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.
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.
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.
1 . Determine the chemical equilibrium constant of the process of dissociation of phosphorus pentachloride, proceeding according to the equation PCl 5<=>PCl 3 + Cl 2,
according to the equilibrium concentrations of K C and the partial pressures of the components K R. It is known that by the time equilibrium was established at a temperature of 500 K, 54% of PCl 5 had reacted with respect to the initial concentration of 1 mol/l.
Solution. According to the law of mass action for chemical equilibrium
K C = :.
Applying this expression to our equation, we get
K C \u003d C μ (PCl 3) C μ (Cl 2) / C μ (PCl 5).
This expression includes the equilibrium concentrations of the reactant and products. Let's calculate them.
If by the time equilibrium is established, 54% of phosphorus pentachloride has reacted, then 46% is still left in the system, which means that the equilibrium concentration of PCl 5 will be: C μ (PCl 5) = C 0 (PCl 5) 0.46 = 1 0, 46 = 0.46 (mol/l).
Before the start of the reaction, there were no products in the reaction space; therefore, their concentration in the equilibrium state is determined by the proportion of the spent reagent (see the reaction equation), i.e. C μ (PCl 3) \u003d C μ (Cl 2) \u003d 0.54 mol / l.
Let's substitute the found values into the expression for К С and make calculations:
K C \u003d (0.54) (0.54) / (0.46) \u003d 0.63 (mol / l).
The equilibrium constant in relation to the partial pressures of the gas system is determined by the equation: K R = K C (RT) Δ n, while, Δn = Σn prod - Σn react. It can be seen from the chemical equation that Δn = 2 - 1 = 1 mol.
We substitute the found and known values \u200b\u200binto the expression for K P and make calculations: K P \u003d (0.63) (8.31 500) 1 \u003d 2617 (Pa / l K) \u003d 2.62 (kPa / l K ).
Answer: K C = 0.63 mol/l; K P \u003d 2.62 kPa / l K.
2 . In a state of system equilibrium
CO 2 (g) + H 2 (g)<=>CO (g) + H 2 O (g)
the reaction mixture had the following composition by volume: 22% CO 2 , 41% H 2 , 17% CO and 20% H 2 O. Determine K P and K C for this equilibrium at 1900 K and a pressure of 98,501 Pa.
Solution. From the volume fractions of each gas in a mixture, one can determine their partial pressures, knowing that the total pressure in any gas mixture is equal to the sum of the partial pressures of the gases that make up this mixture:
P \u003d P (CO 2) + P (H 2) + P (CO) + P (H 2 O).
And then P (CO 2) \u003d 0.22 P \u003d 0.22 98 501 \u003d 21 670 (Pa);
P (H 2) \u003d 0.41 98 501 \u003d 40 385 (Pa);
Р(СО) = 0.17 98 501 = 16 745 (Pa); P (H 2 O) \u003d 0.2 98 501 \u003d 19 700 (Pa).
We introduce the found values of the partial pressures of gases into an equation that establishes the dependence of the equilibrium constant on the partial pressures of the process components:
K P \u003d P (CO) P (H 2 O) / P (CO 2) P (H 2);
K P \u003d 16745 19700 / 21670 40385 \u003d 0.38.
The equilibrium constant in relation to the equilibrium concentrations of reagents K C is calculated by the equation K C = K R ·(RT) -Δ n . For our case, Δn \u003d 2 - 2 \u003d 0, therefore, the factor (RT) -Δ n \u003d (RT) 0 \u003d 1. and then K C \u003d K P \u003d 0.38.
Answer: K C = K P = 0.38.
3 . Equilibrium in a homogeneous system
4HCl (g) + O 2 (g)<=>2H 2 O (g) + 2Cl 2 (g)
established at the following concentrations of reactants (mol/l): C μ (H 2 O) = 0.14; C μ (Cl 2) = 0.14; C μ (HCl) = 0.20; C μ (O 2) = 0.32. Calculate the equilibrium constant of this reaction and determine the initial concentrations of hydrogen chloride and oxygen.
Solution. According to the law of mass action, the chemical equilibrium constant for the concentrations of reagents is determined in relation to our case by the equation:
K C \u003d [C μ 2 (H 2 O) C μ 2 (Cl 2)] / [C μ 4 (HCl) C μ (O 2)].
Substituting the values of equilibrium concentrations into it and performing calculations, we obtain:
K C \u003d (0.14) 4 / (0.20) 4 (0.32) \u003d 0.75 (l / mol).
The chemical equation of the reaction shows in what proportions the substances interact with each other: from 4 mol of HCl and 1 mol of O 2, 2 mol of H 2 O and Cl 2 are formed. This means that ΔС μ (HCl) = 2 0.14 = 0.28 mol / l and ΔС μ (О 2) = ½ 0.14 = 0.07 mol / l .
The initial concentrations of hydrogen chloride and oxygen will be equal to:
C μ 0 (HCl) \u003d C μ (HCl) + ΔC μ (HCl), C μ 0 (HCl) \u003d (0.20 + 0.28) mol / l \u003d 0.48 mol / l;
C μ 0 (O 2) \u003d C μ (O 2) + ΔC μ (O 2), C μ 0 (O 2) \u003d (0.32 + 0.07) mol / l \u003d 0.39 mol / l.
Answer: equilibrium constant of this reaction K C = 0.75 l / mol; initial concentrations of reagents С μ 0 (HCl) = 0.48 mol/l; C μ 0 (O 2) = 0.39 mol / l.
4 . Determine the equilibrium constant of the reaction
2NO(g) + Cl 2 (g)<=>2NOCl(g)
at a temperature of 298 K according to the values of the standard enthalpies of formation and the entropies of its participants.
Solution. The dependence of the equilibrium constant of a chemical reaction on temperature is determined according to the Kirchhoff equation:
ΔH˚ equal ΔS˚ equal
ℓnК Р = – – + – .
We calculate ΔH˚ equal and ΔS˚ equal under standard conditions, using the reference data of the table in Appendix No. 1:
ΔH˚ equals = 2ΔН 0 (NOCl) - 2ΔН 0 (NO),
ΔS˚ equals = 2S 0 (NOCl) - 2S 0 (NO) - S 0 (Cl 2), then
ΔH˚ equals = 2 53.55 kJ / mol - 2 90.37 kJ / mol \u003d - 73.64 kJ / mol \u003d - 73640 J / mol;
ΔS˚ equals = 2 263.6 (J / mol K) - 2 210.62 (J / mol K) - 223.0 (J / mol K) \u003d - 117.04 J / mol K .
We substitute the found values into the Kirchhoff equation and perform the calculations:
ℓnK P \u003d - (- 73640) / 8.31 298 + (-117.04 / 8.31) \u003d 29.74 - 14.08 \u003d 15.66.
Then K R \u003d e 15.66 \u003d 5.7 10 6.
Answer: equilibrium constant of the reaction K P = 5.7 10 6 .
5 . For the reaction H 2 (g) + Br 2 (g)<=>2HBr(g) at some temperature, the equilibrium constant is equal to one. Determine the composition (in percent by volume) of the equilibrium gas reaction mixture if the initial mixture consisted of 3 moles of hydrogen and 2 moles of bromine.
Solution. Let us write down the expression of the law of mass action for the equilibrium state of the system:
Kp \u003d C 2 μ (HBr) / C μ (H 2) C μ (Br 2) \u003d 1.
Let us assume that by the time equilibrium is established in the system, it has reacted according to X moles of H 2 and Br 2. Then their equilibrium concentrations will be equal: C μ (H 2) = (3 - X) mol, C μ (Br 2) = (2 - X) mol. Accordingly, by this moment, 2 X mole of hydrogen bromide, i.e. With μ (HBr) = 2 X mol.
Substituting these values of equilibrium concentrations into the expression for the equilibrium constant, we obtain an equation with one unknown: (2 X) 2 /(3-X)·(2- X) = 1.
Deciding it about X, we get: 3 X 2 - 5X - 6 = 0;
X 1.2 = (-5±√25+72)/6 = (-5±10)/6. From here X= 0.75 mol.
Thus, at the time of equilibrium, the mixture contained: 2.25 mol H 2 , 1.25 mol Br 2 and 1.50 mol HBr, which corresponds to 45% hydrogen, 25% bromine and 30% hydrogen bromide.
Answer: at equilibrium, the gas mixture included 45% H 2 , 25% Br 2 and 30% HBr.
Most chemical reactions are reversible, i.e. flow simultaneously in opposite directions. In cases where the forward and reverse reactions proceed at the same rate, chemical equilibrium occurs. For example, in a reversible homogeneous reaction: H 2 (g) + I 2 (g) ↔ 2HI (g), the ratio of the rates of direct and reverse reactions according to the law of mass action depends on the ratio of the concentrations of the reactants, namely: the rate of the direct reaction: υ 1 = k 1 [Н 2 ]. The rate of the reverse reaction: υ 2 \u003d k 2 2.
If H 2 and I 2 are the initial substances, then at the first moment the rate of the forward reaction is determined by their initial concentrations, and the rate of the reverse reaction is zero. As H 2 and I 2 are consumed and HI is formed, the rate of the forward reaction decreases and the rate of the reverse reaction increases. After some time, both velocities are equalized, and chemical equilibrium is established in the system, i.e. the number of formed and consumed HI molecules per unit time becomes the same.
Since at chemical equilibrium the rates of direct and reverse reactions are equal to V 1 \u003d V 2, then k 1 \u003d k 2 2.
Since k 1 and k 2 are constant at a given temperature, their ratio will be constant. Denoting it by K, we get:
K - is called the constant of chemical equilibrium, and the above equation is called the law of mass action (Guldberg - Vaale).
In the general case, for a reaction of the form аА+bB+…↔dD+eE+…, the equilibrium constant is equal to 
. For the interaction between gaseous substances, the expression is often used, in which the reactants are represented by equilibrium partial pressures p. For the mentioned reaction 
.
The state of equilibrium characterizes the limit to which, under given conditions, the reaction proceeds spontaneously (∆G<0). Если в системе наступило химическое равновесие, то дальнейшее изменение изобарного потенциала происходить не будет, т.е. ∆G=0.
The ratio between the equilibrium concentrations does not depend on which substances are taken as starting materials (for example, H 2 and I 2 or HI), i.e. equilibrium can be approached from both sides.
The chemical equilibrium constant depends on the nature of the reactants and on the temperature; the equilibrium constant does not depend on pressure (if it is too high) and on the concentration of reagents.
Influence on the equilibrium constant of temperature, enthalpy and entropy factors. The equilibrium constant is related to the change in the standard isobaric-isothermal potential of a chemical reaction ∆G o by a simple equation ∆G o =-RT ln K.
It shows that large negative values of ∆G o (∆G o<<0) отвечают большие значения К, т.е. в равновесной смеси преобладают продукты взаимодействия. Если же ∆G o характеризуется большими положительными значениями (∆G o >>0), then the initial substances predominate in the equilibrium mixture. This equation allows us to calculate K from the value of ∆G o and then the equilibrium concentrations (partial pressures) of the reagents. If we take into account that ∆G o =∆Н o -Т∆S o , then after some transformation we get 
. It can be seen from this equation that the equilibrium constant is very sensitive to changes in temperature. The influence of the nature of the reagents on the equilibrium constant determines its dependence on the enthalpy and entropy factors.
Le Chatelier's principle
The state of chemical equilibrium is maintained under these constant conditions at any time. When the conditions change, the state of equilibrium is disturbed, since in this case the rates of opposite processes change to different degrees. However, after some time, the system again comes to a state of equilibrium, but already corresponding to the new changed conditions.
The shift of equilibrium depending on changes in conditions is generally determined by the Le Chatelier principle (or the principle of moving equilibrium): if a system in equilibrium is influenced from outside by changing any of the conditions that determine the equilibrium position, then it is shifted in the direction of the process, the flow of which weakens the effect of the effect produced.
Thus, an increase in temperature causes a shift in equilibrium in the direction of that of the processes, the course of which is accompanied by the absorption of heat, and a decrease in temperature acts in the opposite direction. Similarly, an increase in pressure shifts the equilibrium in the direction of a process accompanied by a decrease in volume, and a decrease in pressure acts in the opposite direction. For example, in the equilibrium system 3H 2 + N 2 2H 3 N, ∆H o = -46.2 kJ, an increase in temperature enhances the decomposition of H 3 N into hydrogen and nitrogen, since this process is endothermic. An increase in pressure shifts the equilibrium towards the formation of H 3 N, because the volume decreases.
If a certain amount of any of the substances participating in the reaction is added to a system that is in equilibrium (or vice versa, removed from the system), then the rates of the forward and reverse reactions change, but gradually become equal again. In other words, the system again comes to a state of chemical equilibrium. In this new state, the equilibrium concentrations of all substances present in the system will differ from the initial equilibrium concentrations, but the ratio between them will remain the same. Thus, in a system in equilibrium, it is impossible to change the concentration of one of the substances without causing a change in the concentrations of all the others.
In accordance with the Le Chatelier principle, the introduction of additional amounts of a reagent into the equilibrium system causes a shift in the equilibrium in the direction in which the concentration of this substance decreases and, accordingly, the concentration of the products of its interaction increases.
The study of chemical equilibrium is of great importance both for theoretical research and for solving practical problems. By determining the equilibrium position for various temperatures and pressures, one can choose the most favorable conditions for conducting a chemical process. In the final choice of process conditions, their influence on the process rate is also taken into account.
Example 1 Calculation of the equilibrium constant of the reaction from the equilibrium concentrations of the reactants.
Calculate the equilibrium constant of the reaction A + B 2C, if the equilibrium concentrations [A] = 0.3 mol ∙ l -1; [B]=1.1 mol∙l -1; [C] \u003d 2.1 mol ∙ l -1.
Solution. The expression for the equilibrium constant for this reaction is: . Let us substitute here the equilibrium concentrations indicated in the condition of the problem: =5.79.
Example 2. Calculation of equilibrium concentrations of reactants. The reaction proceeds according to the equation A + 2B C.
Determine the equilibrium concentrations of the reactants if the initial concentrations of substances A and B are respectively 0.5 and 0.7 mol∙l -1, and the equilibrium constant of the reaction is K p =50.
Solution. For each mole of substances A and B, 2 moles of substance C are formed. If the decrease in the concentration of substances A and B is denoted by X mol, then the increase in the concentration of the substance will be 2X mol. The equilibrium concentrations of the reactants will be:
C A \u003d (o.5-x) mol ∙ l -1; C B \u003d (0.7-x) mol ∙ l -1; C C \u003d 2x mol ∙ l -1
x 1 \u003d 0.86; x 2 \u003d 0.44
According to the condition of the problem, the value x 2 is valid. Hence, the equilibrium concentrations of the reactants are:
C A \u003d 0.5-0.44 \u003d 0.06 mol ∙ l -1; C B \u003d 0.7-0.44 \u003d 0.26 mol ∙ l -1; C C \u003d 0.44 ∙ 2 \u003d 0.88 mol ∙ l -1.
Example 3 Determination of the change in the Gibbs energy ∆G o of the reaction by the value of the equilibrium constant K p. Calculate the Gibbs energy and determine the possibility of the reaction CO+Cl 2 =COCl 2 at 700K, if the equilibrium constant is Kp=1.0685∙10 -4. The partial pressure of all reacting substances is the same and equal to 101325 Pa.
Solution.∆G 700 =2.303∙RT 
.
For this process:
Since ∆Go<0, то реакция СО+Cl 2 COCl 2 при 700К возможна.
Example 4. Shift in chemical equilibrium. In which direction will the equilibrium shift in the N 2 + 3H 2 2NH 3 -22 kcal system:
a) with an increase in the concentration of N 2;
b) with an increase in the concentration of H 2;
c) when the temperature rises;
d) when the pressure decreases?
Solution. An increase in the concentration of substances on the left side of the reaction equation, according to the Le Chatelier rule, should cause a process that tends to weaken the effect, lead to a decrease in concentrations, i.e. the equilibrium will shift to the right (cases a and b).
The ammonia synthesis reaction is exothermic. An increase in temperature causes a shift in equilibrium to the left - towards an endothermic reaction that weakens the impact (case c).
A decrease in pressure (case d) will favor the reaction leading to an increase in the volume of the system, i.e. towards the formation of N 2 and H 2 .
Example 5 How many times will the rate of forward and reverse reactions in the system 2SO 2 (g) + O 2 (g) 2SO 3 (r) change if the volume of the gas mixture decreases three times? In which direction will the equilibrium of the system shift?
Solution. Let us denote the concentrations of reacting substances: = a, =b,=With. According to the law of mass action, the rates of the forward and reverse reactions before a change in volume are
v pr \u003d Ka 2 b, v arr \u003d K 1 s 2
After reducing the volume of a homogeneous system by a factor of three, the concentration of each of the reactants will increase by a factor of three: 3a,[O 2] = 3b; = 3s. At new concentrations of the rate v "np of the direct and reverse reactions:
v" np = K(3a) 2 (3b) = 27 Ka 2 b; v o 6 p = K 1 (3c) 2 = 9K 1 c 2 .
; 
Consequently, the rate of the forward reaction increased 27 times, and the reverse - only nine times. The equilibrium of the system has shifted towards the formation of SO 3 .
Example 6 Calculate how many times the rate of the reaction proceeding in the gas phase will increase with an increase in temperature from 30 to 70 0 C, if the temperature coefficient of the reaction is 2.
Solution. The dependence of the rate of a chemical reaction on temperature is determined by the Van't Hoff empirical rule according to the formula
Therefore, the reaction rate at 70°C is 16 times greater than the reaction rate at 30°C.
Example 7 The equilibrium constant of a homogeneous system
CO (g) + H 2 O (g) CO 2 (g) + H 2 (g) at 850 ° C is 1. Calculate the concentrations of all substances at equilibrium if the initial concentrations are: [CO] ISC = 3 mol / l, [H 2 O] ISH \u003d 2 mol / l.
Solution. At equilibrium, the rates of the forward and reverse reactions are equal, and the ratio of the constants of these rates is constant and is called the equilibrium constant of the given system:
V np= K 1[CO][H 2 O]; V o b p = TO 2 [CO 2 ][H 2 ];
In the condition of the problem, the initial concentrations are given, while in the expression K r includes only the equilibrium concentrations of all substances in the system. Let us assume that by the moment of equilibrium the concentration [СО 2 ] Р = X mol/l. According to the equation of the system, the number of moles of hydrogen formed in this case will also be X mol/l. The same number of prayers (X mol / l) CO and H 2 O are consumed for the formation of X moles of CO 2 and H 2. Therefore, the equilibrium concentrations of all four substances (mol / l):
[CO 2] P \u003d [H 2] p \u003d X;[CO] P = (3 – x); P =(2-x).
Knowing the equilibrium constant, we find the value X, and then the initial concentrations of all substances:
; x 2 \u003d 6-2x-3x + x 2; 5x \u003d 6, l \u003d 1.2 mol / l.
Consider a reversible chemical reaction of a general form, in which all substances are in the same state of aggregation, for example, liquid:
aA + bB D cC + d D,
where A and B are the starting materials of the direct reaction; C and D are direct reaction products; a, b, c, and d- stoichiometric coefficients.
At the initial moment of time, when the concentration of substances A and B is the highest, the rate of the direct reaction will also be the highest and, according to the law of mass action, is equal to
u pr \u003d k 1 C A a C B in (6.1)
where k 1 is the rate constant of the direct reaction.
Over time, the concentration of substances A and B decreases, and, consequently, the rate of the direct reaction also decreases.
At the initial moment of time, the concentration of substances C and D is equal to zero, and, consequently, the rate of the reverse reaction is equal to zero, over time, the concentration of substances C and D increases, and, consequently, the rate of the reverse reaction also increases and it will be equal to
u arr \u003d k 2 C C with C D d (6.2)
where k 2 is the rate constant of the reverse reaction.
At the moment of reaching equilibrium, the concentrations take on the value of equilibrium, and the speeds are equal to each other u pr \u003d u arr, therefore
k 1 C A a C B c = k 2 C C c C D d (6.3)
Let's move the rate constants in one direction, and the concentrations in the other:
The ratio of two constants is a constant, and it is called the constant of chemical equilibrium:
The equilibrium constant shows how many times the rate of the forward reaction is greater or less than the rate of the reverse reaction.
Equilibrium constant is the ratio of the product of the equilibrium concentrations of the reaction products, taken to the power of their stoichiometric coefficients, to the product of the equilibrium concentrations of the starting materials, taken to the power of their stoichiometric coefficients.
The value of the equilibrium constant depends on the nature of the reacting substances and temperature, and does not depend on the concentration at the moment of equilibrium, since their ratio is always a constant value, numerically equal to the equilibrium constant. If a homogeneous reaction occurs between substances in solution, then the equilibrium constant is denoted by K C, and if between gases, then K P.
where Р С, Р D , Р А and Р В are the equilibrium pressures of the reaction participants.
Using the equation Clapeyron-Mendeleev, you can determine the relationship between K P and K C
Move the volume to the right side
p = RT, i.e. p = CRT (6.9)
We substitute equation (6.9) into (6.7), for each reagent and simplify
where Dn is the change in the number of moles of gaseous participants in the reaction
Dn = (s + d) - (a + c) (6.11)
Hence,
K P \u003d K C (RT) D n (6.12)
From equation (6.12) it can be seen that K P = K C, if the number of moles of gaseous participants in the reaction does not change (Dn = 0) or there are no gases in the system.
It should be noted that in the case of a heterogeneous process, the concentration of the solid or liquid phase in the system is not taken into account.
For example, the equilibrium constant for a reaction of the form 2A + 3B \u003d C + 4D, provided that all substances are gases and has the form
and if D is solid, then
The equilibrium constant is of great theoretical and practical importance. The numerical value of the equilibrium constant makes it possible to judge the practical possibility and depth of a chemical reaction.
If K > 1, then this reaction proceeds with a significant yield of reaction products; if K > 10 4 , then the reaction is irreversible; if K< 1, то такая реакция нетехнологична; если K < 10 -4 , то такая реакция невозможна.
Knowing the equilibrium constant, one can determine the composition of the reaction mixture at the moment of equilibrium and calculate the yield constant of the reaction products. The equilibrium constant can be determined using experimental methods, by analyzing the quantitative composition of the reaction mixture at the moment of equilibrium, or by applying theoretical calculations. For many reactions under standard conditions, the equilibrium constant is a tabular value.
Under external action on the system, the chemical equilibrium is shifted, i.e., the equilibrium concentrations of the initial substances and reaction products change. If, as a result of an external influence, the equilibrium concentrations of the reaction products increase, then one speaks of a shift of the equilibrium to the right (in the direction of the direct reaction). If, as a result of an external influence, the equilibrium concentrations of the starting substances increase, then one speaks of a shift of the equilibrium to the left (in the direction of the reverse reaction).
The influence of various factors on the shift in chemical equilibrium reflects Le Chatelier's principle (1884): if a system in stable chemical equilibrium is acted upon from the outside by changing the temperature, pressure or concentration, then the chemical equilibrium shifts in the direction in which the effect of the effect produced decreases.
It should be noted that the catalyst does not shift the chemical equilibrium, but only accelerates its onset.
Consider the influence of each factor on the shift of chemical equilibrium for a general reaction:
aA + bB = cC + d D±Q.
Effect of concentration change. According to the Le Chatelier principle, an increase in the concentration of one of the components of an equilibrium chemical reaction leads to a shift in the equilibrium towards an increase in the reaction in which the chemical processing of this component occurs. Conversely, a decrease in the concentration of one of the components leads to a shift in the equilibrium towards the formation of this component.
Thus, an increase in the concentration of substance A or B shifts the equilibrium in the forward direction; an increase in the concentration of substance C or D shifts the equilibrium in the opposite direction; a decrease in the concentration of A or B shifts the equilibrium in the opposite direction; a decrease in the concentration of substance C or D shifts the equilibrium in the forward direction. (Schematically, you can write: -C A or C B ®; -C C or C D ¬; ¯ C A or C B ¬; ¯ C C or CD ®).
The effect of temperature. The general rule that determines the effect of temperature on equilibrium has the following formulation: an increase in temperature contributes to a shift in equilibrium towards an endothermic reaction (- Q); lowering the temperature contributes to a shift in the equilibrium towards an exothermic reaction (+ Q).
Reactions that proceed without thermal effects do not shift the chemical equilibrium with a change in temperature. An increase in temperature in this case only leads to a more rapid establishment of equilibrium, which would be achieved in the given system even without heating, but over a longer time.
Thus, in an exothermic reaction (+ Q), an increase in temperature leads to a shift in the equilibrium in the opposite direction and, conversely, in an endothermic reaction (- Q), an increase in temperature leads to a shift in the forward direction, and a decrease in temperature in the opposite direction. (Schematically, you can write: at +Q -T ¬; ¯T ®; at -Q -T ®; ¯T ¬).
Influence of pressure. As experience shows, pressure has a noticeable effect on the displacement of only those equilibrium reactions in which gaseous substances participate, and in this case, the change in the number of moles of gaseous participants in the reaction (Dn) is not equal to zero. With an increase in pressure, the equilibrium shifts towards the reaction that is accompanied by the formation of a smaller number of moles of gaseous substances, and with a decrease in pressure - towards the formation of a larger number of moles of gaseous substances.
Thus, if Dn = 0, then pressure does not affect the shift in chemical equilibrium; if Dn< 0, то увеличение давления смещает равновесие в прямом направлении, уменьшение давления в сторону обратной реакции; если Dn >0, then an increase in pressure shifts the equilibrium in the opposite direction, and a decrease in pressure - in the direction of a direct reaction. (Schematically, it can be written: at Dn = 0 P does not affect; at Dn 0 -P ¬, ¯P ®). Le Chatelier's principle is applicable to both homogeneous and heterogeneous systems and gives a qualitative characteristic of the equilibrium shift.
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.
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.
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.
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.
Reactions are the "verbs" of chemistry - the activities that chemists study. Many reactions go to completion and then stop, meaning that the reactants are completely converted to products without being able 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.
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.
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 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 forward 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.
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.
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.
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.
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.
