Chemical equilibrium is defined as the state in which both products and reactants are present in the present concentrations which have no further tendency to change with time so that there is no observable change in the properties of the system which is present. This results in the state when the forward reaction proceeds at the same rate as the reverse reaction. The reaction rates of the backward and forward reactions are generally not zero, but equal. Thus, no net changes are seen in the concentrations of reactants and products. Such a state is called a dynamic equilibrium state.
Berthollet developed a chemical equilibrium in 180, he found that some chemical reactions are reversible. Forward and backward (reverse) reaction rates are equal for any reaction mixture to exist at equilibrium. Arrows pointing both ways to indicate equilibrium in the following chemical equation, where A and B are reactant chemical species, product species are S and T, and the stoichiometric coefficients α, β, σ, and τ are of the respective reactants and products.
"Far to the right" is said to lie to the equilibrium concentration position of a reaction if, at equilibrium, nearly all the reactants are consumed. "Far to the left" is said Conversely to the equilibrium position if hardly any product is formed from the reactants.
Based on Berthollet's ideas, Guldberg and Waage in 1865, proposed the law of mass action, where A, B, S, and T are active masses and k+ and k− are rate constants.
The numerator is formed from the products by convention. For the reaction of one-step, the mass law action is only valid, that proceed through a single transition state and is not valid in general because the rate of equations do not, in general, follow the reaction of stoichiometry as Guldberg and Waage had proposed (see, for example, nucleophilic aliphatic substitution by SN1 or reaction of hydrogen and bromine to form hydrogen bromide). However, for chemical equilibrium, the equality of forwarding and backward reaction rate is a necessary condition, though it is not sufficient to explain why equilibrium occurs.
The constant of equilibrium for a reaction is indeed a constant even despite the failure of this derivation, activities independent of any of the various species involved, though it is temperature-dependent as observed by the van 't Hoff equation. A catalyst addition will affect both the forward reaction and the reverse reaction in the same way and will not have an effect on the equilibrium constant. Both reactions will speed up by the catalyst and increase the speed at which the equilibrium is reached.
Reactions do occur at the molecular level, Although the macroscopic equilibrium concentrations are constant in time. For example, when the acetic acid dissolved in water it forms acetate and hydronium ions,
CH3CO2H + H2O ⇌ CH3CO−2 + H3O+
Types of Chemical Equilibrium
There are two types of chemical equilibrium:
Homogeneous Chemical Equilibrium
In Homogeneous Equilibrium, the reactants and the products of chemical equilibrium are all in the same phase. It can be further divided into two types:
a) Reactions in which the number of molecules of the products is equal to the number of molecules of the reactants. For example,
b) Reactions in which the number of molecules of the products is not equal to the total number of reactant molecules. For example,
Heterogeneous Chemical Equilibrium
In, Heterogeneous Equilibrium the reactants and the products of chemical equilibrium are present in different phases. Examples of heterogeneous equilibrium are listed below.
Thus, the different types of chemical equilibrium are based on the phase of the reactants and products.
Factors Affecting Chemical Equilibrium
Temperature, pressure, and concentration of the system which affect equilibrium are several factors. Some of the important factors affecting chemical equilibrium are discussed below.
Change in Pressure:
Due to the change in the volume, the change in pressure happens. The gaseous reaction can be affected If there is a change in pressure, as the total number of gaseous reactants and products are now different. As per Le Chatelier’s principle, in heterogeneous chemical equilibrium, as the volume is independent of pressure, the change of pressure in both liquids and solids can be ignored.
Change in Temperature:
The temperature effect on chemical equilibrium depends upon the sign of ΔH of the reaction and follows Le-Chatelier’s Principle.
The rate of reaction is also affected along with equilibrium constant, by the change in temperature. According to Le Chatelier’s principle, the equilibrium shifts towards the reactant side when the temperature increases in case of exothermic reactions, for endothermic reactions the equilibrium shifts towards the product side with an increase in temperature.
The Importance of Chemical Equilibrium in Real Life
Chemical equilibrium plays a vital role in numerous industrial applications. For instance, it enables the production of ammonia through Haber's process. In this method, nitrogen and hydrogen combine to create ammonia. Achieving a higher yield of ammonia is possible at lower temperatures, higher pressures, and when iron serves as a catalyst.
Additionally, in the contact process, we can produce sulphuric acid. The core reaction involves oxidizing sulphur trioxide, and the concept of chemical equilibrium is crucial in making this process successful.
Examples of Chemical Equilibrium
In chemical reactions, we have reactants turning into products through the forward reaction, and then the products can go back to being reactants through the backward reaction. These reactants and products are different in their makeup.
As the reaction progresses, there comes a point when the rates of the forward and backward reactions balance out. At this moment, the number of reactants turning into products is equal to the number of products reverting to reactants. This is when we say the reactants and products are in a state of chemical equilibrium.
For instance, let's take a look at some equilibrium reactions like N2O4 ⇌ 2NO2, PCl5 ⇌ PCl3 + PCl2, and N2 + H2 ⇌ 2NH3.
Chemical equilibrium is crucial in many industrial processes. It helps us make things like ammonia using the Haber's process, where nitrogen and hydrogen join forces to create ammonia. We get the best results when it's chilly, under high pressure, and with a little help from iron as a catalyst.
Another superhero moment is seen in the preparation of sulfuric acid through the contact process. Here, we see the transformation of sulfur dioxide into sulfur trioxide, and it's the equilibrium that plays a pivotal role in making this process a success. So, chemical equilibrium isn't just a fancy concept; it's a crucial part of many industrial processes, making our lives better in more ways than we realize.
Chemical equilibrium is a fascinating concept that defines the state in which reactants and products coexist in constant concentrations, leading to no observable changes in the system's properties. This equilibrium occurs when the forward and reverse reactions proceed at the same rate. The historical roots of this idea trace back to Berthollet's observations in 1800, followed by Guldberg and Waage's law of mass action in 1865. While this law doesn't hold for all reactions, it's essential for understanding equilibrium.
Chemical equilibrium is classified into two main types: homogeneous and heterogeneous, depending on the phase of reactants and products. Temperature, pressure, and concentration are key factors influencing equilibrium, following Le Chatelier's Principle.
This equilibrium is not just a theoretical concept but holds immense practical importance in various industrial processes, such as the production of ammonia and sulfuric acid, where precise control of conditions is crucial for higher yields and successful reactions.