Chemical equilibrium inherent reversible reactions and is not typical for irreversible chemical reactions.
Often, when carrying out a chemical process, the initial reactants are completely converted into reaction products. For example:
Cu + 4HNO 3 = Cu(NO 3) 2 + 2NO 2 + 2H 2 O
It is impossible to obtain metallic copper by carrying out the reaction in the opposite direction, because given the reaction is irreversible. In such processes, reactants are completely converted into products, i.e. the reaction proceeds to completion.
But the bulk of chemical reactions reversible, i.e. the reaction is likely to occur in parallel in the forward and reverse directions. In other words, the reactants are only partially converted into products and the reaction system will consist of both reactants and products. System in in this case is in a state chemical equilibrium.
In reversible processes, initially the direct reaction has maximum speed, which gradually decreases due to a decrease in the number of reagents. The reverse reaction, on the contrary, initially has minimum speed, which increases as products accumulate. Eventually, a moment comes when the rates of both reactions become equal—the system reaches a state of equilibrium. When a state of equilibrium occurs, the concentrations of the components remain unchanged, but the chemical reaction does not stop. That. – this is a dynamic (moving) state. For clarity, here is the following figure:
Let's say there is a certain reversible chemical reaction:
a A + b B = c C + d D
then, based on the law of mass action, we write down expressions for straightυ 1 and reverseυ 2 reactions:
v1 = k 1 ·[A] a ·[B] b
v2 = k 2 ·[C] c ·[D] d
Able chemical equilibrium, the rates of forward and reverse reactions are equal, i.e.:
k 1 ·[A] a ·[B] b = k 2 ·[C] c ·[D] d
we get
TO= k 1 / k 2 = [C] c [D] d ̸ [A] a [B] b
Where K =k 1 / k 2 – equilibrium constant.
For any reversible process, under given conditions k is a constant value. It does not depend on the concentrations of substances, because When the amount of one of the substances changes, the amounts of other components also change.
When the conditions of a chemical process change, the equilibrium may shift.
Factors influencing the shift in equilibrium:
- changes in concentrations of reagents or products,
- pressure change,
- temperature change,
- adding a catalyst to the reaction medium.
Le Chatelier's principle
All of the above factors influence the shift in chemical equilibrium, which obeys Le Chatelier's principle: If you change one of the conditions under which the system is in a state of equilibrium - concentration, pressure or temperature - then the equilibrium will shift in the direction of the reaction that counteracts this change. Those. equilibrium tends to shift in a direction leading to a decrease in the influence of the influence that led to a violation of the state of equilibrium.
So, let us consider separately the influence of each of their factors on the state of equilibrium.
Influence changes in concentrations of reactants or products let's show with an example Haber process:
N 2(g) + 3H 2(g) = 2NH 3(g)
If, for example, nitrogen is added to an equilibrium system consisting of N 2 (g), H 2 (g) and NH 3 (g), then the equilibrium should shift in a direction that would contribute to a decrease in the amount of hydrogen towards its original value, those. in the direction of the formation of additional ammonia (to the right). At the same time, the amount of hydrogen will decrease. When hydrogen is added to the system, the equilibrium will also shift towards the formation of a new amount of ammonia (to the right). Whereas the introduction of ammonia into the equilibrium system, according to Le Chatelier's principle , will cause a shift in equilibrium towards the process that is favorable for the formation of starting substances (to the left), i.e. The ammonia concentration should decrease through the decomposition of some of it into nitrogen and hydrogen.
A decrease in the concentration of one of the components will shift the equilibrium state of the system towards the formation of this component.
Influence pressure changes makes sense if gaseous components take part in the process under study and there is a change in the total number of molecules. If the total number of molecules in the system remains permanent, then the change in pressure does not affect on its balance, for example:
I 2(g) + H 2(g) = 2HI (g)
If the total pressure of an equilibrium system is increased by decreasing its volume, then the equilibrium will shift towards decreasing volume. Those. towards decreasing the number gas in system. In reaction:
N 2(g) + 3H 2(g) = 2NH 3(g)
from 4 gas molecules (1 N 2 (g) and 3 H 2 (g)) 2 gas molecules are formed (2 NH 3 (g)), i.e. the pressure in the system decreases. As a result, an increase in pressure will contribute to the formation of an additional amount of ammonia, i.e. the equilibrium will shift towards its formation (to the right).
If the temperature of the system is constant, then a change in the total pressure of the system will not lead to a change in the equilibrium constant TO.
Temperature change system affects not only the displacement of its equilibrium, but also the equilibrium constant TO. If additional heat is imparted to an equilibrium system at constant pressure, then the equilibrium will shift towards the absorption of heat. Consider:
N 2(g) + 3H 2(g) = 2NH 3(g) + 22 kcal
So, as you can see, the direct reaction proceeds with the release of heat, and the reverse reaction with absorption. As the temperature increases, the equilibrium of this reaction shifts towards the decomposition reaction of ammonia (to the left), because she appears and weakens external influence- temperature increase. On the contrary, cooling leads to a shift in equilibrium in the direction of ammonia synthesis (to the right), because the reaction is exothermic and resists cooling.
Thus, an increase in temperature favors a shift chemical equilibrium towards the endothermic reaction, and the temperature drop towards the exothermic process . Equilibrium constants all exothermic processes decrease with increasing temperature, and endothermic processes increase.
>> Chemistry: Chemical equilibrium and methods of shifting it In reversible processes, the rate of a direct reaction is initially maximum, and then decreases due to the fact that the concentrations of starting substances consumed in the formation of reaction products decrease. On the contrary, the rate of the reverse reaction, minimal at the beginning, increases as the concentration of reaction products increases. Finally, a moment comes when the rates of the forward and reverse reactions become equal.
The state of a chemical reversible process is called chemical equilibrium if the rate of the forward reaction is equal to the rate of the reverse reaction.
Chemical equilibrium is dynamic (mobile), since when it occurs, the reaction does not stop, only the concentrations of the components remain unchanged, that is, per unit time the same amount of reaction products is formed as is converted into the starting substances. At constant temperature and pressure, the equilibrium of a reversible reaction can be maintained indefinitely.
In production, they are most often interested in the preferential occurrence of a direct reaction. For example, in the production of ammonia, sulfur oxide (VI). nitric oxide (II). How to derive a system from a state of equilibrium? How does a change in the external conditions under which this or that reversible chemical process occurs affect it?
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Let's consider each of these cases.
Disturbance of equilibrium due to a change in the concentration of any of the substances participating in the reaction. Let hydrogen, hydrogen iodide and iodine vapor be in equilibrium with each other at a certain temperature and pressure. Let us introduce an additional amount of hydrogen into the system. According to the law of mass action, an increase in the concentration of hydrogen will entail an increase in the rate of the forward reaction - the HI synthesis reaction, while the rate of the reverse reaction will not change. The reaction will now proceed faster in the forward direction than in the reverse direction. As a result of this, the concentrations of hydrogen and iodine vapor will decrease, which will slow down the forward reaction, and the concentration of HI will increase, which will accelerate the reverse reaction. After some time, the rates of the forward and reverse reactions will become equal again, and a new equilibrium will be established. But at the same time, the concentration of HI will now be higher than it was before adding , and the concentration will be lower.
The process of changing concentrations caused by an imbalance is called a displacement or equilibrium shift. If at the same time there is an increase in the concentrations of substances on the right side of the equation (and, of course, at the same time a decrease in the concentrations of substances on the left), then they say that the equilibrium shifts to the right, i.e., in the direction of the direct reaction; when the concentrations change in the opposite direction, they speak of a shift in equilibrium to the left - in the direction of the reverse reaction. In the example considered, the equilibrium has shifted to the right. At the same time, the substance, the increase in concentration of which caused an imbalance, entered into a reaction - its concentration decreased.
Thus, with an increase in the concentration of any of the substances participating in the equilibrium, the equilibrium shifts towards the consumption of this substance; When the concentration of any substance decreases, the equilibrium shifts towards the formation of this substance.
Disturbance of equilibrium due to changes in pressure (by decreasing or increasing the volume of the system). When gases are involved in a reaction, equilibrium may be disrupted when the volume of the system changes.
Consider the effect of pressure on the reaction between nitrogen monoxide and oxygen:
Let a mixture of gases be in chemical equilibrium at a certain temperature and pressure. Without changing the temperature, we increase the pressure so that the volume of the system decreases by 2 times. At the first moment, the partial pressures and concentrations of all gases will double, but at the same time the ratio between the rates of forward and reverse reactions will change - the equilibrium will be disrupted.
In fact, before the pressure increased, the gas concentrations had equilibrium values , and , and the rates of the forward and reverse reactions were the same and were determined by the equations:
At the first moment after compression, the gas concentrations will double compared to their initial values and will be equal to , and , respectively. In this case, the rates of forward and reverse reactions will be determined by the equations:
Thus, as a result of increasing pressure, the rate of the forward reaction increased 8 times, and the reverse reaction only 4 times. The equilibrium in the system will be disrupted - the forward reaction will prevail over the reverse one. After the speeds become equal, equilibrium will be established again, but the quantity in the system will increase, and the equilibrium will shift to the right.
It is easy to see that the unequal change in the rates of forward and reverse reactions is due to the fact that on the left and right sides of the equation of the reaction under consideration the number of gas molecules is different: one molecule of oxygen and two molecules of nitrogen monoxide (three gas molecules in total) are converted into two gas molecules - nitrogen dioxide. The pressure of a gas is the result of its molecules hitting the walls of the container; other things being equal, the higher the number of molecules contained in a given volume of gas, the higher the gas pressure. Therefore, a reaction that occurs with an increase in the number of gas molecules leads to an increase in pressure, and a reaction that occurs with a decrease in the number of gas molecules leads to a decrease in pressure.
With this in mind, the conclusion about the effect of pressure on chemical equilibrium can be formulated as follows:
When the pressure increases by compressing the system, the equilibrium shifts towards a decrease in the number of gas molecules, i.e. towards a decrease in pressure; when the pressure decreases, the equilibrium shifts towards an increase in the number of gas molecules, i.e. towards an increase in pressure.
In the case when the reaction proceeds without changing the number of gas molecules, the equilibrium is not disturbed during compression or expansion of the system. For example, in the system
equilibrium is not disturbed when volume changes; the HI output is independent of pressure.
Disequilibrium due to temperature changes. The equilibrium of the vast majority of chemical reactions shifts with temperature changes. The factor that determines the direction of the equilibrium shift is the sign of the thermal effect of the reaction. It can be shown that when the temperature increases, the equilibrium shifts in the direction of the endothermic reaction, and when it decreases, in the direction of the exothermic reaction.
Thus, ammonia synthesis is an exothermic reaction
Therefore, as the temperature increases, the equilibrium in the system shifts to the left - towards the decomposition of ammonia, since this process occurs with the absorption of heat.
Conversely, the synthesis of nitric oxide (II) is an endothermic reaction:
Therefore, as the temperature increases, the equilibrium in the system shifts to the right - towards the formation.
The patterns that appear in the considered examples of chemical imbalance are special cases general principle, which determines the influence various factors to equilibrium systems. This principle, known as Le Chatelier's principle, when applied to chemical equilibria, can be formulated as follows:
If any impact is exerted on a system that is in equilibrium, then as a result of the processes occurring in it, the equilibrium will shift in such a direction that the impact will decrease.
Indeed, when one of the substances participating in the reaction is introduced into the system, the equilibrium shifts towards the consumption of this substance. “When the pressure increases, it shifts so that the pressure in the system decreases; when the temperature increases, the equilibrium shifts towards the endothermic reaction - the temperature in the system drops.
Le Chatelier's principle applies not only to chemical, but also to various physicochemical equilibria. A shift in equilibrium when the conditions of processes such as boiling, crystallization, and dissolution change occurs in accordance with Le Chatelier’s principle.
Chemical equilibrium is maintained as long as the conditions in which the system is located remain unchanged. Changing conditions (concentration of substances, temperature, pressure) causes an imbalance. After some time, the chemical equilibrium is restored, but under new, different from previous conditions. Such a transition of a system from one equilibrium state to another is called displacement(shift) of equilibrium. The direction of displacement obeys Le Chatelier's principle.
As the concentration of one of the starting substances increases, the equilibrium shifts towards greater consumption of this substance, and the direct reaction intensifies. A decrease in the concentration of the starting substances shifts the equilibrium towards the formation of these substances, as the reverse reaction intensifies. An increase in temperature shifts the equilibrium towards an endothermic reaction, while a decrease in temperature shifts the equilibrium towards an exothermic reaction. An increase in pressure shifts the equilibrium towards decreasing quantities gaseous substances, that is, towards smaller volumes occupied by these gases. On the contrary, as pressure decreases, the equilibrium shifts towards increasing amounts of gaseous substances, that is, towards larger volumes formed by gases.
Example 1.
How will an increase in pressure affect the equilibrium state of the following reversible gas reactions:
a) SO 2 + C1 2 =SO 2 CI 2;
b) H 2 + Br 2 = 2НВr.
Solution:
We use Le Chatelier's principle, according to which an increase in pressure in the first case (a) shifts the equilibrium to the right, towards a smaller amount of gaseous substances occupying a smaller volume, which weakens the external influence of the increased pressure. In the second reaction (b), the quantities of gaseous substances, both the starting materials and the reaction products, are equal, as are the volumes they occupy, so pressure has no effect and the equilibrium is not disturbed.
Example 2.
In the reaction of ammonia synthesis (–Q) 3H 2 + N 2 = 2NH 3 + Q, the forward reaction is exothermic, the reverse reaction is endothermic. How should the concentration of reactants, temperature and pressure be changed to increase the yield of ammonia?
Solution:
To shift the balance to the right you need to:
a) increase the concentrations of H 2 and N 2;
b) reduce the concentration (removal from the reaction sphere) of NH 3;
c) lower the temperature;
d) increase the pressure.
Example 3.
The homogeneous reaction between hydrogen chloride and oxygen is reversible:
4HC1 + O 2 = 2C1 2 + 2H 2 O + 116 kJ.
1. What effect will the following have on the equilibrium of the system?
a) increase in pressure;
b) increase in temperature;
c) introduction of a catalyst?
Solution:
a) In accordance with Le Chatelier's principle, an increase in pressure leads to a shift in equilibrium towards the direct reaction.
b) An increase in t° leads to a shift in equilibrium towards the reverse reaction.
c) The introduction of a catalyst does not shift the equilibrium.
2. In what direction will the chemical equilibrium shift if the concentration of reactants is doubled?
Solution:
υ → = k → 0 2 0 2 ; υ 0 ← = k ← 0 2 0 2
After increasing concentrations, the rate of the forward reaction became:
υ → = k → 4 = 32 k → 0 4 0
that is, it increased by 32 times compared to the initial speed. Similarly, the rate of the reverse reaction increases 16 times:
υ ← = k ← 2 2 = 16k ← [H 2 O] 0 2 [C1 2 ] 0 2 .
The increase in the rate of the forward reaction is 2 times greater than the increase in the rate of the reverse reaction: the equilibrium shifts to the right.
Example 4.
IN Which side will the equilibrium of a homogeneous reaction shift:
PCl 5 = PC1 3 + Cl 2 + 92 KJ,
if you increase the temperature by 30 °C, knowing that the temperature coefficient of the forward reaction is 2.5, and the reverse reaction is 3.2?
Solution:
Since the temperature coefficients of the forward and reverse reactions are not equal, increasing the temperature will have different effects on the change in the rates of these reactions. Using Van't Hoff's rule (1.3), we find the rates of forward and reverse reactions when the temperature increases by 30 °C:
υ → (t 2) = υ → (t 1)=υ → (t 1)2.5 0.1 30 = 15.6υ → (t 1);
υ ← (t 2) = υ ← (t 1) =υ → (t 1)3.2 0.1 30 = 32.8υ ← (t 1)
An increase in temperature increased the rate of the forward reaction by 15.6 times, and the reverse reaction by 32.8 times. Consequently, the equilibrium will shift to the left, towards the formation of PCl 5.
Example 5.
How will the rates of forward and reverse reactions change in the isolated system C 2 H 4 + H 2 ⇄ C 2 H 6 and where will the equilibrium shift when the volume of the system increases by 3 times?
Solution:
The initial rates of forward and reverse reactions are as follows:
υ 0 = k 0 0 ; υ 0 = k 0 .
An increase in the volume of the system causes a decrease in the concentrations of reactants by 3 times, hence the change in the rate of forward and reverse reactions will be as follows:
υ 0 = k = 1/9υ 0
υ = k = 1/3υ 0
The decrease in the rates of forward and reverse reactions is not the same: the rate of the reverse reaction is 3 times (1/3: 1/9 = 3) higher than the rate of the reverse reaction, therefore the equilibrium will shift to the left, to the side where the system occupies a larger volume, that is, towards the formation of C 2 H 4 and H 2.
All chemical reactions are, in principle, reversible.
This means that both the interaction of reagents and the interaction of products occurs in the reaction mixture. In this sense, the distinction between reactants and products is conditional. The direction of a chemical reaction is determined by the conditions of its conduct (temperature, pressure, concentration of substances).
Many reactions have one preferential direction and to carry out such reactions in the opposite direction requires extreme conditions. In such reactions, almost complete conversion of reactants into products occurs.Example. Iron and sulfur, when heated moderately, react with each other to form iron (II) sulfide; FeS is stable under such conditions and practically does not decompose into iron and sulfur:
At 200 atm and 400 0C, the maximum NH3 content in the reaction mixture is reached, equal to 36% (by volume). With a further increase in temperature, due to the increased occurrence of the reverse reaction, the volume fraction of ammonia in the mixture decreases.
Forward and reverse reactions occur simultaneously in opposite directions.
In all reversible reactions, the rate of the forward reaction decreases and the rate of the reverse reaction increases until both rates are equal and equilibrium is established.
In a state of equilibrium, the rates of forward and reverse reactions become equal.
LE CHATELIER'S PRINCIPLE. SHIFT OF CHEMICAL EQUILIBRIUM.
The position of chemical equilibrium depends on the following reaction parameters: temperature, pressure and concentration. The influence that these factors have on chemical reaction, obey the pattern that was expressed in general view in 1884 by the French scientist Le Chatelier. The modern formulation of Le Chatelier's principle is as follows:
1. Effect of temperature. In each reversible reaction, one of the directions corresponds to an exothermic process, and the other to an endothermic process.
2. Effect of pressure. In all reactions involving gaseous substances, accompanied by a change in volume due to a change in the amount of substance during the transition from starting substances to products, the equilibrium position is affected by the pressure in the system.
The influence of pressure on the equilibrium position obeys the following rules:Thus, during the transition from starting substances to products, the volume of gases was halved. This means that with increasing pressure, the equilibrium shifts towards the formation of NH3, as evidenced by the following data for the ammonia synthesis reaction at 400 0C:
3. Effect of concentration. The influence of concentration on the state of equilibrium is subject to the following rules:
