 times, and the denominator by 16 times, i.e., the equilibrium will be disrupted. To restore it, the concentration of ammonia must increase and the concentrations of nitrogen and hydrogen must decrease. The balance will shift to the right. A change in pressure has virtually no effect on the volume of liquids and solids, i.e. it does not change their concentration. Hence, the state of chemical equilibrium of reactions that do not involve gases does not depend on pressure.
Effect of temperature change
As the temperature increases, the rates of all reactions (exo- and endothermic) increase. Moreover, an increase in temperature has a greater effect on the rate of those reactions that have a higher activation energy, which means endothermic.
Thus, the rate of the reverse reaction (endothermic) increases more than the rate of the forward reaction. The equilibrium will shift towards the process accompanied by the absorption of energy.
The direction of the equilibrium shift can be predicted using Le Chatelier's principle:
If an external influence is exerted on a system that is in equilibrium (concentration, pressure, temperature changes), then the equilibrium shifts to the side that weakens this influence.
Thus:
As the concentration of reactants increases, the chemical equilibrium of the system shifts towards the formation of reaction products;
As the concentration of reaction products increases, the chemical equilibrium of the system shifts towards the formation of the starting substances;
As pressure increases, the chemical equilibrium of the system shifts towards the reaction in which the volume of gaseous substances formed is smaller;
As the temperature increases, the chemical equilibrium of the system shifts towards the endothermic reaction;
As the temperature decreases, it moves towards an exothermic process.
Le Chatelier's principle is applicable not only to chemical reactions, but also to many other processes: evaporation, condensation, melting, crystallization, etc. In the production of the most important chemical products, Le Chatelier's principle and calculations arising from the law of mass action make it possible to find such conditions to carry out chemical processes that provide maximum yield of the desired substance.
Reference material for taking the test:
Mendeleev table
Solubility table
Among the numerous classifications of types of reactions, for example those that are determined by the thermal effect (exothermic and endothermic), by changes in the oxidation states of substances (redox), by the number of components participating in them (decomposition, compounds) and so on, reactions occurring in two mutual directions, otherwise called reversible . An alternative to reversible reactions are reactions irreversible, during which the final product (precipitate, gaseous substance, water) is formed. Among these reactions are the following:
Exchange reactions between salt solutions, during which either an insoluble precipitate is formed - CaCO 3:
Ca(OH) 2 + K 2 CO 3 → CaCO 3↓ + 2KON (1)
or a gaseous substance - CO 2:
3 K 2 CO 3 + 2H 3 RO 4 →2K 3 RO 4 + 3 CO 2+ 3H 2 O (2)
or a slightly dissociable substance is obtained - H 2 O:
2NaOH + H 2 SO 4 → Na 2 SO 4 + 2 H 2O(3)
If we consider a reversible reaction, then it proceeds not only in the forward direction (in reactions 1,2,3 from left to right), but also in the opposite direction. An example of such a reaction is the synthesis of ammonia from gaseous substances - hydrogen and nitrogen:
3H 2 + N 2 ↔ 2NH 3 (4)
Thus, a chemical reaction is called reversible if it proceeds not only in the forward direction (→), but also in the reverse direction (←) and is indicated by the symbol (↔).
The main feature of this type of reaction is that reaction products are formed from the starting substances, but at the same time, the starting reagents are formed from the same products. If we consider reaction (4), then in a relative unit of time, simultaneously with the formation of two moles of ammonia, their decomposition will occur with the formation of three moles of hydrogen and one mole of nitrogen. Let us denote the rate of direct reaction (4) by the symbol V 1, then the expression for this rate will take the form:
V 1 = kˑ [Н 2 ] 3 ˑ , (5)
where the value “k” is defined as the rate constant of a given reaction, the values [H 2 ] 3 and correspond to the concentrations of the starting substances raised to powers corresponding to the coefficients in the reaction equation. In accordance with the principle of reversibility, the rate of the reverse reaction will take the expression:
V 2 = kˑ 2 (6)
At the initial moment of time, the rate of the forward reaction takes on the greatest value. But gradually the concentrations of the starting reagents decrease and the reaction rate slows down. At the same time, the rate of the reverse reaction begins to increase. When the rates of forward and reverse reactions become the same (V 1 = V 2), state of equilibrium , at which there is no longer a change in the concentrations of both the initial and the resulting reagents.
It should be noted that some irreversible reactions should not be taken literally. Let us give an example of the most frequently cited reaction of a metal with an acid, in particular, zinc with hydrochloric acid:
Zn + 2HCl = ZnCl 2 + H 2 (7)
In fact, zinc, when dissolved in acid, forms a salt: zinc chloride and hydrogen gas, but after some time the rate of the direct reaction slows down as the concentration of salt in the solution increases. When the reaction practically stops, a certain amount of hydrochloric acid will be present in the solution along with zinc chloride, so reaction (7) should be given in the following form:
2Zn + 2HCl = 2ZnНCl + H2 (8)
Or in the case of the formation of an insoluble precipitate obtained by merging solutions of Na 2 SO 4 and BaCl 2:
Na 2 SO 4 + BaCl 2 = BaSO 4 ↓ + 2NaCl (9)
the precipitated salt BaSO 4, albeit to a small extent, will dissociate into ions:
BaSO 4 ↔ Ba 2+ + SO 4 2- (10)
Therefore, the concepts of irreversible and irreversible reactions are relative. But nevertheless, both in nature and in the practical activities of people, these reactions are of great importance. For example, combustion processes of hydrocarbons or more complex organic substances, such as alcohol:
CH 4 + O 2 = CO 2 + H 2 O (11)
2C 2 H 5 OH + 5O 2 = 4CO 2 + 6H 2 O (12)
are completely irreversible processes. It would be considered a happy dream of humanity if reactions (11) and (12) were reversible! Then it would be possible to synthesize gas and gasoline and alcohol again from CO 2 and H 2 O! On the other hand, reversible reactions such as (4) or oxidation of sulfur dioxide:
SO 2 + O 2 ↔ SO 3 (13)
are basic in the production of ammonium salts, nitric acid, sulfuric acid, and other inorganic and organic compounds. But these reactions are reversible! And in order to obtain the final products: NH 3 or SO 3, it is necessary to use such technological methods as: changing the concentrations of reagents, changing pressure, increasing or decreasing the temperature. But this will already be the subject of the next topic: “Shift in chemical equilibrium.”
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All chemical reactions can be divided into two groups: irreversible and reversible reactions. Irreversible reactions proceed to completion - until one of the reactants is completely consumed. Reversible reactions do not proceed to completion: in a reversible reaction, none of the reactants are completely consumed. This difference is due to the fact that an irreversible reaction can only proceed in one direction. A reversible reaction can occur in both forward and reverse directions.
Let's look at two examples.
Example 1. The interaction between zinc and concentrated nitric acid proceeds according to the equation:
With a sufficient amount of nitric acid, the reaction will only end when all the zinc has dissolved. In addition, if you try to carry out this reaction in the opposite direction - passing nitrogen dioxide through a solution of zinc nitrate, then metallic zinc and nitric acid will not work - this reaction cannot proceed in the opposite direction. Thus, the interaction of zinc with nitric acid is an irreversible reaction.
Example 2. Ammonia synthesis proceeds according to the equation:
If you mix one mole of nitrogen with three moles of hydrogen, create conditions in the system that are favorable for the reaction to occur, and after a sufficient time, analyze the gas mixture, the results of the analysis will show that not only the reaction product (ammonia) will be present in the system, but also the initial substances (nitrogen and hydrogen). If now, under the same conditions, not a nitrogen-hydrogen mixture, but ammonia is placed as the starting substance, then it will be possible to find that part of the ammonia will decompose into nitrogen and hydrogen, and the final ratio between the quantities of all three substances will be the same as in that case , when starting from a mixture of nitrogen and hydrogen. Thus, ammonia synthesis is a reversible reaction.
In equations of reversible reactions, arrows can be used instead of the equal sign; they symbolize the reaction occurring in both forward and reverse directions.
In Fig. Figure 68 shows the change in the rates of forward and reverse reactions over time. At first, when mixing the starting substances, the rate of the forward reaction is high, and the rate of the reverse reaction is zero. As the reaction proceeds, the starting substances are consumed and their concentrations fall.
Rice. 63. Change in the speed of forward and reverse reactions over time.
As a result, the rate of the forward reaction decreases. At the same time, reaction products appear and their concentration increases. As a result, a reverse reaction begins to occur, and its speed gradually increases. When the rates of forward and reverse reactions become equal, chemical equilibrium occurs. Thus, in the last example, an equilibrium is established between nitrogen, hydrogen and ammonia.
Chemical equilibrium is called dynamic equilibrium. This emphasizes that at equilibrium both forward and reverse reactions occur, but their rates are the same, as a result of which changes in the system are not noticeable.
A quantitative characteristic of chemical equilibrium is a value called the chemical equilibrium constant. Let's consider it using the example of the iodide-hydrogen synthesis reaction:
According to the law of mass action, the rates of forward and reverse reactions are expressed by the equations:
At equilibrium, the rates of forward and reverse reactions are equal to each other, hence
The ratio of the rate constants of the forward and reverse reactions is also a constant. It is called the equilibrium constant of this reaction (K):
From here finally
On the left side of this equation are those concentrations of interacting substances that are established at equilibrium - equilibrium concentrations. The right side of the equation is a constant (at constant temperature) quantity.
It can be shown that in the general case of a reversible reaction
the equilibrium constant will be expressed by the equation:
Here, large letters indicate the formulas of substances, and small letters indicate coefficients in the reaction equation.
Thus, at a constant temperature, the equilibrium constant of a reversible reaction is a constant value showing the ratio between the concentrations of reaction products (numerator) and starting substances (denominator) that is established at equilibrium.
The equilibrium constant equation shows that under equilibrium conditions, the concentrations of all substances participating in the reaction are related to each other. A change in the concentration of any of these substances entails changes in the concentrations of all other substances; as a result, new concentrations are established, but the ratio between them again corresponds to the equilibrium constant.
The numerical value of the equilibrium constant, to a first approximation, characterizes the yield of a given reaction. For example, when the reaction yield is high, because in this case
i.e., at equilibrium, the concentrations of the reaction products are much greater than the concentrations of the starting substances, and this means that the yield of the reaction is high. When (for a similar reason) the yield of the reaction is low.
In the case of heterogeneous reactions, the expression of the equilibrium constant, as well as the expression of the law of mass action (see § 58), includes the concentrations of only those substances that are in the gas or liquid phase. For example, for the reaction
the equilibrium constant has the form:
The value of the equilibrium constant depends on the nature of the reacting substances and on the temperature. It does not depend on the presence of catalysts. As already mentioned, the equilibrium constant is equal to the ratio of the rate constants of the forward and reverse reactions. Since the catalyst changes the activation energy of both forward and reverse reactions by the same amount (see § 60), it does not affect the ratio of their rate constants.
Therefore, the catalyst does not affect the value of the equilibrium constant and, therefore, can neither increase nor decrease the yield of the reaction. It can only speed up or slow down the onset of equilibrium.