Chemically irreversible reactions under these conditions they go almost to the end, until the complete consumption of one of the reactants (NH4NO3 → 2H2O + N2O - no attempt to obtain nitrate from H2O and N2O leads to a positive result).
Chemically reversible reactions proceed simultaneously under the given conditions both in the forward and in the opposite direction. There are fewer irreversible reactions than reversible ones. An example of a reversible reaction is the interaction of hydrogen with iodine.
After some time, the rate of HI formation will become equal to the rate of its decomposition.
In other words, chemical equilibrium will come.
Chemical equilibrium is the state of the system at which the rate of formation of the reaction products is equal to the rate of their transformation into the initial reagents.
Chemical equilibrium is dynamic, that is, its establishment does not mean the termination of the reaction.
Mass action law:
The mass of the substances that have entered the reaction is equal to the mass of all the reaction products.
The law of mass action establishes the ratio between the masses of reactants in chemical reactions at equilibrium, as well as the dependence of the rate chemical reaction on the concentration of the starting substances.
Signs of true chemical equilibrium:
1.the state of the system remains unchanged over time in the absence of external influences;
2. the state of the system changes under the influence of external influences, no matter how small they are;
3. the state of the system does not depend on which side it comes to equilibrium.
When equilibrium is established, the product of the concentrations of the reaction products divided by the product of the concentrations of the starting substances, in powers equal to the corresponding stoichiometric coefficients, for a given reaction at a given temperature is a constant value called the equilibrium constant.
The concentrations of reagents at steady state are called equilibrium concentrations.
In the case of heterogeneous reversible reactions, the expression for Kc includes only equilibrium concentrations of gaseous and dissolved substances. So, for the reaction CaCO3 ↔ CaO + CO2
Under unchanged external conditions, the equilibrium position remains indefinitely. When external conditions change, the position of equilibrium can change. A change in temperature, concentration of reagents (pressure for gaseous substances) leads to a violation of the equalities of the rates of direct and reverse reactions and, accordingly, to an imbalance. After some time, the equality of speeds will be restored. But the equilibrium concentrations of reagents in the new conditions will be different. System transition from one equilibrium state to another called balance shift or shift ... Chemical equilibrium can be compared to the position of a balance beam. Just as it changes from the pressure of the load on one of the cups, the chemical equilibrium can shift towards a forward or reverse reaction depending on the process conditions. Each time, a new equilibrium is established, corresponding to new conditions.
The numerical value of a constant usually changes with temperature. At constant temperature, the Kc values \u200b\u200bdo not depend on pressure, volume, or concentration of substances.
Knowing the numerical value of Kc, one can calculate the values \u200b\u200bof equilibrium concentrations or pressures of each of the reaction participants.
Direction chemical equilibrium displacement as a result of changes in external conditions, it is determined le Chatelier's principle:
if an external influence is exerted on the equilibrium system, then the equilibrium is shifted to the side opposing this influence.
Dissolution as a physicochemical process. Solvation. Solvates. Special properties water as a solvent. Hydrates. Crystalline hydrates. Solubility of substances. Dissolution of solid, liquid and gaseous substances. Influence of temperature, pressure and nature of substances on solubility. Ways of expressing the composition of solutions: mass fraction, molar concentration, equivalent concentration and mole fraction.
There are two main theories of solutions: physical and chemical.
Physical theory of solutions was proposed by the laureates Nobel Prize the Dutchman J. Van't Hoff (1885) and the Swedish physicist-chemist S. Arrhenius (1883). A solvent is considered as a chemically inert medium in which particles (molecules, ions) of a solute are evenly distributed. It is assumed that there is no intermolecular interaction, both between the particles of the solute and between the molecules of the solvent and the particles of the solute. Particles of solvent and solute are evenly distributed in the volume of the solution due to diffusion. Subsequently, it turned out that the physical theory satisfactorily describes the nature of only a small group of solutions, the so-called ideal solutions, in which the particles of the solvent and the solute do not really interact with each other. Many gas solutions are examples of ideal solutions.
Chemical (or solvation) theory of solutions proposed by D.I. Mendeleev (1887). For the first time, using huge experimental material, he showed that a chemical interaction occurs between the particles of a solute and solvent molecules, as a result of which unstable compounds of variable composition are formed, called solvates or hydrates ( if the solvent is water). DI. Mendeleev defined a solution as a chemical system, all forms of interaction in which are associated with the chemical nature of the solvent and the substances being dissolved. The main role in education solvates fragile intermolecular forces and hydrogen bonds play.
Dissolution process cannot be represented by a simple physical model, for example, the statistical distribution of a solute in a solvent as a result of diffusion. It is usually accompanied by a noticeable thermal effect and a change in the volume of the solution, due to the destruction of the structure of the solute and the interaction of the particles of the solvent with the particles of the solute. Both of these processes are accompanied by energetic effects. To destroy the structure of the solute, it is required energy expenditure , while the interaction of the particles of the solvent and the solute is accompanied by the release of energy. Depending on the ratio of these effects, the dissolution process can be endothermic or exothermic.
When copper sulfate dissolves, the presence of hydrates can be easily detected by a color change: anhydrous salt whitedissolving in water forms a blue solution. Sometimes hydrated water firmly binds to the solute and, when it is released from the solution, is part of its crystals. Crystalline substancescontaining water, are called crystalline hydrates , and the water included in the structure of such crystals is called crystallization water. The composition of crystalline hydrates determines the formula of a substance, which indicates the number of molecules of crystallization water per one molecule. So, the formula for the crystalline hydrate of copper sulfate (copper sulfate) is CuSO4 × 5H2O. The retention of the color characteristic of the corresponding solutions by crystalline hydrates serves as direct evidence of the existence of similar hydrate complexes in solutions. The color of the crystalline hydrate depends on the number of crystallization water molecules.
Exists different ways solution composition expressions... Most commonly used mass fraction solute, molar and normal concentration.
In general terms, concentration can be expressed as the number of particles per unit volume or as the ratio of the number of particles of a given species to the total number of particles in solution. The amount of solute and solvent is measured in units of mass, volume, or moles. Generally, solution concentration Is the amount of a solute in a condensed system (mixture, alloy, or in a certain volume of solution). There are various ways of expressing the concentration of solutions, each of which has a predominant application in a particular field of science and technology. Usually, the composition of solutions is expressed using dimensionless (mass and molar fractions) and dimensional quantities (molar concentration of a substance, molar concentration of a substance - equivalent and molality).
Mass fraction- value, equal ratio the mass of the solute (m1) to the total mass of the solution (m).
Reversible and irreversible chemical reactions. Chemical equilibrium. Equilibrium shift under the influence of various factors
Chemical equilibrium
Chemical reactions proceeding in one direction are called irreversible.
Most chemical processes are reversible... This means that under the same conditions, both direct and reverse reactions occur (especially when it comes to closed systems).
For instance:
a) reaction
$ CaCO_3 (→) ↖ (t) CaO + CO_2 $
irreversible in an open system;
b) the same reaction
$ CaCO_3⇄CaO + CO_2 $
in a closed system is reversible.
Let us consider in more detail the processes occurring during reversible reactions, for example, for a conditioned reaction:
Based on the law of mass action, the speed of the direct reaction
$ (υ) ↖ (→) \u003d k_ (1) C_ (A) ^ (α) C_ (B) ^ (β) $
Since the concentrations of substances $ A $ and $ B $ decrease with time, the rate of the direct reaction also decreases.
The appearance of reaction products means the possibility of a reverse reaction, and with time the concentrations of substances $ C $ and $ D $ increase, which means that the rate of the reverse reaction also increases:
$ (υ) ↖ (→) \u003d k_ (2) C_ (C) ^ (γ) C_ (D) ^ (δ) $
Sooner or later, a state will be reached in which the rates of the forward and reverse reactions become equal
${υ}↖{→}={υ}↖{←}$
The state of the system in which the rate of the forward reaction is equal to the rate of the reverse reaction is called chemical equilibrium.
At the same time, the concentrations of reactants and reaction products remain unchanged. They are called equilibrium concentrations... At the macro level, nothing seems to change overall. But in fact, both direct and reverse processes continue to go on, but with equal speed... Therefore, such an equilibrium in the system is called mobile and dynamic.
Equilibrium constant
Let us denote the equilibrium concentrations of substances by $ [A], [B], [C], [D] $.
Then since $ (υ) ↖ (→) \u003d (υ) ↖ (←), k_ (1) · [A] ^ (α) · [B] ^ (β) \u003d k_ (2) · [C] ^ (γ) · [D] ^ (δ) $, whence
$ ([C] ^ (γ) · [D] ^ (δ)) / ([A] ^ (α) · [B] ^ (β)) \u003d (k_1) / (k_2) \u003d K_ (equal) $
where $ γ, δ, α, β $ - exponents equal to the coefficients in the reversible reaction; $ K_ (equal) $ - chemical equilibrium constant.
The resulting expression quantitatively describes the state of equilibrium and is a mathematical expression of the law of mass action for equilibrium systems.
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 the initial substances (denominator), which is established at equilibrium.
The equilibrium constants are calculated from experimental data by determining the equilibrium concentrations of the starting materials and reaction products at a certain temperature.
The value of the equilibrium constant characterizes the yield of the reaction products, the completeness of its course. If you get $ K_ (equal) \u003e\u003e 1 $, this means that in equilibrium $ [C] ^ (γ) · [D] ^ (δ) \u003e\u003e [A] ^ (α) · [B] ^ ( β) $, i.e., the concentrations of the reaction products prevail over the concentrations of the starting substances, and the yield of the reaction products is high.
For $ K_ (equal)
$ CH_3COOC_2H_5 + H_2O⇄CH_3COOH + C_2H_5OH $
equilibrium constant
$ K_ (equal) \u003d () / () $
at $ 20 ° С $ is $ 0.28 $ (i.e. less than $ 1 $). This means that a significant part of the ether was not hydrolyzed.
In the case of heterogeneous reactions, the expression for the equilibrium constant includes the concentrations of only those substances that are in the gas or liquid phase. For example, for the reaction
the equilibrium constant is expressed as follows:
$ K_ (equal) \u003d (^ 2) / () $
The value of the equilibrium constant depends on the nature of the reacting substances and temperature.
The constant does not depend on the presence of a catalyst, since it changes the activation energy of both the direct and reverse reactions by the same amount. The catalyst can only accelerate the onset of equilibrium without affecting the value of the equilibrium constant.
Equilibrium shift under the influence of various factors
The state of equilibrium is maintained for an arbitrarily long time under constant external conditions: temperature, concentration of starting substances, pressure (if gases are involved or formed in the reaction).
By changing these conditions, it is possible to transfer the system from one equilibrium state to another that meets the new conditions. Such a transition is called displacement or balance shift.
Let us consider different ways of shifting equilibrium using the example of the reaction of interaction of nitrogen and hydrogen with the formation of ammonia:
$ N_2 + 3H_2⇄2HN_3 + Q $
$ K_ (equal) \u003d (^ 2) / (^ 3) $
Effect of changes in the concentration of substances
When nitrogen $ N_2 $ and hydrogen $ Н_2 $ are added to the reaction mixture, the concentration of these gases increases, which means that the rate of the direct reaction increases. The equilibrium shifts to the right, towards the reaction product, i.e. towards ammonia $ NH_3 $.
The same conclusion can be made by analyzing the expression for the equilibrium constant. With an increase in the concentration of nitrogen and hydrogen, the denominator increases, and since $ K_ (equal) $ is a constant value, the numerator must increase. Thus, the amount of the reaction product $ NH_3 $ will increase in the reaction mixture.
An increase in the concentration of the reaction product of ammonia $ NH_3 $ will lead to a shift of 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 changes
A change in pressure affects only those systems where at least one of the substances is in a gaseous state. With increasing pressure, the volume of gases decreases, which means that their concentration increases.
Suppose that the pressure in a closed system has increased, for example, $ 2 $ times. This means that the concentrations of all gaseous substances ($ N_2, H_2, NH_3 $) in the reaction we are considering will increase 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 disturbed. To restore it, the concentration of ammonia must increase and the concentration of nitrogen and hydrogen must decrease. The balance will shift to the right. The change in pressure has practically no effect on the volume of liquids and solids, i.e. does not change their concentration. Consequently, the state of chemical equilibrium of reactions in which gases are not involved does not depend on pressure.
Effect of temperature change
As the temperature rises, as you know, the rates of all reactions (exo- and endothermic) increase. Moreover, an increase in temperature has a greater effect on the rate of those reactions that have a high activation energy, and therefore endothermic.
Thus, the rate of the reverse reaction (in our example, endothermic) increases more than the rate of the forward reaction. The equilibrium will shift towards the process, accompanied by the absorption of energy.
The direction of the displacement of the equilibrium can be predicted using Le Chatelier's principle (1884):
If an external influence is exerted on a system in equilibrium (concentration, pressure, temperature changes), then 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 the initial substances;
- with increasing pressure, the chemical equilibrium of the system shifts in the direction of the reaction in which the volume of formed gaseous substances is less;
- as the temperature rises, the chemical equilibrium of the system shifts towards the endothermic reaction;
- with decreasing temperature - towards the exothermic process.
Le Chatelier's principle is applicable not only to chemical reactions, but also to many other processes: evaporation, condensation, melting, crystallization, etc. In the production of the most important chemical products, Le Chatelier's principle and calculations arising from the law of mass action make it possible to find such conditions for carrying out chemical processes that ensure the maximum yield of the desired substance.
All chemical reactions can be divided into two groups: irreversible and reversiblee reactions. Irreversible reactions flow to the end (until the complete consumption of one of the reagents), and in reversible none of the reactants is completely consumed, because a reversible reaction can proceed both in the forward and in the opposite direction.
An example of an irreversible reaction:
Zn + 4HNO 3 → Zn (NO 3) 2 + 2NO 2 + 2H 2 O
An example of a reversible reaction:
Initially, the speed of the direct reaction v pr is large, and the speed of the reverse reaction v about is zero
 | Dependence of the rates of forward and reverse reactions on time τ. When these rates are equal, chemical equilibrium occurs. |
As the reaction proceeds, the starting materials are consumed and their concentrations decrease. At the same time, reaction products appear, their concentrations increase. As a result, a reverse reaction begins to occur, and its speed gradually increases. When the rates of the forward and reverse reactions become the same, chemical equilibrium occurs. It is dynamic, because, although the concentrations of substances in the system remain constant, the reaction continues to proceed both in the forward and in the opposite direction.
With equality v pr and v about one can equate their expressions according to the law of action of the masses *. For example, for the reversible interaction of hydrogen with iodine:
k pr \u003d k about 2 or
Attitude rate constants of forward and reverse reactions (K) is called the equilibrium constant. At a constant temperature, the equilibrium constant is a constant value that shows the ratio between the concentrations of products and starting substances, which is established at equilibrium. The quantity K depends on the nature of the reactants and on the temperature.
The system is in a state of equilibrium as long as the external conditions remain constant. With an increase in the concentration of any of the substances participating in the reaction, the equilibrium shifts towards the consumption of this substance; with a decrease in the concentration of any of the substances, the equilibrium shifts towards the formation of this substance.
Chemical reactions on the basis of reversibility are divided into irreversible and reversible. Irreversible reactions include those reactions that proceed until one of the reagents is completely consumed. Signs of irreversible reactions occurring in solutions are: a) precipitation, b) gas formation, c) formation of a weak electrolyte.
Reversible reactions are those reactions that proceed simultaneously in two mutually opposite directions. For such reactions, instead of an equal sign, use oppositely directed arrows (-).
Over time, the rate of any reaction, measured by the decreasing concentrations of the initial substances, will decrease, since as the substances interact, their concentrations decrease (the rate of the direct reaction). If the reaction is reversible, then as the concentration of the products increases, its rate will increase (the rate of the reverse reaction). As soon as the rates of the forward and reverse reactions become the same, chemical equilibrium is established in the system and further change the concentration of all substances in the system stops.
The quantitative characteristic of the state of equilibrium is the constant of chemical equilibrium K, which is determined by the ratio of the rate constants of the forward and reverse reactions
In the overwhelming majority of cases, the rate constants of the forward and reverse reactions are not equal. The equilibrium constant is a constant value at a given temperature and determines the ratio between the equilibrium concentrations of the reaction products and the initial substances, raised to the degree of their stoichiometric coefficients. For example, for the process
The square bracket denotes the concentration of each substance at the moment of equilibrium, the so-called equilibrium concentration.
The equilibrium constant depends on the nature of the reactants and the temperature. The catalyst does not affect the equilibrium state. The presence of a catalyst in the system only changes the time to reach it. The system can be in equilibrium until at least one of the external influences changes: temperature, concentration of one of the reagents, pressure (for gases). Changes occurring in an equilibrium system as a result of external influences are determined by the principle of mobile equilibrium (Le Chatelier's principle): an external influence on a system in a state of equilibrium leads to a shift of this equilibrium in the direction in which the effect of the produced influence is weakened.
The balance shift is influenced by:
- 1) temperature change: the endothermic process is accelerated to a greater extent with increasing temperature, and, conversely, with decreasing temperature, the exothermic process is accelerated;
- 2) a change in pressure (for reactions proceeding in the gas phase): with increasing pressure, the equilibrium of the reaction shifts in the direction of the formation of substances occupying a smaller volume, and, conversely, a decrease in pressure promotes a process accompanied by an increase in volume. If the reaction proceeds without a change in volume, then a change in pressure in the system does not affect the chemical equilibrium.
- 3) change in concentration: an increase in the concentration of the initial substances leads to an increase in the rate of the direct reaction, while the process proceeding in the system will end when the rates of the direct and reverse reactions become equal and a new equilibrium is established. A decrease in the concentration of one of the reaction products (withdrawal from the system) leads to a shift in equilibrium towards its formation.
One of the most important characteristics of a chemical reaction is the depth (degree) of conversion, which shows how much the starting materials are converted into reaction products. The larger it is, the more economical the process can be carried out. The conversion rate depends, among other factors, on the reversibility of the reaction.
Reversible reactions , unlike irreversible, do not proceed completely: none of the reactants is completely consumed. At the same time, the reaction products interact with the formation of the starting materials.
Let's consider some 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 chemical bonds can be rearranged to form an intermediate compound (activated complex, see p. 1.3.1). Further restructuring of bonds can lead to the decomposition of the intermediate 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 prevents 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, in this system, two reactions will occur simultaneously - the formation of hydrogen iodide and its decomposition. They can be expressed in one general equation
H 2 + I 2 "2HI
The reversibility of the process is indicated 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 you mix two moles of sulfur dioxide with one mole of oxygen, create conditions in the system favorable for the reaction, and after a time, analyze the gas mixture, the results will show that 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 a starting material, then it will be possible to find that part of it decomposes into oxygen and sulfur oxide (+4), and the final ratio between the amounts of all three substances will be the same as in the case when they started from a mixture of sulfur dioxide and oxygen.
Thus, the interaction of sulfur dioxide with oxygen is also one example 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 a sufficient amount of 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 - passing hydrogen through a solution of ferric 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 reversibly under certain conditions, i.e. in principle, all reactions can be considered reversible. But very often one of the reactions clearly predominates. This happens in cases when the interaction products are removed from the reaction sphere: a precipitate is formed, gas is released, and practically non-dissociating products are formed during ion-exchange reactions; or when the opposite process is practically suppressed due to the obvious excess of the starting substances. Thus, the natural or artificial exclusion of the possibility of a back reaction can 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 also somewhat soluble, and the complex cation 2+ is not absolutely stable, and water dissociates, albeit to an extremely insignificant extent.