Reversible and irreversible reactions, chemical equilibrium. Reversible and irreversible reactions – Knowledge Hypermarket

retraining of education workers.

Topic: “Reversible and irreversible reactions.

Chemical balance. Le Chatelier's principle.

Work completed:

Group X listener – 1

Chemistry teacher Municipal Educational Institution Secondary School No. 6

Dimitrovgrad

Ulyanovsk region

Lepikhova Tatyana Vasilievna.

Scientific adviser:

Head of the department

natural sciences

Akhmetov Marat Anvarovich

Reversible and irreversible chemical reactions.

Chemical balance.

Le Chatelier's principle.

2) Generalization and concretization of knowledge about the laws of chemical reactions, the formation of skills 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 ideas about chemical equilibrium as a special case of the unified law of natural equilibrium, the desire for compensation, the stability of equilibrium in unity with the basic form of existence of matter, movement, dynamics.

1. 
   Consider the topic: “Reversible and irreversible reactions” using specific examples, using previous ideas about the rate of chemical reactions.
2. 
   Continue studying the features of reversible chemical reactions and developing ideas about chemical equilibrium as a dynamic state of a reacting system.
3. 
   Study the principles of shifting chemical equilibrium and teach students to determine the conditions for shifting chemical equilibrium.
4. 
   To give students an idea of ​​the significance 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 at various levels: domestic, defense, industrial, technical, environmental and others, thousands, millions of completely different reactions occur or are carried out, which can be viewed from different perspectives. points of view and classifications. We will consider chemical reactions from the point of view of their reversibility and irreversibility.

It is difficult to overestimate the importance of these concepts: as long as a thinking person exists, the human thought about the reversibility and irreversibility of the processes occurring in his body, the eternal problem of prolonging a person’s life, the problem of the irreversibility of the consequences of his life activity, a thoughtless attitude towards 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. Present a theoretical basis with subsequent testing, self-testing of knowledge on this topic, using tests of various typologies. I assume that by “traversing the path” from simple to more complex tasks, students will have clear, good knowledge not only on this topic, but will also deepen their knowledge of chemistry.

Chemical reactions are phenomena in which one (or some) substances are transformed into others, evidence of this is visible and invisible changes. Visible: changes in color, odor, taste, precipitation, change in indicator color, absorption and release of heat. Invisible: changes 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.

An example of such a reaction is the decomposition of potassium chlorate (Bertholette salt) when heated:

2KClO 3 = 2KCl + 3O 2

The reaction will stop when all the potassium chlorate is converted into potassium chloride and oxygen. There are not many irreversible reactions.

If acid and alkali solutions are combined, salt and water are formed, for example,

HCl + NaOH = NaCl + H 2 O, and if the substances were taken in the required 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 resulting substances - sodium chloride and water, then no changes will be found. In such cases, they say that the reaction of an acid with an alkali is irreversible, i.e. there is no backlash. Many reactions are practically irreversible at room temperature, for example,

H 2 + Cl 2 = 2HCl, 2H 2 + O 2 = 2H 2 O, etc.

Reversible reactions. Reversible reactions are those that simultaneously occur in two mutually opposite directions.

Most reactions are reversible. In equations of reversible reactions, two arrows pointing in opposite directions are placed between the left and right sides. An example of such a reaction is the synthesis of ammonia from hydrogen and nitrogen:

,

∆H = -46.2 kJ/mol

In technology, reversible reactions are usually disadvantageous. Therefore, various methods (changes in temperature, pressure, etc.) make them practically irreversible.

Irreversible reactions are those reactions that occur:

1) the resulting products leave the reaction sphere - they precipitate, are released in the form of gas, for example

BaCl 2 + H 2 SO 4 = BaSO 4 ↓ + 2HCl

Na 2 CO 3 + 2HCl = 2NaCl + CO 2 ↓ + H 2 O

2) a slightly dissociated compound is formed, for example water:

HCl + NaOH = 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 = MgO, ∆H = -602.5 kJ / mol

In equations of irreversible reactions, an equal sign or an arrow is placed between the left and right sides.

Many reactions are reversible even under normal conditions, which means that the reverse reaction occurs to a noticeable extent. For example, if you try to neutralize an aqueous solution of very weak hypochlorous acid with an alkali, it turns out that the neutralization reaction does not proceed to completion 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 transform into the original compounds. As a result, the solution has an alkaline reaction. The reaction for the 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. Typically, a reaction is considered irreversible if, after completion, the concentrations of the starting substances are so low that they cannot be detected (of course, this depends on the sensitivity of the analytical methods). When external conditions change (primarily temperature and pressure), an irreversible reaction can become reversible and vice versa. Thus, at atmospheric pressure and temperatures below 1000 ° C, the reaction 2H 2 + O 2 = 2H 2 O can still be considered irreversible, while at a temperature of 2500 ° C and above water dissociates into hydrogen and oxygen by approximately 4%, and at a temperature of 3000 °C – 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 iodide: H 2 + I 2 2HI. By changing the temperature, he could achieve preferential occurrence of only the forward or only the reverse reaction, but in the general case, both reactions proceeded simultaneously in opposite directions. There are many similar examples. One of the most famous is the reaction of ammonia synthesis 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 the gas and high concentrations of nitrogen and hydrogen, ammonia is formed; on the contrary, at a high concentration of ammonia it decomposes, the reaction proceeds in the opposite direction. Upon completion of a reversible reaction, i.e., upon reaching chemical equilibrium, the system contains both starting materials and reaction products. A reaction is called irreversible if it can occur only in one direction and ends with the complete conversion of the starting substances into products; an example is the decomposition of explosives. The same reaction, depending on the conditions (temperature, pressure), can be significantly reversible or practically irreversible. A simple (one-stage) reversible reaction consists of two elementary reactions occurring simultaneously, which differ from each other only in the direction of the chemical transformation. The direction of the final reaction accessible to direct observation is determined by which of these mutually inverse reactions has a higher speed. For example, the simple reaction N2O4 Û 2NO2 (2) consists of the elementary reactions N2O4 ? 2NO2 and 2NO2 ? N2O4. For the reversibility of a complex (multistage) reaction, for example reaction (1), it is necessary that all its constituent stages are reversible.? M. I. Tyomkin.

CHEMICAL EQUILIBRIUM.

Chemical equilibrium- 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 is dynamic in nature: forward and reverse reactions do not stop at equilibrium.

The state of chemical equilibrium is quantitatively characterized by an equilibrium constant, which is the ratio of the constants of the forward (K 1) and reverse (K 2) reactions.

For the reaction mA + nB  pC + dD the equilibrium constant is equal to

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, molecules do not stop colliding, and interactions between them do not stop, but the concentrations of substances remain constant. These concentrations are called equilibrium.

With equilibrium (H 2) = or R equilibrium (HI) = .

Like any other concentration, the equilibrium concentration is measured in moles per liter.

.

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 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 subscript "c" in the designation of this value indicates that concentrations were used to calculate the constant.

C (g) + CO 2 2CO (g)

hard graphite C (g) is involved. Formally, using the law of mass action, we write down 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.

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.

Shift in chemical equilibrium. Le Chatelier's principle

The transfer of an equilibrium chemical system from one equilibrium state to another is called displacement (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 is achieved. Both can be useful in chemical technology. Since almost all reactions are reversible to one degree or another, two problems arise in industry and laboratory practice: how to obtain the product of a “useful” reaction with maximum yield and how to reduce the yield of products of a “harmful” reaction. In both cases, there is a need to shift the equilibrium either towards the reaction products or towards the starting substances. To learn how to do this, you need to know what the equilibrium position of any reversible reaction depends on.

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 participating in the reaction and
3) on pressure (for gas systems it is proportional to the concentrations of substances).
To qualitatively assess the influence on the chemical equilibrium of all these very different factors, an 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 influenced from the outside, then the equilibrium in the system will shift in the direction in which this influence is partially compensated.

As an example of the influence on the equilibrium position of the concentrations of substances participating in the reaction, let us consider the reversible reaction for the production of hydrogen iodide

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 in 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 reactor volume is 1 liter, n(H 2) = 1 mol, n(I 2) = 1 mol and n(HI) = 1 mol. At time t 1 we introduce another 1 mol of HI into the reactor, its concentration will become equal to 2 mol/l. But to TO WITH remained constant, the concentrations of hydrogen and iodine should increase, and this is only possible due to the decomposition of part of the hydrogen iodide according to the equation

2HI (g) = H 2 (g) + I 2 (g).

Let t 2 decompose by the time the new equilibrium state is reached x mole 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 forward reaction v 1 = k 1 [A][B], reverse reaction speed 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 starting substances A and B decrease, and the rate of the direct reaction decreases accordingly. 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 there will come a moment when the rates of forward and reverse reactions become equal. After this, the concentrations of all substances - A, B, C and D do not change over time. This means that the reaction has reached an equilibrium position, and concentrations of substances that do not change over time are called equilibrium. But, unlike mechanical equilibrium, in which all movement stops, in chemical equilibrium both reactions - both direct and reverse - continue to occur, but their speeds are equal and therefore it seems that no changes occur in the system. There are many ways to prove the occurrence of forward and reverse reactions after equilibrium is achieved. For example, if a little hydrogen isotope, deuterium D2, 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 you introduce a little deuterated ammonia NH 2 D into the system, then deuterium will immediately appear in the starting 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. A silver plate was placed in a solution of silver nitrate, and no changes were observed. Then a tiny amount of radioactive silver ions was introduced into the solution, after which the silver plate became radioactive. Neither rinsing the plate with water nor washing it with hydrochloric acid could “wash away” this radioactivity. Only etching with nitric acid or mechanically treating the surface with fine sandpaper rendered it inactive. This experiment can be explained in only one way: 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(s) – e – = Ag +. Therefore, the addition of radioactive Ag + ions to the solution led to their “incorporation” 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 sediments. For example, a solid will dissolve most quickly if it is placed in a pure solvent when the system is far from equilibrium, in this case a saturated solution. Gradually, the rate of dissolution 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, in which the rates of dissolution and crystallization are equal, and the mass of the precipitate does not change over time. How can a system “counteract” changes in external conditions? If, for example, the temperature of an equilibrium mixture is increased by heating, the system itself, of course, cannot “weaken” the external heating, but the equilibrium in it shifts in such a way that heating the reaction system to a certain temperature requires a larger amount of heat than in the case if the equilibrium did not shift. In this case, the equilibrium shifts so that heat is absorbed, i.e. towards an endothermic reaction. This can be interpreted as “the desire of the system to weaken external influence.” 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 be shifted by changing the pressure. As the pressure increases, the equilibrium shifts to the side where the number of gaseous molecules is smaller (and in this way, as it were, “counteracts” the 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 pressure does not affect the equilibrium position. It should be noted that when the temperature changes, the equilibrium constant of the reaction also changes, whereas when only the pressure changes, it remains constant.

Several examples of the use of Le Chatelier's principle to predict shifts in chemical equilibrium. The reaction 2SO 2 + O 2 2SO 3 (g) is exothermic. If the temperature is increased, the endothermic reaction of SO 3 decomposition will take advantage and the equilibrium will shift to the left. If you lower the temperature, the equilibrium will shift to the right. Thus, a mixture of SO 2 and O 2 taken in a stoichiometric ratio of 2:1 ( cm . STOICHIOMERIS), at a temperature of 400 ° C and atmospheric pressure turns into SO 3 with a yield of about 95%, i.e. the equilibrium state 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 hydrogen abstraction from cyclohexane to form 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 occurs 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 exactly the opposite of 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, therefore 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 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 in the equilibrium to the right, but in this case there would be too little substance in the reactor, and the reaction rate would also decrease, so that the overall productivity would not increase, but would decrease. This example once again shows that economically sound industrial synthesis is a successful maneuver between “Scylla and Charybdis”.

Le Chatelier's principle also 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, releases heat and therefore, with increasing temperature, 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. This cycle also works in halogen lamps, where tungsten evaporated from the coil and settled on the colder walls forms volatile compounds with halogens, which disintegrate again on the hot coil, and the tungsten is transferred to its original place.

In addition to changing temperature and pressure, there is another effective way to influence the equilibrium position. Let us imagine that from an equilibrium mixture

A + B C + D a 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 in different ways. For example, if there is sulfur dioxide in a tightly closed container of 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, the 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 sulfurous acid completely decomposes. A similar process can be observed every time you open a bottle of lemonade or mineral water: the equilibrium CO 2 (g) CO 2 (p) + H 2 O H 2 CO 3 shifts to the left as CO 2 evaporates.

Removal of a reagent from the system is possible not only through the formation of gaseous substances, but also by binding one or another reagent to form 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 CaCO 3 precipitate by reacting with carbonic acid; the equilibrium CO 2 (p) + H 2 O H 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. Thus, when dilute solutions of FeCl 3 and KSCN are combined, a reddish-orange color appears as a result of the formation of iron thiocyanate (rhodanide):

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 in equilibrium to the right (as if weakening the external influence). If you add excess KCl to the solution, the equilibrium will shift to the left with the color weakening to light yellow.

It is not for nothing that the formulation of Le Chatelier’s principle indicates that it is possible to predict the results of external influences only for systems that are in a state of equilibrium. If this instruction 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 is applicable only to equilibrium systems, we can draw the incorrect conclusion that with increasing temperature, 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 a “weakening of the external influence.” However, the process of dissolving KOH in water is not at all an equilibrium process, 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 sediment to 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 satisfied in this case as well. 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.

In textbooks and popular literature you can find another interesting and instructive example of the erroneous use of Le Chatelier's principle. If you place an equilibrium mixture of brown nitrogen dioxide NO 2 and colorless tetroxide N 2 O 4 into a transparent gas syringe, and then quickly compress the gas using a piston, the color intensity will immediately intensify, and after some time (tens of seconds) it will weaken again, although will not reach the original one. This experience is usually explained like this. Rapidly compressing the mixture causes the pressure and therefore the concentration of both components to increase, so the mixture becomes darker. But an increase in pressure, in accordance with Le Chatelier’s principle, 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 becomes lighter, 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 equilibrium is in any case established in millionths of a second, so it is impossible to push the piston so quickly as to upset the equilibrium. This experiment can be 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 Le Chatelier principle, the equilibrium instantly shifts towards the endothermic reaction, which occurs 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 to room temperature, and the equilibrium again shifts towards the tetroxide - the mixture becomes lighter.

Le Chatelier's principle also 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 can be purchased with that money. What will happen if the “external influence” turns out to be the government’s desire to print more money to pay off its debts? In strict accordance with Le Chatelier's principle, the balance between goods and money will shift in such a way as to weaken the citizens' pleasure in having more money. Namely, prices for goods and services will rise, and in this way a new equilibrium will be achieved. Another example. In one of the US cities, it was decided to get rid of constant traffic jams by expanding highways and building transport interchanges. This helped for a while, but then delighted residents began to buy more cars, so soon traffic jams reappeared - but with a new “balance” between the roads and more cars.

So, let's draw the main conclusions about ways to shift the chemical equilibrium.

1. Pressure. An increase in pressure (for gases) shifts the equilibrium towards a reaction leading to a decrease in volume (i.e., the formation of fewer molecules).

2. An increase in temperature shifts the equilibrium position towards an endothermic reaction (i.e. towards a reaction that occurs 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 starting substances [A] or [B] or [A] and [B]: V 1 > V 2.

1. 
   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 the population of mice - nut year - there is a lot of food for mice, the mouse population is growing rapidly. As the number of mice increases, the amount of food decreases; as a result of the accumulation of rodents, various infectious diseases begin to grow among mice, so there is a gradual decrease in the size of the rodent population. After a certain period of time, a dynamic equilibrium in the number of mice being born and dying occurs; a shift in this equilibrium can occur in one direction or another under the influence of external, favorable or unfavorable conditions.

Biochemical processes occur in the human body, which can also be regulated according to Le Chatelier’s principle. Sometimes, as a result of such a reaction, the body begins to produce poisonous substances that cause a particular disease. How to prevent this process?

Let us recall such a treatment method as homeopathy. The method consists of using very small doses of those drugs that, in large doses, cause signs of some disease in a healthy person. How does the poison medicine work in this case? A product of an unwanted reaction is introduced into the body, and according to Le Chatelier’s principle, the equilibrium shifts towards the starting substances. The process causing painful disorders in the body fades away.

Practical part.

Monitoring the level of mastery of the studied topic is carried out in the form of tests. A test system of succinctly and precisely formulated and standardized tasks, which must be given within a limited time, brief and accurate answers, assessed according to a point system. When compiling tests, I focused on the following levels:

• 
  Reproductive - students at this level perform mainly based on memory.
• 
  Productive - achieving this level requires students to understand the studied formulations, concepts, laws, and the ability to establish relationships between them.
• 
  Creative - the ability to predict based on existing knowledge, design, analyze, draw conclusions, comparisons, generalizations.

Closed tests or tests in which the test taker must choose the correct answer from the given options.

A) Reproductive level: tests with alternative answers in which the subject must answer yes or no. Score 1 point.

1. 
   Phosphorus combustion reaction -

a) yes b) no

1. 
   Decomposition reaction

reversible reaction

a) yes b) no

1. 
   Temperature increase

mercury oxide II per mercury

and oxygen

a) yes b) no

1. 
   In living systems

and irreversible processes

a) yes b) no.

Tests with a choice of one correct answer

1. 
   In which system will the chemical equilibrium shift to the right as the pressure increases?

1. 
   2HI(g)↔H2(g)+I2(g)
2. 
   C (tv)+S2(g)↔CS2(g)
3. 
   C3H6(g)+H2(g)↔С3H8(g)
4. 
   H2(g)+F2(g)↔2HF(g) 1 point

1. 
   temperature rise
2. 
   using a catalyst
3. 
   decrease in temperature; 1 point

1. 
   On the state of chemical equilibrium in the system

does not affect

1. 
   increase in pressure
2. 
   increasing iodine concentration
3. 
   temperature increase
4. 
   decrease in temperature; 1 point

1. 
   In which system does an increase in hydrogen concentration shift the chemical equilibrium to the left?

1. 
   C(s)+2H2(g)↔СH4(g)
2. 
   2NH3(g)↔N2(g)+3H2(g)
3. 
   2H2(g)+O2(g)↔2H2O(g)
4. 
   FeO(s)+H2(g)↔Fe+H2O(g) 1 point

1. 
   In which system does an increase in pressure not affect the shift in chemical equilibrium?

1. 
   H2(g)+J2(g)↔2HJ(g)
2. 
   SO2(g)+H2O(l)↔H2SO3(g)
3. 
   CH4(g)+H2O(g)↔CO(g)+3H2(g)
4. 
   4HCl(g)+O2(g)↔2H2O(g)+2Сl2(g) 1 point

1. 
   On chemical equilibrium in the system

has no effect

1. 
   temperature increase
2. 
   increase in pressure
3. 
   removing ammonia from the reaction zone
4. 
   use of catalyst 1 point

1. 
   Chemical equilibrium in the system

shifts towards the formation of the reaction product at

1. 
   increased pressure
2. 
   temperature rise
3. 
   decrease in pressure
4. 
   catalyst application 1 point

1. 
   In the production of sulfuric acid at the stage of oxidation of SO2 to SO3 to increase product yield

1. 
   increase oxygen concentration
2. 
   increase the temperature
3. 
   lower blood pressure
4. 
   a catalyst is introduced; 1.5 points

Alkene + H2 ↔ alkane

when heating the reaction mixture

1. 
   balance will shift to the right
2. 
   balance will shift to the left
3. 
   equilibrium will flow in both directions with equal probability
4. 
   these substances are not in a state of equilibrium under the specified conditions; 1.5 points

1. 
   Chemical equilibrium in the system

shifts towards the formation of starting substances

1) increasing pressure

1. 
   temperature rise
2. 
   decrease in pressure
3. 
   catalyst application; 1 point

1. 
   On the shift of equilibrium to the right in the system

influences

1. 
   temperature drop
2. 
   increase in pressure
3. 
   use of catalyst
4. 
   temperature increase; 1 point

1. 
   An irreversible reaction corresponds to the equation

1. 
   nitrogen+hydrogen=ammonia
2. 
   acetylene+oxygen=carbon dioxide+water
3. 
   hydrogen+iodine=hydrogen iodide
4. 
   sulfur dioxide + oxygen = sulfuric anhydride; 1.5 points

Multiple Choice Tests, during which the subject must choose 1-2 correct answers, or compare 2 proposed conditions when choosing an answer.

1. 
   In which system will the chemical equilibrium shift towards the reaction products both with increasing pressure and with decreasing temperature?

1. 
   N2+O2↔2NO-Q
2. 
   N2+3H2↔2NH3+Q
3. 
   H2+CL2↔2HCL+Q
4. 
   C2H2↔2C(tv)+H2-Q 1.5 points

1. 
   Chemical equilibrium in the system

NH3+H2O↔NH4+OH

will shift towards the formation of ammonia when ammonia is added to an aqueous solution

1. 
   sodium chloride
2. 
   sodium hydroxide
3. 
   of hydrochloric acid
4. 
   aluminum chloride; 1.5 points

19) The hydration reaction of ethylene CH2=CH2+H2O ↔ is of great practical importance, but it is reversible; to shift the equilibrium of the reaction to the right it is necessary

1. 
   increase the temperature (>280 degrees C)
2. 
   reduce the amount of water in the reaction mixture
3. 
   increase pressure (more than 80 atmospheres)
4. 
   replace the acid catalyst with platinum; 1 point

1. 
   The butane dehydrogenation reaction is endothermic. To shift the reaction equilibrium to the right it is necessary

1. 
   use a more active catalyst, such as platinum
2. 
   lower the temperature
3. 
   increase blood pressure
4. 
   increase the temperature; 1 point

1. 
   For the reaction of acetic acid with methanol to form ether and water, a shift of equilibrium to the left will contribute to

1. 
   appropriate catalyst
2. 
   adding concentrated sulfuric acid
3. 
   use of dehydrated starting materials
4. 
   adding ether; 1.5 points

Tests to eliminate unnecessary things (if you see something unnecessary, remove it)

1. 
   The balance shift is affected by

1. 
   pressure change
2. 
   use of catalyst
3. 
   change in the concentrations of substances involved in the reaction
4. 
   temperature change; 1 point

1. 
   An increase or decrease in pressure affects the shift in chemical equilibrium in reactions

1. 
   moving with heat release
2. 
   reactions involving gaseous substances
3. 
   reactions occurring with a decrease in volume
4. 
   reactions occurring with an increase in volume; 1.5 points

1. 
   The reaction is irreversible

1. 
   burning coal
2. 
   phosphorus burning
3. 
   synthesis of ammonia from nitrogen and hydrogen
4. 
   methane combustion; 1.5 points

Grouping tests include a list of proposed formulas, equations, terms that should be distributed according to specified characteristics

1. 
   With a simultaneous increase in temperature and decrease in pressure, the chemical equilibrium will shift to the right in the system

1. 
   H2(g)+S(g)↔H2S(g)+Q
2. 
   2SO2(g)+O2(g)↔2SO3(g)+Q
3. 
   2NH3(g)↔N2(g)+3H2(g)-Q
4. 
   2HCL(g)↔H2(g)+CL2(g)-Q; 2 points

1. 
   The hydrogenation reaction of propene is exothermic. To shift the chemical equilibrium to the right it is necessary

1. 
   temperature drop
2. 
   increase in pressure
3. 
   decrease in hydrogen concentration
4. 
   decrease in propene concentration; 1 point

Compliance tasks.

When performing tests, the subject is asked to establish the correspondence of the elements of two lists, with several possible answers.

1. 
   The reaction equilibrium shifts to the right. Bring into compliance.

B) N2+3H2↔2NH3+Q 2) With increasing temperature

B) CO2+C(solid)↔2CO-Q 3) When the pressure decreases

D) N2O(g)+S(s)↔2N2(g) 4) With increasing contact area; 2 points

1. 
   The equilibrium of the reaction shifts towards the formation of reaction products. Bring into compliance.

B) 2H2+O2↔2H2O(g)+Q 2) With increasing temperature

B) 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 free-response tests, in which the subject needs to add concepts to the definition of an equation or offer an independent judgment in an evidentiary way.

Tasks of this type constitute the final, most highly evaluated part of the Unified State Exam tests in chemistry.

Addition tasks.

The subject must formulate answers taking into account the restrictions provided for in the task.

1. 
   Complete the equation of reactions that are reversible and at the same time exothermic

B) Hydrogen + Iodine

B) Nitrogen + Hydrogen

D) Sulfur dioxide + Oxygen

E) Carbon dioxide + Carbon 2 points

1. 
   Write the reaction equation according to the diagram, from them select those reversible reactions in which an increase in temperature will cause a shift of equilibrium to the right:

N2 → NO→ NO2→ HNO3→ NH4NO3 2 points

Tests for free presentation tasks.

The subject must independently formulate the answers, since no restrictions are imposed on them in the task.

CO + 2H2↔ CH3OH(g)+Q 2 points

C (sol) + 2H2(g)↔CH4(g) + Q 2 points

Answers to tests.

Test number Correct answer

1. 
   A-2,3

1. 
   В- N2+3H2↔2NH3+Q

1. 
   1) N2+O2↔2NO-Q

3) 4NO2+2H2O+O2↔4HNO3+Q

4) NH3+HNO3=NH4NO3

first reaction

1. 
   CO+2H2↔CH3OH+Q

1. 
   decreasing temperature
2. 
   increasing pressure
3. 
   increasing CO concentration
4. 
   increasing H2 concentration
5. 
   decrease in alcohol concentration

1. 
   C+2H2↔CH4+Q

2) decrease in pressure

3) decreasing hydrogen concentration

4) increasing methane concentration.

Bibliography

1. 
   Akhmetov, M.A. System of tasks and exercises in organic chemistry in test form [Text] / M.A. Akhmetov, I.N. Prokhorov. - Ulyanovsk: IPKPRO, 2004.
2. 
   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 science (Publishing house “First of September”) - 2007.- No. 22.-p.4-13.
3. 
   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.-160 p.
4. 
   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.
5. 
   Leenson, I.A. Chemical reactions, thermal effect, equilibrium, speed [Text] /I.A.Leenson.M.: Astrel, 2002.-190p.
6. 
   Radetsky, A.M. Test work in chemistry in grades 8-11: a manual for teachers [Text] / A.M. Radetsky. M.: Education, 2009.-272 p.
7. 
   Ryabinina, O.A. Demonstration of the action of Le Chatelier’s principle [Text] / O.O. Ryabinina, A. Illarionov // Chemistry at school. - 2008. - No. 7. - p. 64-67.
8. 
   Tushina.E.N. Le Chatelier's principle and some treatment methods [Text] / E.N. Tushina.// Chemistry at school.-1993. No. 2.-p.54.
9. 
   Shelinsky, G.I. Fundamentals of the theory of chemical processes [Text] / G.I. Shelinsky. M.: Education, 1989.-234 p.
10. 
    Strempler, G.I. Pre-profile preparation in chemistry [Text]

All chemical reactions are divided into two types: reversible and irreversible.

Irreversible are called reactions that proceed in only one direction, i.e. the products of these reactions do not interact with each other to form the starting substances.

An irreversible reaction ends when at least one of the starting substances is completely consumed. Combustion reactions are irreversible; many reactions of thermal decomposition of complex substances; most reactions that result in the formation of precipitation or the release of gaseous substances, etc. For example:

C 2 H 5 OH + 3O 2 → 2CO 2 + 3H 2 O

2KMnO 4 = K 2 MnO 4 + MnO 2 + O 2

BaCl 2 + H 2 SO 4 = BaSO 4 ↓ + 2HCl

Reversible Reactions that simultaneously occur in the forward and reverse directions are called:

In equations of reversible reactions, the reversibility sign is used.

An example of a reversible reaction is the synthesis of hydrogen iodide from and:

Some time after the start of the chemical reaction, not only the final product of the reaction, HI, but also the starting substances -H 2 and I 2 -can be detected in the gas mixture. No matter how long the reaction lasts, the reaction mixture at 350 o C will always contain approximately 80% HI, 10% H 2 and 10% I 2. If we take HI as the starting substance and heat it to the same temperature, we can find that after some time the ratio between the amounts of all three substances will be the same. Thus, during the formation of hydrogen iodide from hydrogen and iodine, direct and reverse reactions occur simultaneously.

If hydrogen and iodine in concentrations and are taken as starting substances, then the rate of the direct reaction at the initial moment of time was equal to: v pr = k pr ∙ . The rate of the reverse reaction v arr = k arr 2 at the initial moment of time is zero, since there is no hydrogen iodide in the reaction mixture ( = 0). Gradually, the rate of the forward reaction decreases as hydrogen and iodine react and their concentrations decrease. In this case, the rate of the reverse reaction increases because the concentration of hydrogen iodide formed gradually increases. When the rates of forward and reverse reactions become equal, chemical equilibrium occurs. In a state of equilibrium, over a certain period of time, the same number of HI molecules are formed as they disintegrate into H 2 and I 2 .

The state of a reversible reaction, in which the rate of the forward reaction is equal to the rate of the reverse reaction, is called chemical equilibrium.

Chemical equilibrium is a dynamic equilibrium. In an equilibrium state, both forward and reverse reactions continue to occur, but since their rates are equal, the concentrations of all substances in the reaction system do not change. These concentrations are called equilibrium concentrations.

Chemical equilibrium shift

Le Chatelier's principle

Chemical equilibrium is mobile. When external conditions change, the rates of forward and reverse reactions may become unequal, which causes a shift in the equilibrium.

If, as a result of an external influence, the rate of the forward reaction becomes greater than the rate of the reverse reaction, then we speak of a shift in equilibrium right(towards direct reaction). If the rate of the reverse reaction becomes greater than the rate of the forward reaction, then we speak of a shift in equilibrium left(towards the opposite reaction). The result of a shift in equilibrium is the transition of the system to a new equilibrium state with a different ratio of concentrations of reacting substances.

The direction of the equilibrium shift is determined by the principle that was formulated by the French scientist Le Chatelier (1884):

If an external influence is exerted on an equilibrium system, then the equilibrium shifts towards the reaction (direct or reverse) that counteracts this influence.

The most important external factors that can lead to a shift in chemical equilibrium are:

a) concentrations of reacting substances;

b) temperature;

c) pressure.

Effect of concentration of reactants

If any of the substances participating in the reaction is introduced into the equilibrium system, then the equilibrium shifts towards the reaction during which this substance is consumed. If any substance is removed from an equilibrium system, then the equilibrium shifts towards the reaction during which this substance is formed.

For example, consider which substances should be introduced and which substances should be removed from the equilibrium system to shift the reversible synthesis reaction to the right:

To shift the equilibrium to the right (towards the direct reaction of ammonia formation), it is necessary to introduce hydrogen into the equilibrium mixture (i.e., increase their concentrations) and remove ammonia from the equilibrium mixture (i.e., decrease its concentration).

Effect of temperature

Forward and reverse reactions have opposite thermal effects: if the forward reaction is exothermic, then the reverse reaction is endothermic (and vice versa). When the system is heated (i.e., its temperature increases), the equilibrium shifts towards the endothermic reaction; upon cooling (lower temperature), the equilibrium shifts towards the exothermic reaction.

For example, the ammonia synthesis reaction is exothermic: N 2 (g) + 3H 2 (g) → 2NH 3 (g) + 92 kJ, and the ammonia decomposition reaction (reverse reaction) is endothermic: 2NH 3 (g) → N 2 (g) + 3H 2 (g) - 92 kJ. Therefore, an increase in temperature shifts the equilibrium towards the reverse reaction of ammonia decomposition.

Effect of pressure

Pressure affects the equilibrium of reactions in which gaseous substances take part. If the external pressure increases, then the equilibrium shifts towards the reaction during which the number of gas molecules decreases. Conversely, the equilibrium shifts towards the formation of more gaseous molecules as the external pressure decreases. If the reaction proceeds without changing the number of molecules of gaseous substances, then pressure does not affect the equilibrium in this system.

For example, to increase the yield of ammonia (shift to the right), it is necessary to increase the pressure in the reversible reaction system, since during a direct reaction the number of gaseous molecules decreases (from four molecules of nitrogen and hydrogen gases two molecules of ammonia gas are formed).

Video tutorial 2: Chemical equilibrium shift

Lecture: Reversible and irreversible chemical reactions. Chemical balance. Shift in chemical equilibrium under the influence of various factors

From the previous lesson, you learned what the rate of a chemical reaction is and what factors influence it. In this lesson we will look at how these reactions occur. This depends on the behavior of the starting substances participating in the reaction - the reagents. If they are completely converted into final substances - products, then the reaction is irreversible. Well, if the final products are converted back into the original substances, then the reaction is reversible. Taking this into account, let us formulate the definitions:

Reversible reaction- this is a certain reaction that occurs under the same conditions in forward and reverse directions.

Remember, in chemistry lessons you were shown a clear example of a reversible reaction for the production of carbonic acid:

CO 2 + H 2 O<->H2CO3

Irreversible reaction- this is a certain chemical reaction that goes to completion in one specific direction.

An example is the phosphorus combustion reaction: 4P + 5O 2 → 2P 2 O 5

Some evidence of the irreversibility of a reaction is the formation of a precipitate or the release of gas.

When the rates of forward and reverse reactions are equal, it occurs chemical equilibrium.

That is, in reversible reactions, equilibrium mixtures of reactants and products are formed. Let us see with an example how a chemical equilibrium is formed. Let's take the reaction of hydrogen iodide formation:

H 2 (g) + I 2 (g)<->2HI(g)

We can heat a mixture of gaseous hydrogen and iodine or ready-made hydrogen iodine, the result in both cases will be the same: the formation of an equilibrium mixture of three substances H 2, I 2, HI.

At the very beginning of the reaction, before the formation of hydrogen iodide, a direct reaction occurs at a rate of ( v etc ). Let us express it by the kinetic equation v pr = k 1, where k 1 is the rate constant of the forward reaction. The product HI is gradually formed, which, under the same conditions, begins to decompose into H 2 and I 2. The equation for this process is as follows: v arr = k 2 2, where v rev – reverse reaction rate, k 2 – reverse reaction rate constant. At the moment when HI is sufficient for leveling v at v chemical equilibrium occurs. The amount of substances in equilibrium, in our case these are H 2, I 2 and HI, does not change over time, but only if there are no external influences. From the above it follows that chemical equilibrium is dynamic. In our reaction, hydrogen iodide is either formed or consumed.

Remember, changing the reaction conditions allows you to move the equilibrium in the desired direction. If we increase the concentration of iodine or hydrogen, it will increase v Thus, there will be a shift to the right, more hydrogen iodide will be formed. If we increase the concentration of hydrogen iodide, it will increase v arr, and the shift will be to the left. We can get more/less reagents and products.

Thus, chemical equilibrium tends to resist external influences. The addition of H 2 or I 2 ultimately leads to an increase in their consumption and an increase in HI. And vice versa. This process in science is called Le–Chatelier principle. It reads:

If a system that is in stable equilibrium is influenced from the outside (by changing temperature, or pressure, or concentration), then a shift will occur in the direction of a process that weakens this influence.

Remember, a catalyst cannot shift the equilibrium. He can only speed up its onset.

Change in concentration . Above, we looked at how this factor shifts the equilibrium either in the forward or in the opposite direction. If the concentration of reactants is increased, the equilibrium shifts to the side where this substance is consumed. If you reduce the concentration, it shifts to the side where this substance is formed. Remember, the reaction is reversible, and the reactants can be substances on both the right and left sides, depending on which reaction we are considering (direct or reverse).

Influencet . Its increase provokes a shift in equilibrium towards the endothermic reaction (- Q), and a decrease towards the exothermic reaction (+ Q). The reaction equations indicate the thermal effect of the forward reaction. The thermal effect of the reverse reaction is the opposite. This rule is only suitable for reactions with a thermal effect. If it is not there, then t is not capable of shifting the equilibrium, but its increase will accelerate the process of the emergence of equilibrium.

Effect of pressure . This factor can be used in reactions involving gaseous substances. If the moles of gas are zero, there will be no changes. As pressure increases, the equilibrium shifts towards smaller volumes. As the pressure decreases, the equilibrium will shift towards larger volumes. Volumes - look at the coefficients of gaseous substances in the reaction equation.

Among the numerous classifications of types of reactions, for example those that are determined by the thermal effect (exothermic and endothermic), by changes in the oxidation states of substances (redox), by the number of components participating in them (decomposition, compounds) and so on, reactions occurring in two mutual directions, otherwise called reversible . An alternative to reversible reactions are reactions irreversible, during which the final product (precipitate, gaseous substance, water) is formed. Among these reactions are the following:

Exchange reactions between salt solutions, during which either an insoluble precipitate is formed - CaCO 3:

Ca(OH) 2 + K 2 CO 3 → CaCO 3↓ + 2KON (1)

or a gaseous substance - CO 2:

3 K 2 CO 3 + 2H 3 RO 4 →2K 3 RO 4 + 3 CO 2+ 3H 2 O (2)

or a slightly dissociable substance is obtained - H 2 O:

2NaOH + H 2 SO 4 → Na 2 SO 4 + 2 H 2O(3)

If we consider a reversible reaction, then it proceeds not only in the forward direction (in reactions 1,2,3 from left to right), but also in the opposite direction. An example of such a reaction is the synthesis of ammonia from gaseous substances - hydrogen and nitrogen:

3H 2 + N 2 ↔ 2NH 3 (4)

Thus, a chemical reaction is called reversible if it proceeds not only in the forward direction (→), but also in the reverse direction (←) and is indicated by the symbol (↔).

The main feature of this type of reaction is that reaction products are formed from the starting substances, but at the same time, the starting reagents are formed from the same products. If we consider reaction (4), then in a relative unit of time, simultaneously with the formation of two moles of ammonia, their decomposition will occur with the formation of three moles of hydrogen and one mole of nitrogen. Let us denote the rate of direct reaction (4) by the symbol V 1, then the expression for this rate will take the form:

V 1 = kˑ [Н 2 ] 3 ˑ , (5)

where the value “k” is defined as the rate constant of a given reaction, the values ​​[H 2 ] 3 and correspond to the concentrations of the starting substances raised to powers corresponding to the coefficients in the reaction equation. In accordance with the principle of reversibility, the rate of the reverse reaction will take the expression:

V 2 = kˑ 2 (6)

At the initial moment of time, the rate of the forward reaction takes on the greatest value. But gradually the concentrations of the starting reagents decrease and the reaction rate slows down. At the same time, the rate of the reverse reaction begins to increase. When the rates of forward and reverse reactions become the same (V 1 = V 2), state of equilibrium , at which there is no longer a change in the concentrations of both the initial and the resulting reagents.

It should be noted that some irreversible reactions should not be taken literally. Let us give an example of the most frequently cited reaction of a metal with an acid, in particular, zinc with hydrochloric acid:

Zn + 2HCl = ZnCl 2 + H 2 (7)

In fact, zinc, when dissolved in acid, forms a salt: zinc chloride and hydrogen gas, but after some time the rate of the direct reaction slows down as the concentration of salt in the solution increases. When the reaction practically stops, a certain amount of hydrochloric acid will be present in the solution along with zinc chloride, so reaction (7) should be given in the following form:

2Zn + 2HCl = 2ZnНCl + H2 (8)

Or in the case of the formation of an insoluble precipitate obtained by merging solutions of Na 2 SO 4 and BaCl 2:

Na 2 SO 4 + BaCl 2 = BaSO 4 ↓ + 2NaCl (9)

the precipitated salt BaSO 4, albeit to a small extent, will dissociate into ions:

BaSO 4 ↔ Ba 2+ + SO 4 2- (10)

Therefore, the concepts of irreversible and irreversible reactions are relative. But nevertheless, both in nature and in the practical activities of people, these reactions are of great importance. For example, combustion processes of hydrocarbons or more complex organic substances, such as alcohol:

CH 4 + O 2 = CO 2 + H 2 O (11)

2C 2 H 5 OH + 5O 2 = 4CO 2 + 6H 2 O (12)

are absolutely irreversible processes. It would be considered a happy dream of humanity if reactions (11) and (12) were reversible! Then it would be possible to synthesize gas and gasoline and alcohol again from CO 2 and H 2 O! On the other hand, reversible reactions such as (4) or oxidation of sulfur dioxide:

SO 2 + O 2 ↔ SO 3 (13)

are basic in the production of ammonium salts, nitric acid, sulfuric acid, and other inorganic and organic compounds. But these reactions are reversible! And in order to obtain the final products: NH 3 or SO 3, it is necessary to use such technological methods as: changing the concentrations of reagents, changing pressure, increasing or decreasing the temperature. But this will already be the subject of the next topic: “Shift in chemical equilibrium.”

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Very often, chemical reactions proceed in such a way that the primary reactants are completely converted into reaction products. For example, if you put a zinc granule in hydrochloric acid, then with a certain (sufficient) amount of acid the reaction will proceed until the zinc is completely dissolved according to the equation: 2HCL + ZN = ZnCl 2 + H 2.

If you carry out this reaction in the opposite direction, in other words, pass hydrogen through a solution of zinc chloride, then metallic zinc is formed - this reaction cannot proceed in the opposite direction, so it is irreversible.

A chemical reaction, as a result of which primary substances are almost completely converted into final products, is called irreversible.

Such reactions include both heterogeneous and homogeneous reactions. For example, combustion reactions of simple substances - methane CH4, carbon disulfide CS2. As we already know, combustion reactions are exothermic reactions. In most cases, exothermic reactions include compound reactions, for example, the lime slaking reaction: CaO + H 2 O = Ca(OH) 2 + Q (heat is released).

It would be logical to assume that endothermic reactions include reverse reactions, i.e. decomposition reaction. For example, the reaction of burning limestone: CaCo 3 = CaO + CO 2 – Q (heat is absorbed).

It must be remembered that the number of irreversible reactions is not so large.

Homogeneous reactions (between solutions of substances) are irreversible if they occur with the formation of an insoluble, gaseous product or water. This rule is called "Berthollet's rule". Let's conduct an experiment. Let's take three test tubes and pour 2 ml of hydrochloric acid solution into them. Add 1 ml of phenolphthalein-colored raspberry alkali solution to the first vessel; it will lose color as a result of the reaction: HCl + NaOH = NaCl + H 2 O.

Add 1 ml of sodium carbonate solution to the second test tube - we will see a violent boiling reaction, which is caused by the release of carbon dioxide: Na 2 CO 3 + 2HCl = 2NaCl + H 2 O + CO 2.

Let's add a few drops of silver nitrate to the third test tube and see how a whitish precipitate of silver chloride has formed in it: HCl + AgNO 3 = AgCl↓ + HNO 3.

Most reactions are reversible. There are not very many irreversible reactions.

Chemical reactions that can occur simultaneously in two opposite directions - forward and reverse - are called reversible.

Let's pour 3 ml of water into a test tube and add a few pieces of litmus, and then begin to pass through it using a gas outlet tube the carbon dioxide coming out of another vessel, which is formed due to the interaction of marble and hydrochloric acid. After some time, we will see the purple litmus turn red, this indicates the presence of acid. We obtained fragile carbonic acid, which was formed by combining carbon dioxide and water: CO 2 + H 2 O = H 2 CO 3.

Let's leave this solution in the tripod. After some time, we will notice that the solution has turned purple again. The acid decomposed into its original components: H 2 CO 3 = H 2 O + CO 2.

This process will occur much faster if we heat the carbonic acid solution. Thus, we have found that the reaction to produce carbonic acid can occur in both forward and reverse directions, which means it is reversible. The reversibility of a reaction is indicated in writing by two oppositely directed arrows: CO 2 + H 2 O ↔ H 2 CO 3 .

Among the reversible reactions that underlie the production of important chemical products, we give as an example the reaction of the synthesis of sulfur oxide (VI) from sulfur oxide (IV) and oxygen: 2SO 2 + O 2 ↔ 2SO 3 + Q.

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