What is the equilibrium constant. chemical balance. The law of active masses. Chemical equilibrium constant and methods of its expression. Ways to Express the Equilibrium Constant
If we consider a reversible gas reaction and consider gases to be ideal, then the following relations hold:
1) The Mendeleev–Clapeyron law pV = nRT (or p i V = n i RT).
Where 
,
(19)
From here, the following relations can be obtained, showing the relationship between k p and k c
, but because p i = c i RT, then we get:
.
Here is the change in the number of moles of gaseous substances as a result of one run of this reaction. 
(20)
k p \u003d k c (RT) . (21)
If the partial pressures of the components are expressed in atmospheres, and the concentrations in 
, then in the relation (21) connecting k р and k с, instead of R, one should put the value R = 0.082 
.
k p does not depend on pressure and concentration, k c does not depend on concentration and pressure.
3.5 Equilibrium in heterogeneous reactions.
So far, we have been talking about homogeneous reactions (occurring in one phase). Consider heterogeneous reactions in which not all substances are in the gaseous state. In the case of heterogeneous systems, in which liquid or solid substances do not form solutions with each other and with gaseous substances, the chemical potentials of these condensed substances at a constant temperature will be constant, as well as the saturation vapor pressure over each of these substances in the mixture. Therefore, the expressions for the equilibrium constants include the pressures of only gaseous substances.
Sao tv. + CO 2 gas CaCO 3 tv. (22)
If solutions take part in a heterogeneous reaction, the expressions for the equilibrium constants will include the activities (concentrations) of these substances.
3.6 Le Chatelier's principle.
As noted earlier, chemical equilibria are dynamic and mobile. Changes in external conditions can lead to a shift in equilibrium towards the formation of either reaction products or starting materials. For the first time, the principle of equilibrium shift was formulated by A. Le Chatelier.
Le Chatelier's principle: if an external influence is exerted on a system that is in equilibrium, then the equilibrium is shifted towards the process that reduces this impact.
Theoretically, this principle was put forward by F. Brown and is now known as the Le Chatelier-Brown principle.
Usually, the influence of temperature, pressure, and changes in the concentration of reactants are considered as external factors affecting the state of equilibrium. For example, with an increase in temperature, the equilibrium shifts towards an endothermic reaction proceeding with the absorption of heat. A decrease in pressure leads to a shift in equilibrium towards the reaction proceeding with an increase in the number of moles of gaseous substances.
Example 3 .2 Consider the reaction
(
)
Let's increase the temperature: since the reaction is exothermic, that is, it proceeds with the release of heat, the equilibrium shifts towards the starting materials (the reverse reaction proceeds with the absorption of heat).
Let's increase the pressure: since the direct reaction proceeds with a decrease in the number of moles of gaseous substances (i.e., the volume decreases), the equilibrium shifts towards the reaction products.
Chemical equilibrium constant- a characteristic of a chemical reaction, by the value of which one can judge the direction of the process at the initial ratio of the concentrations of the reactants, the maximum possible yield of the reaction product under certain conditions.
The chemical equilibrium constant is determined by the law of mass action. Its values are calculated or based on experimental data. The chemical equilibrium constant depends on the nature of the reactants and on the temperature.
Equilibrium constant and Gibbs energy
The equilibrium constant ~K is related to the Gibbs free energy ~\Delta G as follows:
~\Delta G=-RT\cdot\ln K.The above equation makes it possible to calculate K from the value of ΔG°, and then the equilibrium concentrations (partial pressures) of the reagents.
It can be seen from this equation that the equilibrium constant is very sensitive to changes in temperature (if we express the constant from here, then the temperature will be in the exponent). For endothermic processes, an increase in temperature corresponds to an increase in the equilibrium constant, for exothermic processes, to its decrease. The equilibrium constant does not depend on pressure, except for cases of very high pressure (from 100 Pa).
The dependence of the equilibrium constant on the enthalpy and entropy factors indicates the influence of the nature of the reagents on it.
Equilibrium constant and reaction rate
You can express the equilibrium constant in terms of the reaction rate. In this case, the equilibrium constant is defined as
~K=\frac(k_1)(k_(-1)),where ~k_1 is the rate constant of the forward reaction, ~k_(-1) is the rate constant of the reverse reaction.
An arbitrary reversible chemical reaction can be described by an equation of the form:
aA + bB Û dD + eE
In accordance with the law of mass action, in the simplest case, the rate of a direct reaction is related to the concentrations of the starting substances by the equation
v pr = k pr C A a FROM in b,
and the rate of the reverse reaction - with the concentrations of the products by the equation
v arr = k arr C D d FROM E e .
When equilibrium is reached, these speeds are equal to each other:
v pr = v arr
The ratio of the rate constants of the forward and reverse reactions to each other will be equal to equilibrium constant:
Since this expression is based on taking into account the amount of reactants and reaction products, it is a mathematical notation of the law acting masses for reversible reactions.
The equilibrium constant, expressed in terms of the concentrations of the reactants, is called the concentration constant and is denoted K s . For a more rigorous consideration, instead of concentrations, one should use the thermodynamic activities of substances but = fC (where f - activity coefficient). In this case, we are talking about the so-called thermodynamic equilibrium constant
At low concentrations, when the activity coefficients of the starting substances and products are close to unity, K s And K a practically equal to each other.
The equilibrium constant of a reaction occurring in the gas phase can be expressed in terms of partial pressures R substances involved in the reaction:
Between K r And K s there is a relation that can be derived in this way. We express the partial pressures of substances in terms of their concentrations using the Mendeleev-Clapeyron equation:
pV = nRT ,
where p = (n /V )RT = CRT .
Then for the reaction in general form, after replacing partial pressures by concentrations, we obtain
Replacing the expression (d + c) - (a + b) with its equal Dn , we get the final expression
K r = K s (RT )D n or K s = K r (RT ) - D n ,
where Dn - change in the number of moles of gaseous substances during the reaction:
Dn = ån i prod (g) - ån i ref (d) ).
If Dn = 0, i.e. the process goes on without changing the number of moles of gaseous substances, and K r = K s .
For example, for the ethylene hydration reaction occurring in the gas phase:
C 2 H 4 (g) + H 2 O (g) Û C 2 H 5 OH (g),
In this case Dn = 1 - (1 + 1) = -1. Hence, the relation between the constants can be expressed by the following equation:
K r = K s (RT ) - 1 or K s = K r RT .
Thus, knowing K r this reaction at any given temperature, you can calculate the value K s and vice versa.
The dimension of the equilibrium constants depends on the method of expressing the concentration (pressure) and the stoichiometry of the reaction. Often it can be bewildering, for example, in the considered example [mol - 1 m 3] for K s and [Pa - 1] for K r , but there is nothing wrong with that. If the sums of the stoichiometric coefficients of products and starting materials are equal, the equilibrium constant will be dimensionless.
In 1885, the French physicist and chemist Le Chatelier was deduced, and in 1887 by the German physicist Braun, the law of chemical equilibrium and the chemical equilibrium constant were substantiated, and their dependence on the influence of various external factors was studied.
The essence of chemical equilibrium
Equilibrium is a state that means things are always moving. Products are decomposed into reagents, and reagents are combined into products. Things move, but concentrations remain the same. The reaction is written with a double arrow instead of an equals sign to show that it is reversible.
Classic patterns
Back in the last century, chemists discovered certain patterns that provide for the possibility of changing the direction of the reaction in the same container. Knowing how chemical reactions work is incredibly important for both laboratory research and industrial production. At the same time, the ability to control all these phenomena is of great importance. It is human nature to intervene in many natural processes, especially reversible ones, in order to later use them for their own benefit. From knowledge of chemical reactions will be more useful if you are fluent in the levers of controlling them.
The law of mass action in chemistry is used by chemists to correctly calculate the rates of reactions. It gives a clear idea that none will be completed if it takes place in a closed system. The molecules of the resulting substances are in constant and random motion, and a reverse reaction may soon occur, in which the molecules of the starting material will be restored.
In industry, open systems are most often used. Vessels, apparatus and other containers where chemical reactions take place remain unlocked. This is necessary so that during these processes it is possible to extract the desired product and get rid of useless reaction products. For example, coal is burned in open furnaces, cement is produced in open furnaces, blast furnaces operate with a constant supply of air, and ammonia is synthesized by continuously removing ammonia itself.
Reversible and irreversible chemical reactions
Based on the name, one can give the appropriate definitions: irreversible reactions are those that are brought to an end, do not change their direction and proceed along a given trajectory, regardless of pressure drops and temperature fluctuations. Their distinguishing feature is that some products may leave the reaction sphere. Thus, for example, it is possible to obtain gas (CaCO 3 \u003d CaO + CO 2), a precipitate (Cu (NO 3) 2 + H 2 S \u003d CuS + 2HNO 3) or others will also be considered irreversible if a large amount is released during the process thermal energy, for example: 4P + 5O 2 \u003d 2P 2 O 5 + Q.
Almost all reactions that occur in nature are reversible. Regardless of such external conditions as pressure and temperature, almost all processes can proceed simultaneously in different directions. As the law of mass action in chemistry says, the amount of heat absorbed will be equal to the amount released, which means that if one reaction was exothermic, then the second (reverse) will be endothermic.
Chemical equilibrium: chemical equilibrium constant
Reactions are the "verbs" of chemistry - the activities that chemists study. Many reactions go to their completion and then stop, which means that the reactants are completely converted into products, with no way to return to their original state. In some cases, the reaction is indeed irreversible, for example, when combustion changes both physical and chemical. However, there are many other circumstances in which it is not only possible, but also continuous, since the products of the first reaction become reactants in the second.
The dynamic state in which the concentrations of reactants and products remain constant is called equilibrium. It is possible to predict the behavior of substances with the help of certain laws that are applied in industries seeking to reduce the cost of producing specific chemicals. The concept of chemical equilibrium is also useful in understanding processes that maintain or potentially threaten human health. The chemical equilibrium constant is the value of a reaction factor that depends on ionic strength and temperature and is independent of the concentrations of reactants and products in solution.
Calculation of the equilibrium constant
This value is dimensionless, that is, it does not have a certain number of units. Although the calculation is usually written for two reactants and two products, it works for any number of reaction participants. The calculation and interpretation of the equilibrium constant depends on whether the chemical reaction is associated with a homogeneous or heterogeneous equilibrium. This means that all reacting components can be pure liquids or gases. For reactions that reach heterogeneous equilibrium, as a rule, not one phase is present, but at least two. For example, liquids and gases or and liquids.
The value of the equilibrium constant
For any given temperature, there is only one value for the equilibrium constant, which only changes if the temperature at which the reaction occurs changes in one direction or another. Some predictions about a chemical reaction can be made based on whether the equilibrium constant is large or small. If the value is very large, then the equilibrium favors the reaction to the right and more products are obtained than there were reactants. The reaction in this case can be called "total" or "quantitative".
If the value of the equilibrium constant is small, then it favors the reaction to the left, where the amount of reactants was greater than the number of products formed. If this value tends to zero, we can assume that the reaction does not occur. If the values of the equilibrium constant for the direct and reverse reactions are almost the same, then the amount of reactants and products will also be almost the same. This type of reaction is considered to be reversible.
Consider a specific reversible reaction
Take two such chemical elements as iodine and hydrogen, which, when mixed, give a new substance - hydrogen iodide.
For v 1 we take the rate of the direct reaction, for v 2 - the rate of the reverse reaction, k - the equilibrium constant. Using the law of mass action, we obtain the following expression:
v 1 \u003d k 1 * c (H 2) * c (I 2),
v 2 = k 2 * c 2 (HI).
When mixing iodine (I 2) and hydrogen (H 2) molecules, their interaction begins. At the initial stage, the concentration of these elements is maximum, but by the end of the reaction, the concentration of a new compound, hydrogen iodide (HI), will be maximum. Accordingly, the reaction rates will also be different. At the very beginning, they will be maximum. Over time, there comes a moment when these values are equal, and this is the state called chemical equilibrium.
The expression of the chemical equilibrium constant, as a rule, is denoted using square brackets: , , . Since at equilibrium the speeds are equal, then:
k 1 \u003d k 2 2,
so we get the equation of the chemical equilibrium constant:
k 1 /k 2 = 2 / = K.
Le Chatelier-Brown principle
There is the following regularity: if a certain effect is made on a system that is in equilibrium (change the conditions of chemical equilibrium by changing temperature or pressure, for example), then the balance will shift in order to partially counteract the effect of the change. In addition to chemistry, this principle also applies in slightly different forms to the fields of pharmacology and economics.
Chemical equilibrium constant and ways of its expression
The equilibrium expression can be expressed in terms of the concentration of products and reactants. Only chemicals in the aqueous and gaseous phases are included in the equilibrium formula because the concentrations of liquids and solids do not change. What factors affect chemical equilibrium? If a pure liquid or solid is involved in it, it is considered that it has K \u003d 1, and accordingly ceases to be taken into account, with the exception of highly concentrated solutions. For example, pure water has an activity of 1.
Another example is solid carbon, which can be formed by the reaction of two molecules of carbon monoxide to form carbon dioxide and carbon. Factors that can affect the balance include the addition of a reactant or product (changes in concentration affect the balance). The addition of a reactant can bring equilibrium to the right in the chemical equation, where more forms of the product appear. The addition of product can bring equilibrium to the left as more reactant forms become available.
Equilibrium occurs when a reaction proceeding in both directions has a constant ratio of products and reactants. In general, the chemical equilibrium is static, since the quantitative ratio of products and reactants is constant. However, a closer look reveals that equilibrium is actually a very dynamic process, as the reaction moves in both directions at the same rate.
Dynamic equilibrium is an example of a steady state function. For a system at steady state, the currently observed behavior continues into the future. Therefore, once the reaction reaches equilibrium, the ratio of product to reactant concentrations will remain the same even though the reaction continues.
How easy is it to talk about complex things?
Concepts such as chemical equilibrium and chemical equilibrium constant are quite difficult to understand. Let's take an example from life. Have you ever been stuck on a bridge between two cities and noticed that the traffic in the other direction is smooth and measured while you are hopelessly stuck in traffic? This is not good.
What if the cars were measured and at the same speed moving on both sides? Would the number of cars in both cities remain constant? When the speed of entry and exit to both cities is the same, and the number of cars in each city is stable over time, this means that the whole process is in dynamic equilibrium.
Chemical equilibrium is a state of a chemical system in which one or more chemical reactions reversibly proceed, and the rates in each pair of forward-reverse reactions are equal to each other. For a system in chemical equilibrium, the concentrations of reagents, temperature and other parameters of the system do not change with time
A2 + B2 ⇄ 2AB
A quantitative characteristic of chemical equilibrium is a quantity called the constant of chemical equilibrium.
At a constant temperature, the equilibrium constant of a reversible reaction is a constant value showing the ratio between the concentrations of the reaction products and starting substances, which is established at equilibrium.
The equilibrium constant equation shows that under equilibrium conditions, the concentrations of all substances involved in the reaction are interconnected. A change in the concentration of any of these substances entails a change in the concentration of all other substances. As a result, a new concentration is established, but the ratio between them corresponds to the equilibrium constant.
58. Factors that determine the direction of chemical reactions. Chemical processes must proceed in the direction of decreasing the internal energy of the system, i.e. in the direction corresponding to the positive thermal effect of the reaction.
The second factor influencing the direction of chemical reactions is the principle of the direction of processes to the most probable state, i.e. in chemical reactions, due to the principle of directing processes to a minimum of internal energy, atoms combine into molecules, during the formation of which the greatest amount of energy is released.
The tendency to transition to the state with the lowest internal energy is manifested at the temperature to the same extent. The tendency to achieve the most probable state becomes stronger the higher the temperature. At low temperatures, in most cases, only the influence of the first of these tendencies has a practical effect, as a result of which exothermic processes proceed spontaneously. As the temperature rises, the equilibrium in chemical systems shifts more and more towards the reaction of decomposition or an increase in the number of states of atoms. In this case, each temperature corresponds to an equilibrium state, characterized by a certain ratio of the concentration of reactants and reaction products
59. Displacement of chemical equilibrium. La Chatelier's principle. If the system is in a state of equilibrium, then it will remain in it as long as the external conditions remain constant. Of greatest importance are cases of violation of chemical equilibrium due to a decrease in the concentration of any of the substances participating in the equilibrium; change in pressure and temperature. These imbalances are governed by the Le Chatelier principle: if a system in equilibrium is affected, then as a result of the processes occurring in it, the equilibrium will shift in such a direction that the impact will decrease.
60. Gibbs phase rule. For any system in equilibrium, the sum of the number of phases (P) and the number of possible states of the system (V) is greater than the number of components (C) by 2: P + V = C + 2
61. Solutions. dissolution process. A solution is a solid, gaseous or liquid homogeneous system consisting of two or more components, the amount of which can vary over a wide range. The most important type of solutions is liquid. Any solution consists of solutes and a solvent, i.e. environment in which these substances are evenly distributed in the form of molecules or ions. Usually, a solvent is considered to be that component that exists in its pure form in the same state of aggregation as the resulting solution.
The homogeneity of solutions makes them very similar to chemical compounds.
The difference between a solution and chemical compounds is that the composition of a solution can vary over a very wide range. In addition, many properties of its individual components can be found in the properties of the solution, which is not observed in the case of chemical compounds.
The variability of the composition of solutions brings them closer to mechanical mixtures, but they differ sharply from them in their uniformity. Solutions occupy an intermediate position between mechanical mixtures and chemical compounds.
The separation of molecules from the surface of the crystal during dissolution is caused, on the one hand, by their own thermal vibrations of the molecules, and, on the other hand, by the attraction of molecules by the solvent.
A solution in equilibrium with a solute is called a saturated solution.
62. Ways of expressing the composition of the solution. a) Mass fraction: ω \u003d m 1 / (m 1 + m 2) * 100% where m 1 is the dissolved substance; m 1 + m 2 - the mass of the solution; m 2 is the mass of the solvent;
b) Mole fraction N= ν 1 / ν 1 + ν 2 - this is the ratio of the number of moles of a solute to the sum of the amount of all substances that make up the solution;
c) Molar concentration C \u003d V 1 / m 2 - the ratio of the amount of substance contained in the solution to the volume of the solution (mol / l);
d) Mole concentration C \u003d V e1 / V - the ratio of the amount of substance contained in the solution to the mass of the solvent (mol \ kg);
e) The molar concentration of the equivalent is the ratio of the amount of the equivalent substance contained in the solution to the volume of this solution (mol / l).
63. Solubility, Henry's Law. Solubility is the ability of a substance to dissolve in a particular solvent. A measure of the solubility of a substance under given conditions is its content in a saturated solution. However, as a rule, substances consisting of polar molecules and substances with an ionic type of bond dissolve better in polar solvents (water, alcohol, ammonia), and non-polar substances in non-polar solvents (benzene, etc.). If the dissolution of a solid in a liquid is accompanied by the absorption of heat, then an increase in temperature leads to an increase in the solubility of the substance.
When solids are dissolved, the volume of the system usually changes insignificantly. Therefore, the solubility of solids in liquids does not depend on pressure.
Henry's law: the mass of a gas that dissolves at a constant temperature in a given volume of liquid is directly proportional to the partial pressure of the gas. C \u003d cr, where C is the mass-volume concentration, p is the partial gas pressure, k is the Henry coefficient.
Consequence of Henry's Law:
a) The volume of a gas dissolved at a constant temperature in a given volume of liquid does not depend on its partial pressure.
b) If there is a mixture of gases above the liquid, then the dissolution of each of them is determined by the partial pressure.
64. Law of distribution. Extraction. Distribution law: a substance that can dissolve in 2 immiscible solvents is distributed between them so that the ratio of its concentrations in these solvents at a constant temperature remains constant, regardless of the total amount of the dissolved substance.
Extraction is a method of extracting a substance from a solution using a suitable solvent (extractant). For extraction from a solution, solvents are used that are immiscible with this solution, but in which the substance dissolves better than in the first solvent.
65. Osmosis. Van't Hoff's law. Osmosis is the one-way diffusion of molecules through a semi-permeable membrane.
When measuring the osmotic pressure of various solutions, it was found that the magnitude of the osmotic pressure depends on the concentration of the solution and on its temperature, but does not depend on the nature of the dissolved substances and the solvent.
P=CRT - van't Hoff's law
where P is the osmotic pressure of the solution (Pa), C is the molarity, R is the universal gas constant, T is the absolute temperature
66. Vapor pressure of solution. Raul's Law. At a given temperature, the vapor pressure over the liquid is a constant value. When a substance dissolves in a liquid, the vapor pressure of the liquid decreases.
A vapor in equilibrium with its liquid is said to be saturated.
Saturated vapor pressure depends on the nature of the liquid and temperature, but does not depend on the volume of the vessel in which the vapor is located.
Thus, the saturation vapor pressure of a solvent over a solution is always lower than over a pure solvent at the same temperature.
The difference between the saturation vapor pressure over the pure solvent and over the solution is called the vapor pressure drop of the solution, and the ratio of the vapor pressure drop of the solution to the saturation vapor pressure over the pure solvent is called the relative vapor pressure drop over the solution.
Raoult's law: the relative decrease in the pressure of the saturated vapor of the solvent over the solution is equal to the mole fraction of the solute.
The phenomenon of a decrease in the pressure of saturated vapor over a solution follows from the Le Chatelier principle.
67. Aqueous solutions of electrolytes. Theory of electrolytic dissociation. a) Dissociation of salts, i.e. crystals with ionic structure.
b) Dissociation upon dissolution of acids, i.e. polar molecules.
Theory of electrolytic dissociation.
Electrolytic dissociation is the process of decomposition of an electrolyte into ions during its dissolution or melting.
The classical theory of electrolytic dissociation was created by S. Arrhenius and W. Ostwald in 1887. Arrhenius adhered to the physical theory of solutions, did not take into account the interaction of electrolyte with water, and believed that free ions were present in solutions. Russian chemists I. A. Kablukov and V. A. Kistyakovsky used the chemical theory of solutions of D. I. Mendeleev to explain the electrolytic dissociation and proved that when the electrolyte is dissolved, it chemically interacts with water, as a result of which the electrolyte dissociates into ions.
The classical theory of electrolytic dissociation is based on the assumption of incomplete dissociation of a solute, characterized by the degree of dissociation α, i.e. fraction of decomposed electrolyte molecules. Dynamic equilibrium between non-dissociated molecules and ions is described by the law of mass action.
68. Strong and weak electrolytes. Degree of dissociation. The degree of dissociation of an electrolyte is the ratio of the number of molecules decomposed into ions in a given solution to the total number of molecules of a given substance in a solution.
Electrolytes, the degree of dissociation of which tends to 1 are called strong: NaCl, NaOH, HCl.
Electrolytes, the degree of dissociation of which tends to 0, are called weak: H 2 O, H 2 CO 2, NH 4 OH.
69. Dissociation constant. Ostwald's dilution law. It is possible to apply the laws that are valid for chemical equilibrium to the equilibria that are established during the dissociation of a weak electrolyte.
The equilibrium constant corresponding to the dissociation of a weak electrolyte is called the dissociation constant.
The value of the equilibrium constant depends on the nature of the electrolyte and solvent, on temperature, but does not depend on the concentration of the solution. This value characterizes the ability of a given acid, base or salt to decompose into ions. The higher the value of the equilibrium constant, the easier the electrolyte dissociates into ions.
Ostwald's law - the degree of dissociation increases as the electrolyte is diluted.
70. The state of strong electrolytes in solution. Activity. Ionic strength. To assess the state of ions in solutions of strong electrolytes, a quantity called activity is used. The activity of an ion is understood as that effective conditional concentration of it, in accordance with which it acts in chemical reactions:
where a is the ion activity, c is the ion concentration, f is the activity coefficient.
For the dilution of solutions, an expression is valid that relates the activity coefficient and the value of the ionic strength of the solution.
lgF= - 0.5Z 2 square root I
If we use activity values, then the laws of chemical equilibrium can also be applied to solutions of strong electrolytes.
71. Properties of acids, bases and salts from the point of view of the theory of electrolytic dissociation. Acids are able to interact with bases. This produces salt and water.
The theory of electrolytic dissociation defines acids as electrolytes that dissociate to form positively charged hydrogen ions.
Similarly, bases are defined as electrolytes that dissociate to form a negatively charged hydroxide ion.
Salts: Hydroxide and hydrogen ions are not formed. They are considered as electrolytes that dissociate to form positively charged ions other than hydrogen ions and negatively charged ions other than hydroxide ion.