reversibility of chemical reactions. Chemical equilibrium
Reversible and irreversible chemical reactions. chemical balance. Equilibrium shift under the influence of various factors
Chemical equilibrium
Chemical reactions that proceed in the same 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 the direct reaction
$(υ)↖(→)=k_(1) C_(A)^(α) C_(B)^(β)$
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 over time, the concentrations of substances $C$ and $D$ increase, which means that 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 the forward and reverse reactions will become equal
${υ}↖{→}={υ}↖{←}$
The state of a 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 the reactants and reaction products remain unchanged. They are called equilibrium concentrations. At the macro level, it seems that in general nothing changes. But in fact, both direct and reverse processes continue to go on, but at the same speed. Therefore, this equilibrium in the system is called mobile And dynamic.
Equilibrium constant
Let us denote the equilibrium concentrations of substances $[A], [B], [C], [D]$.
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 the 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 ratio between the concentrations of the reaction products (numerator) and starting materials (denominator), which is established at equilibrium.
Equilibrium constants are calculated from experimental data by determining the equilibrium concentrations of the initial substances and reaction products at a certain temperature.
The value of the equilibrium constant characterizes the yield of reaction products, the completeness of its course. If $K_(equal) >> 1$ is obtained, this means that at equilibrium $[C]^(γ) [D]^(δ) >> [A]^(α) [B]^( β)$, i.e., the concentrations of the reaction products predominate over the concentrations of the initial substances, and the yield of the reaction products is large.
For $K_(equal)
$CH_3COOC_2H_5+H_2O⇄CH_3COOH+C_2H_5OH$
equilibrium constant
$K_(equal)=( )/( )$
at $20°С$ it has a value of $0.28$ (i.e. less than $1$). This means that a significant part 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 the 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.
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, corresponding to the new conditions. Such a transition is called displacement or balance shift.
Consider different ways to shift the equilibrium using the example of the reaction of the interaction of nitrogen and hydrogen with the formation of ammonia:
$N_2+3H_2⇄2HN_3+Q$
$K_(equal)=(^2)/( ^3)$
The 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 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 drawn 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 ammonia reaction product $NH_3$ will shift the equilibrium to the left, towards the formation of the initial substances. This conclusion can be drawn on the basis of 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 the pressure increases, the volume of gases decreases, which means that their concentration increases.
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 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 upset. 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 practically no effect on the volume of liquid and solid bodies, i.e. does not change their concentration. Therefore, the state of chemical equilibrium of reactions in which gases do not participate is independent of pressure.
Effect of temperature change
With increasing temperature, 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 large activation energy, and hence, endothermic ones.
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 equilibrium shift 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 the equilibrium shifts in the direction 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 starting substances;
- with increasing pressure, the chemical equilibrium of the system shifts towards the reaction in which the volume of gaseous substances formed is less;
- as the temperature rises, the chemical equilibrium of the system shifts towards an endothermic reaction;
- when the temperature drops - in the direction of the exothermic process.
The Le Chatelier 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, the Le Chatelier principle and calculations arising from the law of mass action make it possible to find such conditions for carrying out chemical processes that provide the maximum yield of the desired substance.
retraining of educators.
Department of Natural Science
Topic: “Reversible and irreversible reactions.
chemical balance. Le Chatelier's principle.
Work completed:
Listener group X - 1
chemistry teacher, secondary school №6
Dimitrovgrad
Ulyanovsk region
Lepikhova Tatyana Vasilievna
Scientific adviser:
Department head
natural science
Akhmetov Marat Anvarovich
Ulyanovsk 2009
Reversible and irreversible chemical reactions.
chemical balance.
Le Chatelier's principle.
Goal of the work: 1) The study of the features and patterns of the course of chemical reactions, as a continuation of the formation of ideas about various types of chemical reactions on the basis of reversibility.
2) Generalization and concretization of knowledge about the laws of chemical reactions, the formation of skills and abilities to determine, explain the features and the resulting conditions necessary for the occurrence of a particular reaction. 3) Expand and deepen knowledge about the variety of chemical processes, teach students to compare, analyze, explain, draw conclusions and generalizations. 4) Consider this section of chemical science as the most important in the applied aspect and consider the concept of chemical equilibrium as a particular case of a single law of natural equilibrium, the desire for compensation, the stability of equilibrium in unity with the main form of the existence of matter, movement, dynamics.
Tasks.
Consider the topic: “Reversible and irreversible reactions” on concrete examples, using previous ideas about the rate of chemical reactions.
Continue studying the features of reversible chemical reactions and the formation of ideas about chemical equilibrium as a dynamic state of a reacting system.
To study the principles of shifting chemical equilibrium and teach students to determine the conditions for shifting chemical equilibrium.
To give students an idea of the importance of this topic not only for chemical production, but also for the normal functioning of a living organism and nature as a whole.
Introduction
In nature, in the organisms of living beings, in the process of human physiological activity, in his actions to create conditions different levels: household, defense, industrial, technical, environmental and others - thousands, millions of completely different reactions occur or are carried out, which can be considered from different points vision and classification. We will consider chemical reactions in terms of their reversibility and irreversibility.
It is difficult to overestimate the significance of these concepts: as long as there is a thinking person, the human thought about the reversibility and irreversibility of the processes occurring in his body beats just as much, eternal problem about the extension of human life, the problem of the irreversibility of the consequences of his life, thoughtless attitude to nature.
I want to consider the concept of reversibility and irreversibility of chemical reactions, the concept of chemical equilibrium and the conditions for its shift in a “useful” direction. Introduce theoretical basis with subsequent verification, self-examination of knowledge on this topic, using testing of various typologies. I suppose that “having gone the way” from simple to more difficult tasks, students will have a clear, good knowledge not only on this topic, but will also deepen their knowledge of chemistry.
Chemical reactions are phenomena in which one (or one) substance is converted into another, evidence of this is visible and invisible changes. Visible: changes in color, odor, taste, precipitation, change in color of the indicator, absorption and release of heat. Invisible: A change in the composition of a substance that can be determined using qualitative and analytical reactions. All these reactions can be divided into two types: reversible and irreversible reactions.
irreversible reactions. Reactions that proceed in only one direction and end with the complete conversion of the initial reactants into final substances are called irreversible.
An example of such a reaction is the decomposition of potassium chlorate (bertolet salt) when heated:
2KClO 3 \u003d 2KCl + 3O 2
The reaction will stop when all the potassium chlorate has been converted to potassium chloride and oxygen. There are not many irreversible reactions.
If acid and alkali solutions are drained, salt and water are formed, for example,
HCl + NaOH \u003d NaCl + H 2 O, and if the substances were taken in the right proportions, the solution has a neutral reaction and not even traces of hydrochloric acid and sodium hydroxide remain in it. If you try to carry out a reaction in a solution between the formed substances - sodium chloride and water, then no changes will be found. In such cases, it is said that the reaction of an acid with an alkali is irreversible, i.e. there is no back reaction. Very many reactions are practically irreversible at room temperature, for example,
H 2 + Cl 2 \u003d 2HCl, 2H 2 + O 2 \u003d 2H 2 O, etc.
reversible reactions. Reversible reactions are those that simultaneously proceed in two mutually opposite directions.
Most reactions are reversible. In the equations of reversible reactions, two arrows pointing in opposite directions are placed between the left and right parts. An example of such a reaction is the synthesis of ammonia from hydrogen and nitrogen:
,
∆H = -46.2 kJ/mol
In engineering, reversible reactions are generally unfavorable. Therefore, various methods (changes in temperature, pressure, etc.) make them practically irreversible.
Irreversible are such reactions, during the course of which:
1) the resulting products leave the reaction sphere - they precipitate in the form of a precipitate, are released in the form of a gas, for example
ВаСl 2 + Н 2 SO 4 = ВаSO 4 ↓ + 2НCl
Na 2 CO 3 + 2HCl \u003d 2NaCl + CO 2 ↓ + H 2 O
2) a slightly dissociated compound is formed, for example water:
Hcl + NaOH \u003d H 2 O + NaCl
3) the reaction is accompanied by a large release of energy, for example, the combustion of magnesium
Mg + 1 / 2 O 2 \u003d MgO, ∆H \u003d -602.5 kJ / mol
In the equations of irreversible reactions, an equal sign or an arrow is placed between the left and right parts.
Many reactions are already reversible under ordinary conditions, which means that the reverse reaction proceeds to a noticeable extent. For example, if you try to neutralize with alkali an aqueous solution of a very weak hypochlorous acid, it turns out that the neutralization reaction does not go to the end and the solution has a strongly alkaline environment. This means that the reaction HClO + NaOH NaClO + H 2 O is reversible, i.e. the products of this reaction, reacting with each other, partially pass into the starting compounds. As a result, the solution has an alkaline reaction. The reaction of formation of esters is reversible (the reverse reaction is called saponification): RCOOH + R "OH RCOOR" + H 2 O, many other processes.
Like many other concepts in chemistry, the concept of reversibility is largely arbitrary. Usually, a reaction is considered irreversible, after which the concentrations of the starting substances are so low that they cannot be detected (of course, this depends on the sensitivity of the methods of analysis). When external conditions change (primarily temperature and pressure), an irreversible reaction can become reversible and vice versa. So, at atmospheric pressure and temperatures below 1000 ° C, the reaction 2H 2 + O 2 \u003d 2H 2 O can still be considered irreversible, while at a temperature of 2500 ° C and above, water dissociates into hydrogen and oxygen by about 4%, and at a temperature of 3000 ° С - already by 20%.
At the end of the 19th century German physical chemist Max Bodenstein (1871–1942) studied in detail the processes of formation and thermal dissociation of hydrogen iodine: H 2 + I 2 2HI. By changing the temperature, he could achieve a predominant flow of only the forward or only the reverse reaction, but in the general case, both reactions went simultaneously in opposite directions. There are many such examples. One of the most famous is the ammonia synthesis reaction 3H 2 + N 2 2NH 3; many other reactions are also reversible, for example, the oxidation of sulfur dioxide 2SO 2 + O 2 2SO 3 , reactions of organic acids with alcohols, etc.
A reaction is called reversible if its direction depends on the concentrations of the substances participating in the reaction. For example, in the case of the heterogeneous catalytic reaction N2 + 3H2 = 2NH3 (1), at a low concentration of ammonia in gaseous water and high concentrations of nitrogen and hydrogen, ammonia is formed; on the contrary, at a high concentration of ammonia, it decomposes, the reaction goes in the opposite direction. Upon completion of a reversible reaction, i.e., upon reaching chemical equilibrium, the system contains both the starting materials and the reaction products. The reaction is called irreversible if it can occur only in one direction and ends with the complete transformation of the starting substances into products; an example is the decomposition of explosives. The same reaction, depending on the conditions (temperature, pressure), can be essentially reversible or practically irreversible. A simple (one-stage) reversible reaction consists of two elementary reactions occurring simultaneously, which differ from one another only in the direction of the chemical transformation. The direction of the final reaction accessible to direct observation is determined by which of these reciprocal reactions has a greater speed. For example, the simple reaction N2O4 Û 2NO2 (2) consists of the elementary reactions N2O4?2NO2 and 2NO2?N2O4. M. I. Tyomkin.
CHEMICAL EQUILIBRIUM.
Chemical equilibrium- the state of the system in which the rate of the forward reaction (V 1) is equal to the rate of the reverse reaction (V 2). In chemical equilibrium, the concentrations of substances remain unchanged. Chemical equilibrium has a dynamic character: forward and reverse reactions do not stop at equilibrium.
The state of chemical equilibrium is quantitatively characterized by the equilibrium constant, which is the ratio of the constants of direct (K 1) and reverse (K 2) reactions.
For the reaction mA + nB pC + dD, the equilibrium constant is
K = K 1 / K 2 = ([C] p [D] d) / ([A] m [B] n)
The equilibrium constant depends on the temperature and the nature of the reactants. The larger the equilibrium constant, the more the equilibrium is shifted towards the formation of direct reaction products. In a state of equilibrium, the molecules do not cease to experience collisions, and the interaction between them does not stop, but the concentrations of substances remain constant. These concentrations are called equilibrium.
|
Equilibrium concentration- the concentration of a substance participating in a reversible chemical reaction that has reached a state of equilibrium. |
The equilibrium concentration is indicated by the formula of the substance, taken in square brackets, for example:
With equilibrium (H 2) \u003d or R equilibrium (HI) = .
Like any other concentration, the equilibrium concentration is measured in moles per liter.
If we had taken other concentrations of the initial substances in the examples we have considered, then after reaching equilibrium we would have obtained other values of the equilibrium concentrations. These new values (denoted by asterisks) will be related to the old ones as follows:
.
In general, for a reversible reaction
a A+ b B d D+ f F
in a state of equilibrium at a constant temperature, the relation is observed
This ratio is called law of mass action, which is formulated as follows:
at a constant temperature, the ratio of the product of the equilibrium concentrations of the reaction products, taken in powers equal to their coefficients, to the product of the equilibrium concentrations of the starting substances, taken in powers equal to their coefficients, is a constant value.
Constant value ( TO WITH) is called equilibrium constant this reaction. The index "c" in the designation of this quantity indicates that concentrations were used to calculate the constant.
If the equilibrium constant is large, then the equilibrium is shifted towards the products of the direct reaction, if it is small, then towards the starting materials. If the equilibrium constant is very large, then they say that the reaction " practically irreversible, if the equilibrium constant is very small, then the reaction " practically doesn't work."
Equilibrium constant - for each reversible reaction, the value is constant only at a constant temperature. For the same reaction at different temperatures, the equilibrium constant takes on different values.
The above expression for the law of mass action is valid only for reactions in which all participants are either gases or dissolved substances. In other cases, the equation for the equilibrium constant changes somewhat.
For example, in a reversible reaction proceeding at high temperature
C (gr) + CO 2 2CO (g)
hard graphite C (gr) is involved. Formally, using the law of mass action, we write an expression for the equilibrium constant of this reaction, denoting it TO":
Solid graphite lying at the bottom of the reactor reacts only from the surface, and its "concentration" does not depend on the mass of graphite and is constant for any ratio of substances in the gas mixture.
Multiply the right and left sides of the equation by this constant:
The resulting value is the equilibrium constant of this reaction:
Similarly, for the equilibrium of another reversible reaction also occurring at high temperature,
CaCO 3 (cr) CaO (cr) + CO 2 (g),
we get the equilibrium constant
TO WITH = .
In this case, it is simply equal to the equilibrium concentration of carbon dioxide.
From a metrological point of view, the equilibrium constant is not a single physical quantity. This is a group of quantities with different units of measurement, depending on the specific expression of the constant through equilibrium concentrations. For example, for the reversible reaction of graphite with carbon dioxide [ K c] = 1 mol/l, the equilibrium constant of the reaction of thermal decomposition of calcium carbonate has the same unit of measurement, and the equilibrium constant of the reaction of synthesis of hydrogen iodine is a dimensionless value. In general [ K c] = 1 (mol/l) n .
Shift in chemical equilibrium. Le Chatelier's principle
The transfer of an equilibrium chemical system from one equilibrium state to another is called shift (shift) of chemical equilibrium, which is carried out by changing the thermodynamic parameters of the system - temperature, concentration, pressure. When the equilibrium is shifted in the forward direction, an increase in the yield of products is achieved, and when shifted in the opposite direction, a decrease in the degree of conversion of the reagent. Both can be useful in chemical engineering. Since almost all reactions are reversible to some extent, two problems arise in industry and laboratory practice: how to obtain the product of a "useful" reaction with a maximum yield and how to reduce the yield of products of a "harmful" reaction. In both cases, it becomes necessary to shift the equilibrium either towards the products of the reaction, or towards the starting materials. To learn how to do this, you need to know what determines the equilibrium position of any reversible reaction.
The equilibrium position depends on:
1) on the value of the equilibrium constant (that is, on the nature of the reactants and temperature),
2) on the concentration of substances involved in the reaction and
3) on pressure (for gas systems it is proportional to the concentrations of substances).
For a qualitative assessment of the influence on the chemical equilibrium of all these very different factors, one uses the inherently universal Le Chatelier's principle(French physical chemist and metallurgist Henri Louis Le Chatelier formulated it in 1884), which is applicable to any equilibrium systems, not only chemical ones.
If a system in equilibrium is acted upon from the outside, then the equilibrium in the system will shift in the direction in which this effect is partially compensated.
As an example of the influence on the equilibrium position of the concentrations of substances participating in the reaction, consider the reversible reaction of obtaining hydrogen iodine
H 2 (g) + I 2 (g) 2HI (g) .
According to the law of mass action in a state of equilibrium
.
Let an equilibrium be established in a reactor with a volume of 1 liter at a certain constant temperature, at which the concentrations of all participants in the reaction are the same and equal to 1 mol/l ( = 1 mol/l; = 1 mol/l; = 1 mol/l). Therefore, at this temperature TO WITH= 1. Since the volume of the reactor is 1 liter, n(H 2) \u003d 1 mol, n(I 2) \u003d 1 mol and n(HI) = 1 mol. At time t 1, let's introduce another 1 mol of HI into the reactor, its concentration will become equal to 2 mol/l. But in order to TO WITH remained constant, the concentrations of hydrogen and iodine should increase, and this is possible only due to the decomposition of part of the hydrogen iodine according to the equation
2HI (g) \u003d H 2 (g) + I 2 (g).
Let by the moment of reaching a new state of equilibrium t 2 decomposed x mol of HI and, therefore, an additional 0.5 x mol H 2 and I 2 . New equilibrium concentrations of reaction participants: = (1 + 0.5 x) mol/l; = (1 + 0.5 x) mol/l; = (2 - x) mol/l. Substituting the numerical values of the quantities into the expression of the law of mass action, we obtain the equation
Where x= 0.667. Therefore, = 1.333 mol/l; = 1.333 mol/l; = 1.333 mol/l.
Reaction speed and balance.
Let there be a reversible reaction A + B C + D. If we assume that the forward and reverse reactions take place in one stage, then the rates of these reactions will be directly proportional to the concentrations of the reagents: the rate of the direct reaction v 1 = k 1 [A][B], reverse reaction rate v 2 = k 2 [C][D] (square brackets indicate the molar concentrations of the reagents). It can be seen that as the direct reaction proceeds, the concentrations of the starting substances A and B decrease, respectively, and the rate of the direct reaction also decreases. The rate of the reverse reaction, which is zero at the initial moment (there are no products C and D), gradually increases. Sooner or later, the moment will come when the rates of the forward and reverse reactions will equalize. After that, the concentrations of all substances - A, B, C and D do not change with time. This means that the reaction has reached an equilibrium position, and concentrations of substances that do not change with time are called equilibrium. But, unlike mechanical equilibrium, at which all movement stops, at chemical equilibrium, both reactions - both direct and reverse - continue to go on, but their rates are equal and therefore it seems that no changes occur in the system. There are many ways to prove the flow of forward and reverse reactions after reaching equilibrium. For example, if a little hydrogen isotope - deuterium D 2 is introduced into a mixture of hydrogen, nitrogen and ammonia, which is in an equilibrium position, then a sensitive analysis will immediately detect the presence of deuterium atoms in ammonia molecules. And vice versa, if a little deuterated ammonia NH 2 D is introduced into the system, then deuterium will immediately appear in the initial substances in the form of HD and D 2 molecules. Another spectacular experiment was carried out at the Faculty of Chemistry of Moscow State University. The silver plate was placed in a solution of silver nitrate, and no changes were observed. Then an insignificant amount of radioactive silver ions was introduced into the solution, after which the silver plate became radioactive. This radioactivity could not be "washed away" either by rinsing the plate with water or by washing it with hydrochloric acid. Only etching with nitric acid or mechanical processing of the surface with fine sandpaper made it inactive. There is only one way to explain this experiment: there is a continuous exchange of silver atoms between the metal and the solution, i.e. in the system there is a reversible reaction Ag (tv) - e - \u003d Ag +. Therefore, the addition of radioactive ions Ag + to the solution led to their "embedding" into the plate in the form of electrically neutral, but still radioactive atoms. Thus, not only chemical reactions between gases or solutions are in equilibrium, but also the processes of dissolution of metals and precipitation. For example, a solid dissolves fastest when placed in a pure solvent when the system is far from equilibrium, in this case- from a saturated solution. Gradually, the dissolution rate decreases, and at the same time the rate of the reverse process increases - the transition of a substance from solution to a crystalline precipitate. When the solution becomes saturated, the system reaches a state of equilibrium, while the dissolution and crystallization rates are equal, and the mass of the precipitate does not change with time. How can the system "counteract" changes in external conditions? If, for example, the temperature of the equilibrium mixture is increased by heating, the system itself, of course, cannot “weaken” external heating, but the equilibrium in it is shifted in such a way that heating the reaction system to a certain temperature requires more heat than in the case unless the balance shifted. In this case, the equilibrium is shifted so that the heat is absorbed, i.e. towards an endothermic reaction. This can be interpreted as "the desire of the system to weaken external influences." On the other hand, if there is an unequal number of gaseous molecules on the left and right sides of the equation, then the equilibrium in such a system can also be shifted by changing the pressure. With increasing pressure, the equilibrium shifts to the side where the number of gaseous molecules is less (and in this way, as it were, “opposes” external pressure). If the number of gaseous molecules does not change during the reaction
(H 2 + Br 2 (g) 2HBr, CO + H 2 O (g) CO 2 + H 2), then the pressure does not affect the equilibrium position. It should be noted that when the temperature changes, the equilibrium constant of the reaction also changes, while when only the pressure changes, it remains constant.
Several examples of the use of Le Chatelier's principle for predicting shifts in chemical equilibrium. The reaction 2SO 2 + O 2 2SO 3 (d) is exothermic. If the temperature is raised, the endothermic decomposition of SO 3 will take precedence and the equilibrium will shift to the left. If the temperature is lowered, the equilibrium will shift to the right. So, a mixture of SO 2 and O 2, taken in a stoichiometric ratio of 2: 1 ( cm . stoichiomerism), at a temperature of 400 ° C and atmospheric pressure turns into SO 3 with a yield of about 95%, i.e. the state of equilibrium under these conditions is almost completely shifted towards SO 3 . At 600°C, the equilibrium mixture already contains 76% SO 3 , and at 800°C, only 25%. That is why when sulfur is burned in air, mainly SO 2 and only about 4% SO 3 are formed. It also follows from the reaction equation that an increase in the total pressure in the system will shift the equilibrium to the right, and with a decrease in pressure, the equilibrium will shift to the left.
The reaction of abstraction of hydrogen from cyclohexane with the formation of benzene
C 6 H 12 C 6 H 6 + 3H 2 is carried out in the gas phase, also in the presence of a catalyst. This reaction goes with the expenditure of energy (endothermic), but with an increase in the number of molecules. Therefore, the effect of temperature and pressure on it will be directly opposite to that observed in the case of ammonia synthesis. Namely: an increase in the equilibrium concentration of benzene in the mixture is facilitated by an increase in temperature and a decrease in pressure, so the reaction is carried out in industry at low pressures (2–3 atm) and high temperatures (450–500 ° C). Here, an increase in temperature is “doubly favorable”: it not only increases the reaction rate, but also contributes to a shift in the equilibrium towards the formation of the target product. Of course, an even greater decrease in pressure (for example, to 0.1 atm) would cause a further shift of the equilibrium to the right, however, in this case, there will be too little substance in the reactor, and the reaction rate will also decrease, so that the overall productivity will not increase, but will decrease. This example once again shows that an economically justified industrial synthesis is a successful maneuvering between Scylla and Charybdis.
Le Chatelier's principle "works" in the so-called halogen cycle, which is used to produce titanium, nickel, hafnium, vanadium, niobium, tantalum and other high purity metals. The reaction of a metal with a halogen, for example, Ti + 2I 2 TiI 4, proceeds with the release of heat, and therefore, as the temperature rises, the equilibrium shifts to the left. Thus, at 600°C, titanium easily forms volatile iodide (the equilibrium is shifted to the right), and at 110°C, the iodide decomposes (the equilibrium is shifted to the left) with the release of a very pure metal. Such a cycle also works in halogen lamps, where tungsten evaporated from the spiral and settled on colder walls forms volatile compounds with halogens, which decompose again on a hot spiral, and tungsten is transferred to its original place.
In addition to changing temperature and pressure, there is another effective way to influence the equilibrium position. Imagine that from an equilibrium mixture
A + B C + D any substance is excreted. In accordance with Le Chatelier's principle, the system will immediately "respond" to such an impact: the equilibrium will begin to shift in such a way as to compensate for the loss of a given substance. For example, if substance C or D (or both at once) is removed from the reaction zone, the equilibrium will shift to the right, and if substances A or B are removed, it will shift to the left. The introduction of any substance into the system will also shift the equilibrium, but in the other direction.
Substances can be removed from the reaction zone different ways. For example, if there is sulfur dioxide in a tightly closed vessel with water, an equilibrium will be established between gaseous, dissolved and reacted sulfur dioxide:
O 2 (g) SO 2 (p) + H 2 O H 2 SO 3. If the vessel is opened, sulfur dioxide will gradually begin to evaporate and will no longer be able to participate in the process - the equilibrium will begin to shift to the left, until the complete decomposition of sulfurous acid. A similar process can be observed every time you open a bottle of lemonade or mineral water: the balance of CO 2 (g) CO 2 (p) + H 2 O H 2 CO 3 shifts to the left as CO 2 volatilizes.
The removal of a reagent from the system is possible not only with the formation of gaseous substances, but also by binding one or another reagent with the formation of an insoluble compound that precipitates. For example, if an excess of calcium salt is introduced into an aqueous solution of CO 2, then Ca 2+ ions will form a precipitate of CaCO 3, reacting with carbonic acid; the equilibrium CO 2 (p) + H 2 OH 2 CO 3 will shift to the right until there is no dissolved gas left in the water.
The equilibrium can also be shifted by adding a reagent. So, when dilute solutions of FeCl 3 and KSCN are drained, a reddish-orange color appears as a result of the formation of iron thiocyanate (thiocyanate):
FeCl 3 + 3KSCN Fe(SCN) 3 + 3KCl. If additional FeCl 3 or KSCN is added to the solution, the color of the solution will increase, which indicates a shift of the equilibrium to the right (as if weakening the external influence). If, however, an excess of KCl is added to the solution, then the equilibrium will shift to the left with a decrease in color to light yellow.
In the formulation of Le Chatelier's principle, it is not for nothing that it is indicated that it is possible to predict the results of external influence only for systems that are in equilibrium. If this indication is neglected, it is easy to come to completely wrong conclusions. For example, it is known that solid alkalis (KOH, NaOH) dissolve in water with the release of a large amount of heat - the solution heats up almost as much as when concentrated sulfuric acid is mixed with water. If we forget that the principle applies only to equilibrium systems, we can make the wrong conclusion that as the temperature rises, the solubility of KOH in water should decrease, since it is precisely this shift in the equilibrium between the precipitate and the saturated solution that leads to "weakening of the external influence." However, the process of dissolving KOH in water is not at all equilibrium, since anhydrous alkali is involved in it, while the precipitate that is in equilibrium with a saturated solution is KOH hydrates (mainly KOH 2H 2 O). The transition of this hydrate from the precipitate to the solution is an endothermic process, i.e. is accompanied not by heating, but by cooling of the solution, so that Le Chatelier's principle for an equilibrium process is also fulfilled in this case. In the same way, when anhydrous salts - CaCl 2, CuSO 4, etc. are dissolved in water, the solution heats up, and when crystalline hydrates CuSO 4 5H 2 O, CaCl 2 6H 2 O are dissolved, it cools.
Another interesting and instructive example of the misuse of Le Chatelier's principle can be found in textbooks and popular literature. If an equilibrium mixture of brown nitrogen dioxide NO 2 and colorless N 2 O 4 tetroxide is placed in a transparent gas syringe, and then the gas is quickly compressed with a piston, the color intensity will immediately increase, and after a while (tens of seconds) it will weaken again, although will not reach the original. This experience is usually explained as follows. The rapid compression of the mixture results in an increase in pressure and therefore in the concentration of both components, so the mixture becomes darker. But an increase in pressure, in accordance with the principle of Le Chatelier, shifts the equilibrium in the 2NO 2 N 2 O 4 system towards colorless N 2 O 4 (the number of molecules decreases), so the mixture gradually brightens, approaching a new equilibrium position, which corresponds to increased pressure.
The fallacy of this explanation follows from the fact that both reactions - the dissociation of N 2 O 4 and the dimerization of NO 2 - occur extremely quickly, so that the equilibrium is established in millionths of a second anyway, so it is impossible to push the piston so fast as to disturb the equilibrium. This experience is explained differently: gas compression causes a significant increase in temperature (everyone who has had to inflate a tire with a bicycle pump is familiar with this phenomenon). And in accordance with the same principle of Le Chatelier, the equilibrium instantly shifts towards an endothermic reaction that goes with the absorption of heat, i.e. towards the dissociation of N 2 O 4 - the mixture darkens. Then the gases in the syringe slowly cool down to room temperature, and the equilibrium shifts again towards the tetroxide - the mixture becomes brighter.
Le Chatelier's principle works well in cases that have nothing to do with chemistry. In a normally functioning economy, the total amount of money in circulation is in equilibrium with the goods that this money can buy. What happens if the “outside influence” is the desire of the government to print more money to pay off debts? In strict accordance with Le Chatelier's principle, the balance between commodity and money will be shifted in such a way as to weaken the citizens' pleasure from having more money. Namely, the prices of goods and services will rise, and in this way a new equilibrium will be reached. Another example. In one of the US cities, it was decided to get rid of constant traffic jams by expanding highways and building interchanges. This helped for a while, but then elated residents started buying more cars, so that traffic jams soon reappeared—but with a new “balance position” between roads and more cars.
So, we will draw the main conclusions about the methods of shifting the chemical equilibrium.
Le Chatelier's principle. If an external influence is made on a system in equilibrium (concentration, temperature, pressure change), then it favors the flow of one of the two opposite reactions that weakens this effect.
|
V 1 | ||
|
A+B |
|
IN |
|
V 2 |
1. Pressure. An increase in pressure (for gases) shifts the equilibrium towards a reaction leading to a decrease in volume (i.e., to the formation of a smaller number of molecules).
2. An increase in temperature shifts the equilibrium position towards an endothermic reaction (i.e. towards a reaction proceeding with the absorption of heat)
3. An increase in the concentration of starting substances and the removal of products from the reaction sphere shifts the equilibrium towards a direct reaction. Increasing the concentrations of the starting materials [A] or [B] or [A] and [B]: V 1 > V 2 .
Catalysts do not affect the equilibrium position.
Le Chatelier's principle in nature.
When studying this topic, I always want to give an example of the desire of all living things for balance, compensation. For example: change in mouse population - nut year - there is a lot of food for mice, the population of mice is growing rapidly. With an increase in the number of mice, the amount of food decreases, as a result of the accumulation of rodents, the growth of various infectious diseases among mice begins, so there is a gradual decrease in the population of rodents. After a certain period of time, a dynamic equilibrium in the number of born and dying mice sets in, a shift in this balance can occur in one direction or another under the influence of external, favorable or unfavorable conditions.
Biochemical processes take place in the human body, which can also be regulated according to the principle of Le Chatelier. Sometimes, as a result of such a reaction, poisonous substances begin to be produced in the body, causing a particular disease. How to prevent this process?
Let's remember such a method of treatment as homeopathy. The method consists in the use of very small doses of those drugs that, in large doses, cause healthy person signs of some disease. How does the drug-poison work in this case? The product of an undesired reaction is introduced into the body, and according to Le Chatelier's principle, the equilibrium is shifted towards the starting substances. The process that causes painful disorders in the body is extinguished.
Practical part.
Control of the level of assimilation of the studied topic is carried out in the form of tests. A test system of concisely and precisely formulated and standardized tasks, some of which must be given within a limited time, short and precise answers, evaluated by a scoring system. When compiling tests, I focused on the following levels:
Reproductive-performance by students of this level occurs mainly based on memory.
Productive achievement of this level requires students to understand the studied formulations, concepts, laws, the ability to establish the relationship between them.
Creative - the ability to predict based on existing knowledge, design, analyze, draw conclusions, comparisons, generalizations.
Tests closed type or tests in which the subject must choose the correct answer from the options provided.
A) Reproductive level: tests with alternative answers, in which the subject must answer yes or no. Score 1 point.
The combustion reaction of phosphorus-
a) yes b) no
decomposition reaction
reversible reaction
a) yes b) no
Temperature increase
mercury oxide II for mercury
and oxygen
a) yes b) no
In living systems
and irreversible processes
a) yes b) no.
Multiple Choice Tests
In which system will the chemical equilibrium shift to the right when the pressure is increased?
2HI(g)↔H2(g)+I2(g)
C (tv) + S2 (g) ↔CS2 (g)
C3H6(g)+H2(g)↔С3H8(g)
H2(g)+F2(g)↔2HF(g) 1 point
rise in temperature
using a catalyst
lowering the temperature; 1 point
On the state of chemical equilibrium in the system
does not affect
increase in pressure
increase in iodine concentration
temperature increase
decrease in temperature; 1 point
In which system does an increase in hydrogen concentration shift the chemical equilibrium to the left?
C(tv)+2H2(g)↔СH4(g)
2NH3(g)↔N2(g)+3H2(g)
2H2(g)+O2(g)↔2H2O(g)
FeO(solid)+H2(g)↔Fe+H2O(g) 1 point
In which system does an increase in pressure not affect the shift in chemical equilibrium?
H2(g)+J2(g)↔2HJ(g)
SO2(g)+H2O(l)↔H2SO3(g)
CH4(g)+H2O(g)↔CO(g)+3H2(g)
4HCl(g)+O2(g)↔2H2O(g)+2Сl2(g) 1 point
On the chemical equilibrium in the system
has no effect
temperature increase
pressure increase
removal of ammonia from the reaction zone
application of a catalyst 1 point
Chemical equilibrium in the system
shifts towards the formation of the reaction product at
increase in pressure
rise in temperature
pressure drop
application of a catalyst 1 point
In the production of sulfuric acid at the stage of oxidation of SO2 to SO3 to increase the yield of the product
increase the concentration of oxygen
increase the temperature
lower blood pressure
introducing a catalyst; 1.5 points
Alkene + H2 ↔ alkane
when the reaction mixture is heated
balance will shift to the right
balance will shift to the left
equilibrium will flow in both directions with the same probability
these substances are not in equilibrium under the specified conditions; 1.5 points
Chemical equilibrium in the system
shifts towards the formation of starting materials
1) pressure increase
rise in temperature
pressure drop
the use of a catalyst; 1 point
To shift the equilibrium to the right in the system
has an impact
temperature drop
pressure increase
use of a catalyst
temperature increase; 1 point
An irreversible reaction corresponds to the equation
nitrogen + hydrogen = ammonia
acetylene + oxygen = carbon dioxide + water
hydrogen + iodine = hydrogen iodide
sulfur dioxide + oxygen = sulfuric anhydride; 1.5 points
Multiple Choice Tests, during which the subject must choose 1-2 correct answers, or match 2 proposed conditions when choosing an answer.
In which system will the chemical equilibrium shift towards the products of the reaction, both with an increase in pressure, as well as with a decrease in temperature?
N2+O2↔2NO-Q
N2+3H2↔2NH3+Q
H2+CL2↔2HCL+Q
C2H2↔2C(tv)+H2-Q 1.5 points
Chemical equilibrium in the system
NH3+H2O↔NH4+OH
will shift towards the formation of ammonia when ammonia is added to an aqueous solution
sodium chloride
sodium hydroxide
of hydrochloric acid
aluminum chloride; 1.5 points
19) The ethylene hydration reaction CH2=CH2+H2O ↔ has a large practical value, but it is reversible, to shift the reaction equilibrium to the right, it is necessary
raise the temperature (>280 degrees C)
reduce the amount of water in the reaction mixture
increase pressure (more than 80 atmospheres)
replace the acid catalyst with platinum; 1 point
The dehydrogenation reaction of butane is endothermic. To shift the reaction equilibrium to the right,
use a more active catalyst, such as platinum
lower the temperature
raise the pressure
raise the temperature 1 point
For the reaction of the interaction of acetic acid with methanol with the formation of ether and water, the equilibrium shift to the left will be promoted by
appropriate catalyst
adding concentrated sulfuric acid
use of dehydrated starting materials
adding ether; 1.5 points
Exclusion tests
The balance shift is affected
pressure change
use of a catalyst
change in the concentrations of substances involved in the reaction
temperature change; 1 point
An increase or decrease in pressure affects the shift in chemical equilibrium in reactions
going with the release of heat
reactions involving gaseous substances
reactions proceeding with a decrease in volume
reactions going with an increase in volume; 1.5 points
The reaction is irreversible
burning coal
burning phosphorus
synthesis of ammonia from nitrogen and hydrogen
burning methane; 1.5 points
Grouping Tests include a list of proposed formulas, equations, terms that should be distributed according to given criteria
With a simultaneous increase in temperature and decrease in pressure, the chemical equilibrium will shift to the right in the system
H2(g)+S(g)↔H2S(g)+Q
2SO2(g)+O2(g)↔2SO3(g)+Q
2NH3(g)↔N2(g)+3H2(g)-Q
2HCL(g)↔H2(g)+CL2(g)-Q; 2 points
The propene hydrogenation reaction is exothermic. To shift the chemical equilibrium to the right, it is necessary
temperature drop
increase in pressure
decrease in hydrogen concentration
decrease in the concentration of propene; 1 point
Compliance tasks.
When performing tests, the subject is asked to match the elements of two lists, with several possible answers.
The equilibrium of the reaction shifts to the right. Bring in line.
B) N2+3H2↔2NH3+Q 2) When the temperature rises
C) CO2 + C (solid) ↔2CO-Q 3) When the pressure drops
D) N2O(g)+S(t)↔2N2(g) 4) With an increase in the contact area; 2 points
The equilibrium of the reaction is shifted towards the formation of reaction products. Bring in line.
B) 2H2 + O2 ↔ 2H2O (g) + Q 2) With increasing temperature
C) CH3OH + CH3COOH↔CH3COOCH3 3) When the pressure decreases
D) N2+O2↔2NO-Q 4) When adding ether
5) When adding alcohol; 2 points
Open-ended or open-ended tests, in which the subject needs to add the concepts of the definition of the equation or offer an independent judgment in evidence.
Tasks of this type constitute the final, most highly valued part USE tests in chemistry.
Supplement tasks.
The subject must formulate answers, taking into account the restrictions provided for in the task.
Add the reaction equation related to reversible and exothermic at the same time
B) Hydrogen + Iodine
C) Nitrogen + Hydrogen
D) Sulfur dioxide + Oxygen
E) Carbon dioxide + Carbon 2 points
Write the reaction equation according to the scheme, from which select those reversible reactions in which an increase in temperature will cause the equilibrium to shift to the right:
N2 → NO→ NO2→ HNO3→ NH4NO3 2 points
Free presentation tests.
The subject must independently formulate the answers, because no restrictions are imposed on them in the task.
31) List the factors that shift the equilibrium to the right in the system:
CO + 2H2↔ CH3OH(g)+Q 2 points
32) List the factors that shift the equilibrium towards the formation of starting substances in the system:
C (tv) + 2H2 (g) ↔CH4 (g) + Q 2 points
Answers to tests.
Test No. Correct answer
B-1
G-3.4
A-2.3
G-2
B- N2+3H2↔2NH3+Q
1) N2+O2↔2NO-Q
3) 4NO2+2H2O+O2↔4HNO3+Q
4) NH3+HNO3=NH4NO3
first reaction
CO+2H2↔CH3OH+Q
decrease in temperature
increase in pressure
increasing the concentration of CO
increase in H2 concentration
decrease in alcohol concentration
C+2H2↔CH4+Q
2) pressure reduction
3) lowering the concentration of hydrogen
4) increase in methane concentration.
Bibliography
Akhmetov, M.A. The system of tasks and exercises in organic chemistry in test form [Text] / M.A. Akhmetov, I.N. Prokhorov.-Ulyanovsk: IPKPRO, 2004.
Gabrielyan, O.S. Modern didactics of school chemistry, lecture No. 6 [Text] / O.S. Gabrielyan, V.G. Krasnova, S.T. Sladkov.// Newspaper for teachers of chemistry and natural sciences ( Publishing House"First of September") -2007.- No. 22.-p.4-13.
Kaverina, A.A. Educational and training materials for preparing for the unified state exam. Chemistry [Text] / A.A. Kaverina et al. - M .: Intellect Center, 2004.-160s.
Kaverina, A.A. Unified State Exam 2009. Chemistry [Text] / A.A. Kaverina, A.S. Koroshchenko, D.Yu. Dobrotin / FIPI.-M .: Intellect Center, 2009.-272 p.
Leenson, I.A. Chemical reactions, thermal effect, equilibrium, speed [Text] / I. A. Leenson. M .: Astrel, 2002.-190s.
Radetsky, A.M. Verification work in chemistry in grades 8-11: a guide for the teacher [Text] / A.M. Radetsky. M.: Enlightenment, 2009.-272p.
Ryabinina, O.A. Demonstration of the principle of Le Chatelier [Text] / O. O. Ryabinina, A. Illarionov / / Chemistry at school. -2008. - No. 7. - p. 64-67.
Tushina.E.N. The principle of Le Chatelier and some methods of treatment [Text] / E.N. Tushina.// Chemistry at school.-1993. No. 2.-p.54.
Shelinskiy, G.I. Fundamentals of the theory of chemical processes [Text] / G.I. Shelinskiy. M.: Enlightenment, 1989.-234p.
Strempler, G.I. Pre-profile training in chemistry [Text]
>> Chemistry: Reversible and irreversible reactions
CO2 + H2O = H2CO3
Leave the resulting acid solution to stand in a tripod. After a while, we will see that the solution has turned purple again, as the acid has decomposed into its original substances.
This process can be carried out much faster if a third is a solution of carbonic acid. Consequently, the reaction of obtaining carbonic acid proceeds both in the forward and in the opposite direction, that is, it is reversible. The reversibility of a reaction is indicated by two oppositely directed arrows:
Among the reversible reactions underlying the preparation of the most important chemical products, we mention as an example the reaction of synthesizing (compounding) sulfur oxide (VI) from sulfur oxide (IV) and oxygen.
1. Reversible and irreversible reactions.
2. Berthollet's rule.
Write down the equations for the combustion reactions that were mentioned in the text of the paragraph, revealing that as a result of these reactions, oxides of those elements are formed from which the initial substances are built.
Give a description of the last three reactions carried out at the end of the paragraph, according to the plan: a) the nature and number of reagents and products; b) state of aggregation; c) direction: d) the presence of a catalyst; e) release or absorption of heat
What inaccuracy is made in the equation for the reaction of limestone roasting proposed in the text of the paragraph?
How true is the statement that the reactions of the compound will, as a rule, be exothermic reactions? Justify your point of view using the facts given in the text of the textbook.
Lesson content lesson summary support frame lesson presentation accelerative methods interactive technologies Practice tasks and exercises self-examination workshops, trainings, cases, quests homework discussion questions rhetorical questions from students Illustrations audio, video clips and multimedia photographs, pictures graphics, tables, schemes humor, anecdotes, jokes, comics parables, sayings, crossword puzzles, quotes Add-ons abstracts articles chips for inquisitive cheat sheets textbooks basic and additional glossary of terms other Improving textbooks and lessonscorrecting errors in the textbook updating a fragment in the textbook elements of innovation in the lesson replacing obsolete knowledge with new ones Only for teachers perfect lessons calendar plan year methodological recommendations of the discussion program Integrated LessonsCodifier Topics: reversible and irreversible reactions. chemical balance. Displacement of chemical equilibrium under the influence of various factors.
According to the possibility of a reverse reaction, chemical reactions are divided into reversible and irreversible.
Reversible chemical reactions are reactions whose products, under given conditions, can interact with each other.
irreversible reactions These are reactions whose products under given conditions cannot interact with each other.
More details about classification of chemical reactions can be read.
The probability of product interaction depends on the conditions of the process.
So if the system open, i.e. exchanges with environment both matter and energy, then chemical reactions in which, for example, gases are formed, will be irreversible. For example , when calcining solid sodium bicarbonate:
2NaHCO 3 → Na 2 CO 3 + CO 2 + H 2 O
gaseous carbon dioxide will be released and volatilize from the reaction zone. Therefore, such a reaction will irreversible under these conditions. If we consider closed system , which can not exchange matter with the environment (for example, a closed box in which the reaction takes place), then carbon dioxide will not be able to escape from the reaction zone, and will interact with water and sodium carbonate, then the reaction will be reversible under these conditions:
2NaHCO 3 ⇔ Na 2 CO 3 + CO 2 + H 2 O
Consider reversible reactions. Let the reversible reaction proceed according to the scheme:
aA + bB = cC + dD
The rate of the forward reaction according to the law of mass action is determined by the expression: v 1 =k 1 ·C A a ·C B b , the rate of the reverse reaction: v 2 =k 2 ·C C c ·C D d . If at the initial moment of the reaction there are no substances C and D in the system, then particles A and B mainly collide and interact, and a predominantly direct reaction occurs. Gradually, the concentration of particles C and D will also begin to increase, therefore, the rate of the reverse reaction will increase. At some point the rate of the forward reaction becomes equal to the rate of the reverse reaction. This state is called chemical equilibrium .
Thus, chemical equilibrium is the state of the system in which the rates of the forward and reverse reactions are equal .
Because the rates of the forward and reverse reactions are equal, the rate of formation of substances is equal to the rate of their consumption, and the current concentrations of substances do not change . Such concentrations are called balanced .
Note that in equilibrium both forward and reverse reactions, that is, the reactants interact with each other, but the products also interact at the same rate. At the same time, external factors may influence shift chemical equilibrium in one direction or another. Therefore, chemical equilibrium is called mobile, or dynamic.
Research in the field of moving balance began in the 19th century. In the writings of Henri Le Chatelier, the foundations of the theory were laid, which were later generalized by the scientist Karl Brown. The principle of moving balance, or the principle of Le Chatelier-Brown, states:
If a system in equilibrium is subjected to external factor, which changes any of the equilibrium conditions, then processes in the system are intensified, aimed at compensating for external influences.
In other words: under an external influence on the system, the equilibrium will shift in such a way as to compensate for this external influence.
This principle, which is very important, works for any equilibrium phenomena (not just chemical reactions). However, we will now consider it in relation to chemical interactions. In the case of chemical reactions, external action leads to a change in the equilibrium concentrations of substances.
Three main factors can affect chemical reactions at equilibrium: temperature, pressure, and concentrations of reactants or products.
1. As you know, chemical reactions are accompanied by a thermal effect. If the direct reaction proceeds with the release of heat (exothermic, or + Q), then the reverse reaction proceeds with the absorption of heat (endothermic, or -Q), and vice versa. If you raise temperature in the system, the equilibrium will shift so as to compensate for this increase. It is logical that with an exothermic reaction, the temperature increase cannot be compensated. Thus, as the temperature rises, the equilibrium in the system shifts towards heat absorption, i.e. towards endothermic reactions (-Q); with decreasing temperature - in the direction of an exothermic reaction (+ Q).
2. In the case of equilibrium reactions, when at least one of the substances is in the gas phase, the equilibrium is also significantly affected by the change pressure in system. When the pressure is increased, the chemical system tries to compensate for this effect, and increases the rate of the reaction, in which the amount of gaseous substances decreases. When the pressure is reduced, the system increases the rate of the reaction, in which more molecules of gaseous substances are formed. Thus: with an increase in pressure, the equilibrium shifts towards a decrease in the number of gas molecules, with a decrease in pressure - towards an increase in the number of gas molecules.
Note! Systems where the number of molecules of reactant gases and products is the same are not affected by pressure! Also, a change in pressure practically does not affect the equilibrium in solutions, i.e. in reactions where there are no gases.
3. Also, the equilibrium in chemical systems is affected by the change concentration reactants and products. As the concentration of the reactants increases, the system tries to use them up and increases the rate of the forward reaction. With a decrease in the concentration of reagents, the system tries to accumulate them, and the rate of the reverse reaction increases. With an increase in the concentration of products, the system also tries to use them up, and increases the rate of the reverse reaction. With a decrease in the concentration of products, the chemical system increases the rate of their formation, i.e. the rate of the forward reaction.
If in a chemical system the rate of the forward reaction increases right , towards the formation of products And reagent consumption . If the rate of the reverse reaction increases, we say that the balance has shifted to the left , towards food consumption And increasing the concentration of reagents .
For example, in the ammonia synthesis reaction:
N 2 + 3H 2 \u003d 2NH 3 + Q
an increase in pressure leads to an increase in the reaction rate, in which a smaller number of gas molecules are formed, i.e. direct reaction (the number of reactant gas molecules is 4, the number of gas molecules in the products is 2). As the pressure increases, the equilibrium shifts to the right, towards the products. At rise in temperature balance will shift towards an endothermic reaction, i.e. to the left, towards the reagents. An increase in the concentration of nitrogen or hydrogen will shift the equilibrium towards their consumption, i.e. to the right, towards the products.
Catalyst does not affect the balance, because speeds up both the forward and reverse reactions.
One of the most important characteristics a chemical reaction is the depth (degree) of transformation, showing how much the starting substances are converted into reaction products. The larger it is, the more economically the process can be carried out. The depth of conversion, among other factors, depends on the reversibility of the reaction.
reversible reactions , Unlike irreversible, do not proceed to the end: none of the reactants is completely consumed. At the same time, the reaction products interact with the formation of starting materials.
Consider examples:
1) equal volumes of gaseous iodine and hydrogen are introduced into a closed vessel at a certain temperature. If the collisions of the molecules of these substances occur with the desired orientation and sufficient energy, then the chemical bonds can be rearranged with the formation of an intermediate compound (an activated complex, see section 1.3.1). Further rearrangement of bonds can lead to the decomposition of the intermediate compound into two molecules of hydrogen iodide. Reaction equation:
H 2 + I 2 ® 2HI
But the molecules of hydrogen iodide will also randomly collide with molecules of hydrogen, iodine and among themselves. When HI molecules collide, nothing will prevent the formation of an intermediate compound, which can then decompose into iodine and hydrogen. This process is expressed by the equation:
2HI ® H 2 + I 2
Thus, two reactions will proceed simultaneously in this system - the formation of hydrogen iodide and its decomposition. They can be expressed by one general equation
H 2 + I 2 "2HI
The reversibility of the process is shown by the sign “.
The reaction directed in this case towards the formation of hydrogen iodide is called direct, and the opposite is called reverse.
2) if we mix two moles of sulfur dioxide with one mole of oxygen, create conditions in the system that are favorable for the reaction to proceed, and after the time has elapsed, analyze the gas mixture, the results will show that the system will contain both SO 3 - the reaction product, and the initial substances - SO 2 and O 2. If sulfur oxide (+6) is placed under the same conditions as the initial substance, then it will be possible to find that part of it will decompose into oxygen and sulfur oxide (+4), and the final ratio between the amounts of all three substances will be the same as when starting from a mixture of sulfur dioxide and oxygen.
Thus, the interaction of sulfur dioxide with oxygen is also one of the examples of a reversible chemical reaction and is expressed by the equation
2SO 2 + O 2 "2SO 3
3) the interaction of iron with hydrochloric acid proceeds according to the equation:
Fe + 2HCL ® FeCL 2 + H 2
With enough hydrochloric acid, the reaction will end when
all the iron is used up. In addition, if you try to carry out this reaction in the opposite direction - to pass hydrogen through a solution of iron chloride, then metallic iron and hydrochloric acid will not work - this reaction cannot go in the opposite direction. Thus, the interaction of iron with hydrochloric acid is an irreversible reaction.
However, it should be borne in mind that theoretically any irreversible process can be represented as reversible under certain conditions, i.e. In principle, all reactions can be considered reversible. But very often one of the reactions clearly prevails. This happens in those cases when the products of interaction are removed from the reaction sphere: a precipitate precipitates, a gas is released, during ion-exchange reactions practically non-dissociating products are formed; or when, due to a clear excess of starting substances, the opposite process is practically suppressed. Thus, the natural or artificial exclusion of the possibility of a reverse reaction allows you to bring the process almost to the end.
Examples of such reactions are the interaction of sodium chloride with silver nitrate in solution
NaCL + AgNO 3 ® AgCl¯ + NaNO 3 ,
copper bromide with ammonia
CuBr 2 + 4NH 3 ® Br 2,
neutralization of hydrochloric acid with sodium hydroxide solution
HCl + NaOH ® NaCl + H 2 O.
These are all examples only practically irreversible processes, since silver chloride is somewhat soluble, and the complex cation 2+ is not absolutely stable, and water dissociates, although to an extremely small extent.
