Reversible in chemical kinetics such reactions are called, which simultaneously and independently proceed in two directions - forward and backward, but with different rates. For reversible reactions, it is characteristic that some time after their beginning, the rates of the forward and reverse reactions become equal and a state of chemical equilibrium sets in.
All chemical reactions are reversible, but under certain conditions some of them can proceed in only one direction until the starting products practically disappear. Such reactions are called irreversible... Usually, irreversible reactions are reactions in which at least one reaction product is removed from the reaction region (in the case of a reaction in solutions, it precipitates or is released in the form of a gas), or reactions that are accompanied by a large positive thermal effect. When ionic reactions, the reaction is practically irreversible if it results in the formation of a very slightly soluble or poorly dissociated substance.
The concept of reaction reversibility considered here does not coincide with the concept of thermodynamic reversibility. A kinetic reversible reaction in a thermodynamic sense can proceed irreversibly. In order for the reaction to be called reversible in the thermodynamic sense, the speed of the direct process must differ infinitely little from the speed of the reverse process and, therefore, the process as a whole must proceed infinitely slowly.
In ideal gas mixtures and in ideal liquid solutions, the rates of simple (one-stage) reactions obey law of mass action... The rate of chemical reaction (1.1) is described by equation (1.2), and in the case of a direct reaction it can be represented as:
where is the rate constant of the direct reaction.
Likewise, the feedback rate is:
In equilibrium, therefore:
This equation expresses the law of mass action for chemical equilibrium in ideal systems; K - konstan t and ravnove s and I.
The reaction constant allows you to find the equilibrium composition of the reaction mixture under the given conditions.
The mass action law for the reaction rates can be explained as follows.
For a reaction to occur, a collision of molecules of the initial substances is necessary, i.e. the molecules must approach each other at a distance of the order of atomic dimensions. The probability of finding in a certain small volume in this moment l molecules of substance L, m molecules of substance M, etc. proportional ....., therefore, the number of collisions per unit of volume per unit of time is proportional to this value; from this equation (1.4) follows.
Chemical reactions often go to the end, i.e. the initial products are completely consumed in the course of a chemical reaction and new substances are formed - reaction products. Such reactions go only in one direction - towards the direct reaction.
Irreversible reactions - reactions, as a result of which the starting materials are completely converted into the final reaction products.
Irreversible reactions occur in three cases if:
1) an insoluble substance is formed, i.e. precipitation falls .
For instance:
BaCl 2 + H 2 SO 4 → BaSO 4 + 2HCl - this is a molecular equation
Now let's write each molecule into ions, except for the substance that precipitated (see the charges of ions in the table "Solubility of hydroxides and salts" on the last flyleaf of the textbook).
Let's cancel the identical ions on the right and left sides of the equation and write out the ions that remain:
| Ba | 2+ | + | SO | 2− | → | BaSO 4 ↓ | is the short ionic equation |
| 4 |
Thus, according to the abbreviated ionic equation, it is seen that the precipitate is formed from barium ions (Ba 2+) and sulfate ions (SO 4 2 –).
2) a gaseous substance is formed, i.e. gas released:
For instance:
Na 2 S + 2HCl → 2NaCl + H 2 S - molecular equation
2Na + + S 2− + 2H + + 2Cl - → 2 Na + + 2 Cl - + H 2 S - complete ionic equation
S 2− + 2H + → H 2 S - short ionic equation
3) formed water:
for instance:
KOH + HNO 3 → KNO 3 + H 2 O - molecular equation
K + + OH - + H + + NO 3 - → K + + NO 3 - + H 2 O - complete ionic equation
OH - + H + → H 2 O - short ionic equation
However, there are not many irreversible reactions; most reactions proceed in two directions (towards the formation of new substances, and vice versa, towards the decomposition of new substances into the initial reaction products), i.e. are reversible.
Reversible reactions - chemical reactions that take place in two opposite directions - forward and backward.
For example: reaction of formation of ammonia from hydrogen(H 2 ) and nitrogen(N 2) follows the reaction:
3H 2 + N 2 → 2NH 3
and the resulting ammonia molecules decompose intoH 2 and N 2 (i.e. for the starting materials):
2NH 3 → 3H 2 + N 2, therefore, in total, these two reactions write: 3H 2 + N 2 ↔ 2NH 3 (arrow ↔ shows the reaction proceeding in two directions).
In reversible reactions, a moment comes when the rate of the direct reaction (the rate of formation of new substances) becomes equal to the rate of the reverse reaction (the rate of formation of the initial reaction products from new substances) - equilibrium occurs.
Chemical equilibrium - the state of a chemically reversible process, in which the rate of the forward reaction is equal to the rate of the reverse reaction.
Chemical equilibrium is dynamic (i.e. mobile) because when it occurs, the reaction does not stop, but only the concentration of substances does not change. This means that the amount of new substances formed is equal to the amount of the original substances. At constant temperature and pressure, equilibrium in a reversible reaction can be maintained indefinitely.
In practice (in the laboratory, in production), they are most often interested in direct reactions.
The equilibrium of a reversible system can be shifted by changing one of the equilibrium conditions (concentration, temperature, or pressure).
The law of displacement of chemical equilibrium (Le Chatelier principle): if a system in equilibrium is acted upon by changing one of the equilibrium conditions, then the state of chemical equilibrium will shift towards a decrease in this effect.
1) When increasing the concentration of reactants, the equilibrium always shifts to the right - towards the direct reaction (i.e. towards the formation of new substances).
2) When increasing pressure by compressing the system, therefore, and increasing the concentration of reactants (only for substances in a gaseous state), the equilibrium of the system shifts towards a smaller number of gas molecules.
3) When increase in temperature balance shifts:
a) with an endothermic reaction (a reaction proceeding with the absorption of heat) - to the right (in the direction of a direct reaction);
b) in case of an exothermic reaction (a reaction proceeding with the release of heat) - to the left (in the direction of the reverse reaction).
4) When falling temperature balance shifts:
a) with an endothermic reaction (a reaction proceeding with the absorption of heat) - to the left (in the direction of the reverse reaction);
b) with an exothermic reaction (a reaction proceeding with the release of heat) - to the right (towards a direct reaction).
Endothermic reactions in writing are indicated by a sign at the end of the reaction "+ Q" or
"∆Н\u003e 0", exothermic - by the sign at the end of the reaction "- Q" or "∆Н< 0».
For example: let's analyze where the balance in the system is shifting:
2NO 2 (g) ↔ 2NO (g) + O 2 (g) + Q
a) an increase in the concentration of reactants
b) decrease in temperature
c) an increase in temperature
d) increase in pressure
Decision:
a) an increase in the concentration of reactants - the equilibrium shifts to the right (because according to the law of mass action, the greater the concentration of substances, the higher the reaction rate);
b) a decrease in temperature (since the reaction is endothermic) - a shift to the left;
c) increase in temperature - shift to the right;
Reversible and irreversible chemical reactions. Chemical equilibrium. Equilibrium shift under the influence of various factors
Chemical equilibrium
Chemical reactions proceeding in one direction are called irreversible.
Most of the chemical processes are reversible... This means that under the same conditions, both direct and reverse reactions occur (especially when it comes to closed systems).
For instance:
a) reaction
$ CaCO_3 (→) ↖ (t) CaO + CO_2 $
irreversible in an open system;
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 conditioned reaction:
Based on the law of mass action, the speed of the direct reaction
$ (υ) ↖ (→) \u003d k_ (1) C_ (A) ^ (α) C_ (B) ^ (β) $
Since the concentrations of substances $ A $ and $ B $ decrease with time, the rate of the direct reaction also decreases.
The appearance of reaction products means the possibility of a reverse reaction, and with time the concentrations of substances $ C $ and $ D $ increase, which means that the rate of the reverse reaction also increases:
$ (υ) ↖ (→) \u003d k_ (2) C_ (C) ^ (γ) C_ (D) ^ (δ) $
Sooner or later, a state will be reached in which the rates of the 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.
At the same time, the concentrations of reactants and reaction products remain unchanged. They are called equilibrium concentrations... At the macro level, nothing seems to change overall. But in fact, both direct and reverse processes continue to go on, but with equal speed. Therefore, such an equilibrium in the system is called mobile and dynamic.
Equilibrium constant
Let us denote the equilibrium concentrations of substances by $ [A], [B], [C], [D] $.
Then since $ (υ) ↖ (→) \u003d (υ) ↖ (←), k_ (1) · [A] ^ (α) · [B] ^ (β) \u003d k_ (2) · [C] ^ (γ) · [D] ^ (δ) $, whence
$ ([C] ^ (γ) · [D] ^ (δ)) / ([A] ^ (α) · [B] ^ (β)) \u003d (k_1) / (k_2) \u003d K_ (equal) $
where $ γ, δ, α, β $ - exponents equal to the coefficients in the reversible reaction; $ K_ (equal) $ - 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.
When constant temperature equilibrium constant is a constant value for a given reversible reaction. It shows the ratio between the concentrations of the reaction products (numerator) and the initial substances (denominator), which is established at equilibrium.
The equilibrium constants are calculated from experimental data by determining the equilibrium concentrations of the starting materials and reaction products at a certain temperature.
The value of the equilibrium constant characterizes the yield of the reaction products, the completeness of its course. If you get $ K_ (equal) \u003e\u003e 1 $, this means that in equilibrium $ [C] ^ (γ) · [D] ^ (δ) \u003e\u003e [A] ^ (α) · [B] ^ ( β) $, i.e., the concentrations of the reaction products prevail over the concentrations of the initial substances, and the yield of the reaction products is high.
For $ K_ (equal)
$ CH_3COOC_2H_5 + H_2O⇄CH_3COOH + C_2H_5OH $
equilibrium constant
$ K_ (equal) \u003d () / () $
at $ 20 ° С $ is $ 0.28 $ (i.e. less than $ 1 $). This means that a significant part of the ether was not hydrolyzed.
In the case of heterogeneous reactions, the expression for 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) \u003d (^ 2) / () $
The value of the equilibrium constant depends on the nature of the reacting substances and temperature.
The constant does not depend on the presence of a catalyst, since it changes the activation energy of both the direct and reverse reactions by the same amount. The catalyst can only accelerate the onset of equilibrium without affecting the value of the equilibrium constant.
Equilibrium shift under the influence of various factors
The state of equilibrium is maintained for an arbitrarily long time under constant external conditions: temperature, concentration of starting substances, pressure (if gases are involved or formed in the reaction).
By changing these conditions, it is possible to transfer the system from one equilibrium state to another that meets the new conditions. Such a transition is called displacement or balance shift.
Let us consider different ways of shifting equilibrium using the example of the reaction of interaction of nitrogen and hydrogen with the formation of ammonia:
$ N_2 + 3H_2⇄2HN_3 + Q $
$ K_ (equal) \u003d (^ 2) / (^ 3) $
Effect of changes in the concentration of substances
When nitrogen $ N_2 $ and hydrogen $ Н_2 $ are added to the reaction mixture, the concentration of these gases increases, which means that 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 made by analyzing the expression for the equilibrium constant. With an increase in the concentration of nitrogen and hydrogen, the denominator increases, and since $ K_ (equal) $ is a constant value, the numerator must increase. Thus, the amount of the reaction product $ NH_3 $ will increase in the reaction mixture.
An increase in the concentration of the reaction product of ammonia $ NH_3 $ will lead to a shift of the equilibrium to the left, towards the formation of the initial substances. This conclusion can be drawn on the basis of similar reasoning.
Influence of pressure changes
A change in pressure affects only those systems where at least one of the substances is in a gaseous state. With increasing pressure, the volume of gases decreases, which means that their concentration increases.
Suppose that the pressure in a closed system has 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 $ 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 disturbed. To restore it, the concentration of ammonia must increase and the concentration of nitrogen and hydrogen must decrease. The balance will shift to the right. The change in pressure has practically no effect on the volume of liquid and solids, i.e. does not change their concentration. Consequently, the state of chemical equilibrium of reactions in which gases do not participate does not depend on pressure.
Effect of temperature change
As the temperature rises, 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 endothermic.
Thus, the rate of the reverse reaction (in our example, 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 displacement of the equilibrium can be predicted using Le Chatelier's principle (1884):
If an external influence is exerted on a system in equilibrium (concentration, pressure, temperature changes), then equilibrium shifts in the direction that weakens this effect.
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 initial substances;
- with increasing pressure, the chemical equilibrium of the system shifts in the direction of the reaction in which the volume of formed gaseous substances is less;
- as the temperature rises, the chemical equilibrium of the system shifts towards the endothermic reaction;
- with decreasing temperature - towards the 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 ensure the maximum yield of the desired substance.
What is a reversible response? This is a chemical process that takes place in two mutually opposite directions. Let us consider the main characteristics of such transformations, as well as their distinctive parameters.
What is the essence of balance
Reversible chemical reactions do not produce specific products. For example, during the oxidation of sulfur oxide (4) simultaneously with the production of sulfur oxide (6), the starting components are again formed.
Irreversible processes involve the complete transformation of the interacting substances, a similar reaction is accompanied by the production of one or more reaction products.
Decomposition reactions are examples of irreversible interactions. For example, when potassium permanganate is heated, metal manganate, manganese oxide (4) are formed, and oxygen gas is also released.
A reversible reaction does not imply the formation of precipitation, the release of gases. This is precisely its main difference from irreversible interaction.
Chemical equilibrium is a state of an interacting system in which one or more chemical reactions can be reversible, provided that the rates of the processes are equal.
If the system is in dynamic equilibrium, there is no change in temperature, concentration of reagents, and other parameters in given interval time.
Equilibrium displacement conditions
The equilibrium of a reversible reaction can be explained using Le Chatelier's rule. Its essence lies in the fact that when an external influence is exerted on the system, which is initially in dynamic equilibrium, a change in the reaction is observed in the direction opposite to the influence. Any reversible reaction using this principle can be shifted in the desired direction in the event of changes in temperature, pressure, and also the concentration of the interacting substances.
Le Chatelier's principle "works" only for gaseous reagents, solid and liquid substances not taken into account. There is a mutually inverse relationship between pressure and volume, determined by the Mendeleev-Clapeyron equation. If the volume of the initial gaseous components is greater than the reaction products, then to change the equilibrium to the right, it is important to increase the mixture pressure.
For example, when carbon monoxide (2) is transformed into carbon dioxide, 2 moles of carbon monoxide and 1 mole of oxygen enter into the reaction. This produces 2 moles of carbon monoxide (4).
If, according to the condition of the problem, this reversible reaction should be shifted to the right, it is necessary to increase the pressure.
The concentration of reactants also has a significant effect on the course of the process. According to Le Chatelier's principle, in the case of an increase in the concentration of the initial components, the equilibrium of the process shifts towards the product of their interaction.
In this case, the decrease (withdrawal from the reaction mixture) of the resulting product contributes to the direct process.
In addition to pressure, concentration, a change in temperature also has a significant effect on the course of the reverse or direct reaction. When the initial mixture is heated, the equilibrium is shifted towards the endothermic process.
Examples of reversible reactions
Let us consider the ways of shifting the equilibrium towards the formation of reaction products on a specific process.
2CO + O 2 -2CO 2
This reaction is a homogeneous process, since all substances are in one (gaseous) state.
On the left side of the equation, there are 3 volumes of components, after the interaction this indicator decreased, 2 volumes are formed. For the direct process to proceed, it is necessary to increase the pressure of the reaction mixture.
Given that the reaction is exothermic, the temperature is lowered to produce carbon dioxide.
The equilibrium of the process will shift towards the formation of the reaction product with an increase in the concentration of one of the initial substances: oxygen or carbon monoxide.
Conclusion
Reversible and irreversible reactions play an important role in human life. The metabolic processes occurring in our body are associated with a systematic shift in chemical equilibrium. In chemical production, optimal conditions are used to direct the reaction in the right direction.
DEFINITION
Chemical reaction called the transformation of substances in which there is a change in their composition and (or) structure.
The reaction is possible with a favorable ratio of energy and entropy factors. If these factors balance each other, the state of the system does not change. In such cases, the system is said to be in equilibrium.
Chemical reactions proceeding in one direction are called irreversible. Most chemical reactions are reversible. This means that under the same conditions, both direct and reverse reactions occur (especially when it comes to closed systems).
The state of the system at 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 (equilibrium concentrations).
Equilibrium constant
Consider the reaction for producing ammonia:
N 2 (g) + 3H 2 (g) ↔ 2 NH 3 (g)
Let us write expressions for calculating the rates of forward (1) and reverse (2) reactions:
1 \u003d k 1 [H 2] 3
2 \u003d k 2 2
The rates of forward and backward reactions are equal, therefore, we can write:
k 1 3 \u003d k 2 2
k 1 / k 2 \u003d 2/3
The ratio of two constants is a constant value. Equilibrium constant is the ratio of the rate constants of the forward and reverse reactions.
K \u003d 2/3
Expressed in general terms, the equilibrium constant is:
mA + nB ↔ pC + qD
K \u003d [C] p [D] q / [A] m [B] n
The equilibrium constant is the ratio of the products of the concentrations of the reaction products raised to powers equal to their stoichiometric coefficients to the product of the concentrations of the starting substances raised to powers equal to their stoichiometric coefficients.
If K is expressed in terms of equilibrium concentrations, then K s is most often denoted. It is also possible to calculate K for gases through their partial pressures. In this case, K is designated as K p. There is a relationship between K c and K p:
K p \u003d K s × (RT) Δn,
where Δn is the change in the number of all moles of gases during the transition from reagents to products, R is the universal gas constant.
K does not depend on the concentration, pressure, volume and presence of the catalyst and depends on the temperature and nature of the reacting substances. If K is much less than 1, then there are more starting materials in the mixture, and in the case of much more than 1, there are more products in the mixture.
Heterogeneous equilibrium
Consider the reaction
CaCO 3 (tv) ↔ CaO (tv) + CO 2 (g)
The expression for the equilibrium constant does not include the concentrations of the components in the solid phase, therefore
Chemical equilibrium occurs in the presence of all components of the system, but the equilibrium constant does not depend on the concentration of substances in the solid phase. Chemical equilibrium is a dynamic process. K gives information about the progress of the reaction, and ΔG - about its direction. They are related to each other by the relationship:
ΔG 0 \u003d -R × T × lnK
ΔG 0 \u003d -2.303 × R × T × logK
Chemical equilibrium shift. Le Chatelier's principle
From point of view technological processes reversible chemical reactions are not profitable, since you need to have knowledge of how to increase the yield of the reaction product, i.e. it is necessary to learn how to shift the chemical equilibrium towards the reaction products.
Consider a reaction in which it is necessary to increase the yield of ammonia:
N 2 (g) + 3H 2 (g) ↔ 2NH 3 (g), ΔН< 0
In order to shift the balance towards a direct or reverse reaction, it is necessary to use le Chatelier principle: if the system, which is in equilibrium, is acted upon by some factor from outside (increase or decrease in temperature, pressure, volume, concentration of substances), then the system counteracts this effect.
For example, if the temperature is increased in an equilibrium system, then out of 2 possible reactions, the one that will be endothermic will go; if the pressure is increased, then the equilibrium will shift towards the reaction with a large number mole of substances; if the volume in the system is increased, then the displacement of equilibrium will be aimed at increasing the pressure; if the concentration of one of the starting substances is increased, then out of 2 possible reactions the one will proceed, which will lead to a decrease in the equilibrium concentration of the product.
So, as applied to the considered reaction, in order to increase the yield of ammonia, it is necessary to increase the concentration of the starting substances; lower the temperature, since the direct reaction is exothermic, increase the pressure or decrease the volume.
Examples of problem solving
EXAMPLE 1
